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552 Cards in this Set
- Front
- Back
Metric Abbreviations: Mega, Kilo, Hecto, Deka, Deci, Centi, Milli, Nano
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Mega=M; Kilo=k; Hecto=h; Deka =da; Deci =d; Centi =c; Milli=m; Nano=n
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Conversion Factor
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A fraction which expresses an equality between two units of measurement and can be used to convert from one to the other (ex. 1kg/1000g)
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Use conversion factors to solve: 1) How many kilograms in 2000 g? 2) How many feet in 60"?
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1) (2000g)(1kg/1000g) = 2kg; 2) (60")(1'/12") = 5 ft
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Significant digits
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Digits which are measured. All non-zero digits are significant. Zeros are significant unless they are placeholders
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Scientific notation
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A number expressed as Ax10B. "A" is between 1.00 and 9.99 and "B" is an integer.
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Write conversion forumlas: 1) Celsius to Fahrenheit; 2) Celsius to Kelvin
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1) ˚f = (1.8 x ˚celsius) + 32; 2) Kelvin = ˚celsius +273
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Energy
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The ability to do work; it is released or absorbed during chemical reactions in the form of heat, light, electricity. (calorie, Joule: 1cal = 4.18J)
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Matter
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A substance that occupies space has mass.
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Compare Weight vs. Mass
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Mass is the amount of matter. Weight measures gravitational force. Mass never varies. Weight can vary.
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Compare potential energy and kinetic energy
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Potential energy is energy due to position. KE is energy of motion.
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Explain how to convert a number greater than 1 to scientific notation
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Move decimal point to left until only 1 digit remains to left. Indicate number of moves as a positive exponent of 10. 3301 = 3.301x10^3
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Explain how to convert a number less than 1 to scientific notation
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Move decimal point to right until only 1 digit remains to left. Indicate number of moves as a negative exponent of 10. (0.00356 = 3.56 x 10-3
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Explain the rules for multiplying the numbers in scientific notation
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Multiply the first numbers and add the exponents. (3x10^5)(2x10^3)= 6x10^8
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Explain the rules for dividing the numbers in scientific notation
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Divide the first numbers and subtract the exponents. (8x10^6)/(2x10^10) = 4x10^-4
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Physical properties
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Properties which can be observed without changing the substance into something different. Color, odor, hardness, density, luster, state, conductivity, solubility, boiling and melting points
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Chemical properties
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A chemical property is observed when a substance changes into a new substance. Iron forms rust in air & water; gasoline burns in oxygen
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Physical change
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Substance changes form or state only. Boiling, melting, freezing, dissolving, grinding, cutting
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Chemical change
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Where new substances are formed with new chemical and physical properties. Oxygen & hydrogen form water; sodium & chlorine form sale (sodium chloride)
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Density: Write the general equation and three standard units
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The mass of a unit volume of a substance. Density = mass/volume = g/mL; g/L; kg/L
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Element. List some examples
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A substance is composed of identical atoms. Gold, silver, oxygen, hydrogen, lead, chlorine, helium, iron, copper, fluorine, arsenic
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Compound. List some examples
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substance composed of two or more elements chemically combined. Water - H2O; Salt - NaCl; Sugar - C6H12O6; Ammonia - NH3
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Mixture. List some examples.
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A combination of substances held together by physical means (dirt, milk, soup, saltwater, granite)
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Homogeneous and Heterogeneous mixtures. Provide examples.
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Homogeneous mixtures are uniform in composition (air, metal alloy, salt water). Heterogeneous mixtures are not uniform in composition (dirt, spaghetti sauce)
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Three postulates of Dalton's Atomic Theory
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1) An element is composed of identical atoms; 2) Atoms of different elements have different properties; 3) Compounds are atoms of 2 or more elements chemically combined
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The Law of Conservation of Mass
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During a chemical reaction, matter is neither created nor destroyed
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The Law of Constant Composition
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A compound always contains the same elements combined in the same proportions by mass (H2O) is 88% oxygen no matter where it is found)
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Law of Multiple Proportions
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The same elements may combine to form more than one compound. The ratios of atoms are in small whole numbers (H2O and H2O2)
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Atomic Mass Unit
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the mass of a proton or neutron is equal to 1 atomic mass unit. Symbol - "amu"; 1 amu=1.66x10^-24
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Atomic Number: What are the atomic numbers of helium, hydrogen, carbon, oxygen?
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The number of protons in the nucleus of an atom of an element. Helium-2; carbon-6; hydrogen-1; oxygen-8
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Mass number
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The sum of protons plus neutrons n the nucleus of an atom
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Isotope
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Atoms which contain the same numbers of protons but different numbers of neutrons (ex. Hydrogen has 3 isotopes with mass numbers of 1,2,3)
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Molecule
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A group of two or more atoms held together by chemical bonds
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Ion (provide examples)
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An atom or group of atoms which contains a positive or negative electrical charge (ex. Na+; Cl-; SO4^2-)
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Cation and Anion (provide examples)
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cation - positively charged ion (Na+; Fe+2; NH4+; Ag+); anion - negatively charged ion (Cl-; SO4-2; OH-; P-3)
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Valence electroncs
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The electrons found in the outermost energy level of an atom
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Oxidation number
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A number (positive or negative) representing the charge on an ion or atom involved in a chemical bond
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Three general rules for determining oxidation numbers
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1) Atoms of uncombined elements equal 0; 2) Hydrogen = +1 (in metallic hydrides =-1); 3) Oxygen = -2 (in peroxides =-1); (bonded with fluorine =+2)
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Empirical Formula. What are the molecular and empirical formulas of peroxide?
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An expression which gives the relative numbers of atoms of the elements in a molecule. Expressed as the lowest possible set of integers (H2O2, HO)
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Molecular Formula. What is the molecular formula for ammonia?
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An expression stating the number and kind of each atom present in a molecule of a substance (NH3 has 1 nitrogen atome and 3 hydrogen atoms in each molecule)
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Positive Ion: which elements tend to form them?
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Metals tend to form positive ions by losing electrons (Na → Na+ e-)
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Negative Ion: which elements tend to form them?
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Non-metals form negative ions by gaining electrons (Cl +e- →Cl-)
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Write oxidation numbers for ions of: Group IA & IIA; Group VIA & VIIA
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IA→+1; IIA→+2; VIA→-2; VIIA→-1
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Write formula for compounds of: 1) Sodium & Sulfate; 2) Magnesium & Nitrite; 3) Aluminum & Phosphate
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1) Na2SO4; 2) Mg(NO2)2; 3) AlPO4
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Names of the ionic compounds: 1) FeCL3; 2) FeO; 3) Cu(OH); 3) Cu3PO4
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1) iron (III) chloride; 2) iron (II) oxide; 3) copper (II) hydroxide; 4) copper (I) phosphate
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10 prefixes used to name covalent compounds
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Mono=1; Di=2; Tri=3; Tetra=4; Penta=5; Hexa=6; Hepta=7; Octa=8; Nona=9; Deca=10
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Forumlas and names of acids formed from: 1) F; 2) Cl; 3) Br; 4) I
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1) HF - hydrofluoric acid; 2) HCl - hydrochloric acid; 3) HBr - hydrobromic acid; 4) HI - hydriodic acid
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Names of acids: 1)H2SO4; 2) HNO2; 3) H3PO4; 4)HClO
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1) sulfuric acid; 2) nitrous acid; 3) phosphoric acid; 4) hypochlorous acid
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Write the equation for Percent Composition. What is the percent composition of Ca in CA(OH)2?
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%=(me/FW)(100%); %= percent composition of the element; me=mass of element in one formula unit; FW = formula weight; (40/74)(100%) = 54% Ca in Ca(OH)2
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Atomic Mass
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A weighted average mass of the atoms of an element (assuming the carbon-12 isotope is exactly 12) (ex. Atomic mass of C=35.45 is calculated from two isotopes. Cl-35 and Cl-36)
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Write the equation to calculate the atomic mass of an element
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Atomic mass of X = ((mx1)(%x1)/100%) + ((mx1)(%x1)/100%)) + etc.; mx1, mx2, mxN = atomic masses of each isotope of element; %x1, %x2, %xN = percent composition of each isotope
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Molecular Mass
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Found by adding all the atomic masses of an element. Ex. H2O. Molecular mass = 18; H2 = 2(1) O = 16
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Avogadro's Number
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The number of atoms or formula units in "x" grams of an element or molecule where "x" is the atomic or molecular mass. (Always equal to 6.02 x 10^23)
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Mole
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6.02x10^23 items, can be anything. The number of atoms in one mole (atomic mass in grams) of a monoatomic element. The number of formula units in one mole (formula mass in grams) of an ionic compound. The number of molecules in one mole (formula mass in grams) of a molecular substance)
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For any substance, write a general formula to convert from Moles to Grams
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g=(n)(MM) where, n = moles MM = molecular mass g = grams
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For any substance, write a general formula to convert from Grams to Moles
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n = g/MM; n = moles; MM = molecular mass; g = grams
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For any substance, write a general formula to convert from Moles to Number of Particles
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(n)(6.02x10^23) = P; n = number of moles; P = number of particles
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For any substance, write a general formula to convert from Number of Particles to Moles
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n = P/6.02x1023; n = number of moles; P = number of particles
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Reactants; Products
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1) The starting materials in a chemical reaction; 2) The substances formed in a chemical reaction
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Balance Equations: H2 + N2 ↔ NH3; NaCL + Br2 ↔ NaBr + Cl2
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1) 3H2 + N2 ↔2NH3; 2) 2NaCl + Br2 ↔2NaBr + Cl2
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Diatomic; list 7 diatomic elements found in nature
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A molecule composed of two atoms. H2, N2, O2, Cl2, F2, Br2, F2, I2
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Synthesis or combination
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A reaction where 2 or more elements form a compound; N2 + 3H2 ↔2NH3; 2H2 + O2 ↔ 2H2O
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Decomposition
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A reaction where a compound breaks down into elements; CO2 ↔ C + O2; 2CaO ↔2Ca + O2
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Single Replacement
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A reaction involving the replacement ina compound of an element by another element. Zn + CuCl2 ↔ Cu + ZnCl2
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Double replacement
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A reaction where two ionic substances "trade" anions; NaCl + AgNO3 ↔ NaNO3 + AgCl
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Period
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A horizontal row of elements in the periodic table. All have the same number of shells of e-. Across the period, the elements' properties change.
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Column or Family
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A verticle group of elements on the periodic table in the same column. They have similar properties and the same number of valence electrons.
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Metalloid. List 5 examples.
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An element with the properties of both metals and non-metals. Ex. Si, As, Ge, Sb, Te
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Metals. List 6 characteristics
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Solid (except Hg); malleable & ductile; conduct heat & electricity; shiny reflective & lustrous; lose e- to form cations; good reducing agents
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Non-metals. List 7 characteristics
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1/2 are gases; solids are brittle; poor conductors of heat and electricity; dull & non-reflective; gain e- to form anions; good oxidizing agents
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Transition element
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A metal having two incomplete shells of electrons; many have multiple oxidation states; less active than family IA & IIA. Ex., Fe, Ag, Au, Cr, W
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Describe the change in atomic radius across the periodic table
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The atomic radius tends to decrease from left to right across the table and increase down the columns
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Ionization energy
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The energy change required for the removal of the outermost electron from the gaseous atom to form a +1 ion
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Describe the change in Ionization Energy across the periodic table
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Ionization energy tends to increase from left to right across the table and decrease down the columns
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Electronegativity
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The ability of an atom to attract to itself the e- in a covalent bond. Values range from 0.7 (Cs) to 4.0 (F)
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Describe the change in electronegativity across the periodic table
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Electronegativity tends to increase across the table left to right and decrease down the columns
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Describe the change in metallic activity across the periodic table
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Metallic activity decreases across the table (left to right) and increases down the columns
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Describe the change in nonmetallic activity across the periodic table
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Non-metallic activity increases across the table (L to R) and decreases down the columns
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Compare the ionic radii of 1) Cl, S, P; 2) Na, Mg, Al; What is the reason for their different size?
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1) Cl-1 is smallest; P-3 is largest (the P atom gained the most e-); 2) Na+1 is largest Al+3 is smallest (the Al atom lost the most e-)
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Quantum
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A packet of energy associated with a specific wavelength of electromagnetic radiation
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Quantum Number
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A number used to describe the energy levels available to an electron. Each electron in an atom has a unique set of four.
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Emission Spectrum
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A bright line spectrum formed when energy absorbed by an element is emitted at specific wavelengths. Each element has a unique spectrum.
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Absorption Spectrum
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A dark-line spectrum formed when white light is passed through a vaporized element and a few specific wavelengths are absorbed.
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List three basic postulates of the bohr model for the hydrogen atom
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1) e- are present only in specific energy states; 2) a quantum of energy is absorbed or emitted to change energy levels; 3) a quantum is the smallest amount of energy that can be gained or lost
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Ground State
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Electron is at its lowest energy level as close to the nucleus as possible
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Excited State
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An electron absorbs energy and moves to a higher energy level above the ground state
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List four quantum numbers and their symbols
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1) Principal energy level "n"; 2) sublevel "l"; 3) orbital "m1"; 4) spin "ms"
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List the names of the first four energy levels (or shells)
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1 or K; 2 or L; 3 or M; 4 or N
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List the names of the four sublevels and their electron capacities
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sublevel s holds two electrons; p holds 6; d holds 10; f holds 14
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Which sublevels are present in energy levels 1, 2, 3, and 4-7?
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1: s; 2: s, p; 3: s, p, d; 4-7: s, p, d, f
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Describe the shapes of the orbitals: 1) s; 2) p; 3) d; 4) f
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1) s- sphere; 2) p - dumbbell shape with 2 lobes; 3) d - double dumbbell; most have 4 lobes; 4) f - most have 8 lobes
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What is the order for filling sublevels (aufbau process) from lowest to highest energy
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1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d
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Electron configuration
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The distribution of electrons into shells and sublevels for an atom of an element. Each element has a unique electron configuration.
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Write the electron configuration for 1) Lithium; 2) Iron
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1) 1s2, 2s1; 2) 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d6
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Valence
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The number of electrons in the atom's highest numbered shell.
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What are the valences of the elements of families IA through VIIIA?
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The valences of elements in groups IA - VIIIA are the element's column number. For example, the valence of Na is 1; O is 6
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Electronegativity difference
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A number found by taking the difference between the electronegativities of two atoms in a bond. Its value determines the type of bond.
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Describe an ionic bond in terms of electronegativity difference
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When the EN values differ by 1.7 or more. The atom with higher EN borrows the electrons from the atom with lower EN. The resulting positive and negative ions attract.
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Nonpolar covalent bond
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when the EN difference is very small (less than 0.5). Two bonded atoms share the valence e-. The resulting molecule has no electrostatic charge.
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Polar covalent bond
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When the EN difference is between 0.5 and 1.7, the bonding electrons stay closer to the more electronegative atom. Electrons are shared unequally.
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Coordinate covalent bond
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When both electrons in a covalent bond are supplied by one atom
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Metallic Bonds
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A sea of electrons surrounding positive metal ions
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Hydrogen bonding
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Formed when hydrogen is bonded to oxygen, fluorine, or nitrogen. The hydrogen of one molecule becomes attracted to the electronegative element of the other molecule. These intermolecular attractions cause higher boiling points than predicted
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Octet Rule
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Atoms tend to gain or lose outer shell electrons in order to achieve a noble gas configuration of 8 electrons
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Double and Triple covalent bond
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In a double bond, wo pairs of electrons are shared. In a triple bond, three pairs of electrons are shared.
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Resonance structures
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Where there is more than one possible bonding structure in a molecule
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Hybrid orbitals (list three types)
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Where 2 or more pure atomic orbitals are mixed to form identical hybrid orbitals (ex. Sp, sp2, sp3)
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Describe hybrid bonding in water, ammonia, methane
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sp3 bonding results in a tetrahedron shape with bond angles of 109.5˚ in methane and slightly less in water and ammonia.
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Sigma bonds & Pi bonds
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A sigma bond is present between any 2 orbitals except when 2 p orbitals share electrons; then this is a pi bond
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List properties of ionic substances
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Solids at 25C; Non-conducting as solids but conducting as aqueous solutions or liquids. Conducting as aqueous solutions or liquids. High melting & boiling points; Brittle; Low volatilities
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List properties of molecular substances
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Non-conducting as liquids and solids; Volatile liquids & solids; Many are gases at 25C; Low melting and boiling points; Soft and waxy solids
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Exothermic
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A chemical reaction which evolves heat
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Endothermic
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A chemical reaction which absorbs heat
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Enthalpy
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The heat content of a system
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Enthalpy change (state the equation)
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The difference in heat content between the products and the reactants (ΔH = ΣHproducts - Σhreactants)
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Exothermic reaction: ΔH is (positive or negative); Enthalphy is (increased or decreased)
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Negative; Decreased
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Endothermic Reaction: ΔH is (positive or negative); Enthalphy is (increased or decreased)
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Positive; Increased
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List in the order of increasing enthalpy: solid, gas, liquid
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Solid (least enthalpy) -> Liquid -> Gas (most enthalpy)
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Thermochemical equation
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A balanced chemical equation which includes the enthalpy change. (H2(g) +1/2O2(g) →H2O(l) ; ΔH = -285kJ
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In a thermochemical equation, what happens to ΔH when the moles of reactants double?
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Enthalpy is directly proportional to mass. Therefore when the moles double, so does ΔH.
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How are ΔH for a forward and ΔH for a reverse reaction related?
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forward is equal in magnitude but opposite in sign to ΔH reverse.
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Hess' Law
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for a reaction is the same regardless of the path travelled from reactants to products.
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Molar Heat of Formation
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The molar heat of formation of a compound is equal to ΔH when 1 mold of compound is formed from its elements at 1 atm and 25C
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Write the general equation to calculate ΔH for a chemical reaction
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ΔH = Σ(ΔHF-products) - Σ(ΔHF-reactants); where ΔHf = Heat of formation of reactants or products
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Specific Heat. Give value for liquid water in calories and joules.
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The amount of heat required to raise one gram of substance 1C. Water: 1cal/g-C or 4.18 J/g-C
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Write an equation to calculate energy change when a fixed mass of substance changes temperature
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ΔH = mCpΔT, where ΔT = temperature change, Cp = Specific Heat; ΔH = heat absorbed or given off
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Charles Law
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V1/T1 = V2/T2 (pressure and amount of gas are constant; V=volume; T=Kelvin)
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Boyles Law
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P1V1 = P2V2 (Temperature and amount of gas are constant, P=pressure, V=vol)
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Combined Gas Law
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P1V1/T1 = P2V2/T2 (amount of gas is contant. P=pressure, V=vol; T=Kelvin)
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Dalton's Law of Partial Pressures
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In a gas mixture, the total pressure equals the sum of the partial pressures of each component. Ptotal = P1 + P2 + P3…
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Ideal Gas Law
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PV = nRT, (P=pressure in atm; V = volume in L; n = # moles; R = 0.0820 Latm/Mol-K; T = Temp in K)
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STP
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Standard temperature (0C or 273 K) and Standard Pressure (1 atm or 760 torr)
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Gay-Lussac's Law of Combining Gas Volumes
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When only gases are involved in a reaction, the volumes of reactants and products are in a small, whole number ratio.
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Avogadro's Law re. gases
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Under the same conditions of temperature and pressure; equal volumes of gases contain equal numbers of moles
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What is the volume of 1 mole of any gas at STP?
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22.4 L
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Phase equilibrium
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For a liquid in a closed container, when the rates of evaporation (liquid to gas) and condensation (gas to liquid) equalize; the concentration of each is stable.
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Dynamic equilibrium
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In a closed container where opposing changes are taking place at equal rates; the concentration of all components remains constant.
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Normal Boiling Point
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The temperature at which a liquid phase becomes a gas phase at a pressure of 1 atm.
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Heat of Fusion (value for water)
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The amount of energy required to change a gram of substance from solid to liquid at its melting point (water = 80 cal/g)
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Heat of Vaporization (value for water?)
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The amount of energy required per gram to change a liquid to a gas at its boiling point (water = 540 cal/g)
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Triple Point
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The only temperature and pressure combination at which the 3 phases of a substance (solid, liquid, gas) can co-exist in equilibrium
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Vapor Pressure
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The pressure the gas phase exerts on its liquid phase in a closed container. This pressure varies with temperature
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Molarity - general equation
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The number of moles of solute it a liter of solution; M = n/L
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Solute (provide example)
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The substance dissolved in another (solvent). Salt is the solute in salt water.
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Solvent (provide example
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A substance, usually a liquid, into which another substance (solute) is dissolved. Water is the solven in iced tea.
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Solution (provide example)
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A liquid, gas or solid phase containing 2 or more components uniformly dispersed (air, coffee, saltwater)
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Solubility curves
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A curve for a given substance which shows how much dissolves in a given amount of solvent at different temperatures.
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How do temperatures and pressure affect the solubility of a solid?
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Solubility usually increases with increasing temperature. Pressure has little effect.
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How do temperature and pressure affect the solubility of a gas?
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Solubility usually decreases with increasing temperature. Solubility increases in direct proportion to an increase in pressure.
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Which three factors affect the rate of solubility?
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Pulverizing; stirring; heating
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What is a general rule for solubilities of polar and nonpolar compounds?
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"Like dissolves like"; Ionic and polar solvents dissolve ionic, polar solutes (water dissolves salt). Non polar solvents dissolve nonpolar solutes (acetone dissolves gasoline)
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List some basic facts about solutions
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Particle size less than 1 mmicron; Clear (may be colored); Particles don't settle; Can pass through membranes; Particles not visible
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List some basic facts about colloids
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Particles measure 1-100 mmicrons; Particles don't pass through a membrane; Show brownian motion and the Tyndall effect; Particles don't settle; Clear and pass through filter paper
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List some basic facts about suspensions
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No brownian motion; Don't pass through filter paper or a membrane; Cloudy but particles settle on standing; Particles visible with microscope or eye
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How many grams of NaCl are required to prepare 500 grams of a 5% solution?
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%Concentration = (gNaCl/gsolution)(100%); 5% =(x/500g)(100%);x=25 g NaCl
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Calculate the molality of 10 moles of H2SO4 dissolved in a 4 kg of water
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Molality = Moles Solute / kg solvent = 10/4 = 2/5 Molal
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Gram-equivalent weight
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The amount of substance which reacts with or displaces 1 mole of H+ ions.
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Normality
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The number of gram-equivalent weights in a liter of solution
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in H2O solutions: 1) How many ˚C is the freezing point depressed for each molal of solute? 2) How many ˚C is the boiling point elevated for each molal of solute?
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1) 1.86C for each molal of particles of solute; 2) 0.51C for each molal of particles of solute
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List five factors that control reaction rate
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nature of reactants; exposed surface area; concentrations; temperatures; presence of catalyst
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State the collision theory of reaction rates
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There must be collisions between reactants. Reaction rate depends on number of collisions per unit time and the percent which are successful (Have sufficient energy)
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How is the reaction rate related to concentration?
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Reaction rate is directly proportional to the concentrations of reactants
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Activation energy
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The energy necessary for a reaction to begin. Obtained from the kinetic energy released during collision
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Catalyst
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a catalyst is introduced into a reaction to speed it up or slow it down. It is not consumed. An increase or decrease of activation energy results from an alternate reaction path.
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Law of Mass Action
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The rate of a chemical reaction is proportional to the product of the concentrations of the reactants
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Specific Rate Constant
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Symbol is "k" in a rate equation. A constant specific to temperature and reaction which is part of every rate equation
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Reversible reaction
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A system where the following opposite reactions are taking place: reactant becoming product; product becoming reactant
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Equilibrium
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The point in a reversible reaction where the forward and reverse reactions are taking place at the same rate.
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Are concentrations of product and reactant equal at equilibrium?
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No. The are constant but not equal. Their relative concentrations are determined by the value of the equilibrium constant at that temperature.
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Write the equilibrium expression for aA + bB ↔ cC + dD
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Keq = [C]^c[D]^d / [A]^a[B]^b, where Keq = Equilibrium constant
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How are reactant & product concentrations related to the magnitude fo Keq?
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Keq is large: [reactant] is small and [product] is large; when Keq is small: [reactant] is large and [product] is small
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Le Chatelier's Principle
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If stress is placed on a system at equilibrium, the equilibrium shifts in order the counteract the effects of the stress and regain equilibrium
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How does a concentration change affect equilibrium?
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If one of the substances is added or removed, all the concentrations of substances adjust to a new equilibrium with the same Keq
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How does a change in temperature affect equilibrium
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The reaction shifts to a new equilibrium point with a new Keq. If the temperature is raised, the equilibrium is shifted to reaction which absorbs heat.
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How does a pressure change affect equilibrium?
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Only in reactions where gases are involved. The reaction will shift to oppose pressure change, resulting in fewer moles of gas particles
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Ionization constant. Write the expression for the ionization of acid "HA"
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For substances in solution that partially ionize. An equilibrium expression may be written with Ki; Ki = [H+][A-] / [HA]
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What are the two driving forces that control reactions?
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A drive towards increased entropy (disorder). A drive towards decreased enthalpy (lower heat content)
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Second Law of Thermodynamics
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The entropy of the universe increases for any spontaneous process.
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Free Energy Change (ΔG). Write the free energy equation.
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A property which reflects a system's capacity to do useful work. ΔG = ΔH - TΔS; G=free energy; S=entropy; H=enthalpy; T=kelvin
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How can ΔG be used to predict if a reaction is spontaneous?
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When ΔG is positive it is not spontaneous; negative, it is spontaneous; equals 0 it is at equilibrium.
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Ionization Constant
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Kw = 1 x 10-14 at 25C
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pH. For what values is a solution acidic, basic and neutral?
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pH = -log[H+] ; The degree of acidity of a solution. <7 =acid; 7 = neutral; >7 = basic.
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pOH. For what values is a solution acidic, basic, neutral?
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pOH = -log[OH-]; The degree of basicity of a solution. <7 = basic; 7 = neutral; >7=acid.
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How are pH and pOH of a solution related?
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The sum of the two values = 14; (pH +pOH = 14)
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How are the concentrations of [H+] and [OH-] related in a solution?
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The product of the concentrations = 1x10-14; [H+][OH-] = 1x10-14
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Solubility Product Constant (Ksp)
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An equilibrium exists in a saturated solution between dissolved and undissolved solute. Ksp is the equilibrium constant for this reaction.
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Write the solubility product expression for AgCl ↔ Ag+ + Cl-
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Ksp = [Ag+][Cl-]
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Six common characteristics of acids
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Form H2O solutions; Conduct electricity; React with active metals; Turn blue litmus red; Neutralize bases; Sour taste
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Seven common characteristics of bases
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Form H2O solutions; Conducts electricity; Turns red litmus blue; Feels slippery; Caustic; Neutralizes acids; Bases + fats form SOAP
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Neutralization Reaction - Write equation for hydrochloric acid and sodium hydroxide
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acid + base → salt + water; HCl + NaOH →NaCl +H2O
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Arhennius Theory
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An acid yields protons in solution (H+ ions); A base yields hydroxide ions in solution (OH- ions)
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Bronsted Theory
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An acid is a proton donor. A base is a proton acceptor.
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Lewis Theory
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An acid is an electron pair acceptor; a base is an electron pair donor.
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Conjugate Base - Write conjugate base of HCl
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When a Bronsted acid donates a proton, it becomes its conjugate base (conjugate base of HCl is Cl-)
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Conjugate Acid - Write conjugate acid of I-
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When a Bronsted base accepts a proton it becomes its conjugate acid (conjugate acid of I- is HI)
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Calculate the volume of 10M NaOH needed to titrate 5L of 2M HCl
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MA x VA = MB x VB; 2M x 5L = 10M x VB; 1L = volume of base
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Electrolyte (give 3 examples)
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A solute whose aqueous solution contains ions and conducts electricity (acids, bases, salts)
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Nonelectrolyte (give 2 examples)
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A solute whose aqueous solution does not conduct electricity (sugar, benzene, most organic compounds)
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Anode
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A positively charged electrode which attracts anions. Where oxidation takes place.
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Cathode
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A negatively charged electrode which attracts cations. Where reduction takes place.
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Oxidation
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The loss of electrons (ex. Cumetal →Cu+2 +2e-
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Reduction
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The gain of electrons (Zn+2 +2e- → Znmetal)
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Electrode Potential
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A measure in volts of the tendency of atoms to gain or lose electrons. (Relative to a H2 oxidation reaction which has an assigned value of zero)
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Half Reaction
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One of the two parts, either the reduction or the oxidation, of an oxidation-reduction reaction
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How can you determine if a redox reaction will take place spontaneously
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Add the electrode potentials of the two half reactions. If the result is positive, the reaction is spontaneous; if negative, the reaction is not spontaneous.
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Faraday
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A unit of electric charge which deposits by electrolysis one equivalent weight of an element. Equals 96,500 coulombs
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Equivalent weight
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The number of grams of an element which will accept or donate 1 mole of electrons
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Oxidizing agent
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A substance which causes another substance to be oxidized (oxidizing agent is simultaneously reduced)
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Reducing agent
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A substance which causes another substance to be reduced (reducing agent is simultaneously oxidized)
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List some basic facts about carbon bonding
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Carbon forms more compounds than any other element; Each atom requires 4 covalent bonds; Carbon can form long chains and rings; Bonds commonly to O, H, N, S, P, and halogens
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Alkane
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A series of hydrocarbons with only single covalent bonds (CnH2n+2)
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Alkene
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A series of hydrocarbons containing at least one double covalent bond (CnH2n)
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Alkyne
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A series of hydrocarbons containing at least one triple covalent bond (CnH2n-2)
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List prefixes for naming hydrocarbons for 1-10 carbons in a molecule
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1 meth; 2 eth; 3 prop; 4 but; 5 pent; 6 hex; 7 hept; 8 oct; 9 non; 10 dec
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List names for alkanes with 1 to 10 carbons in a molecular chain
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1 methane; 2 ethane; 3 propane; 4 butane; 5 pentane; 6 hexane; 7 heptane; 8 octane; 9 nonane; 10 decane
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Substitution reaction. Which elements commonly substitute in alkanes?
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A reaction where one of the hydrogen atoms in a hydrocarbon is replaced by another. Usually a halogen. (CH4 + Br2→Ch3Br + HBr
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Cycloalkane
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An alkane which has a ring structure instead of a chain
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Saturated and Unsaturated. Which are sturated? Alkanes, alkenes, alkynes.
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A compound only containing single covalent bonds is saturated. Alkanes are saturated. A compound containing double or triple bonds is unsaturated. Alkenes and alkynes are unsaturated.
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Addition reaction
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In an unsaturated hydrocarbon, two atoms may be added to the structure across a double or triple bond (C2H2 + Br2 → Ch2Br2)
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Aromatic compounds
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Unsaturated ring structures with six carbon atoms. Benzene is the simplest aromatic. (CnH2n-6)
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Isomer
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Compounds with the same molecular formula but different structural formulas (different connectivity)
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Polymerization
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The combination of two or more unsaturated molecules to form a larger chain molecule. This is how plastics are made.
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Hydrogenation
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The process of adding hydrogen to an unsaturated hydrocarbon
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Dehydrogenation
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The process of removing hydrogen from a hydrocarbon
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Alcohol
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hydrocarbons which contain the hydroxyl functional group (OH-) attached to a saturated carbon (R-O-H)
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Aldehyde
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A hydrocarbon containing the aldehyde functional group (R-C(=O)-H)
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Carboxylic Acid
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A hydrocarbon containing the carboxyl functional group. (R-C(=O)-O-H)
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Ketone
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A hydrocarbon containing a ketone functional group (R-C(=O)-R')
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Ether
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A hydrocarbon containing an ether functional group. (R-O-R')
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Amine
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A hydrocarbon containing an amine functional group. (R-NH2)
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Ester
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A hydrocarbon containing an ester functional group. (R-O-C(=O)-R')
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Amino Acid
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Organic compounds which contain an amine and a carboxyl group. (H2N-CH(-R)-COOH)
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Compare primary and secondary alcohols
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Primary: The OH- group is attaced to the end carbon of the chain. Secondary: The carbon bearing theOH- goup is directly attached to two other carbons.
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Alpha particle
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a helium nucleus; charge =+2; High energy; Low velocity; Ejection reduces atomic number by 2amu and atomic weight by 4amu
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Beta particle
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An electron ejected from the nucleus when a neutron decays to a proton; Increases atomic number by one; High velocity; Low energy
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Gamma radiation
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Usually emitted with beta radiation; Has neither charge nor mass; High energy; Travels at the speed of light
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Half-life
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The time it takes for half of a radioactive sample to decay. It can range from a fraction of a second to many years.
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Transmutation
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The conversion of an element to a new element due to a change in number of protons. Ex. Alpha or beta decay
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Fission
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The breakdown of heavy nuclei into lighter nuclei. The source of nuclear power
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Fusion
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The joining of lighter nuclei to form heavier nuclei. Source of the sun's energy
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Write the nuclear equation: 1) U238 loses an alpha particle; 2) Th234 loses a beta particle.
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92U238 →90Th234 +2He4; 90Th234→ 91Pa234 + -1e0
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Atoms
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Tiny particles making up mass
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Nucleus
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Contains protons and neutrons (nucleons)
Held together by the strong nuclear force Surrounded by one or more electrons Radius = 10^-4 angstroms 1 angstrom = 10^-10 meters |
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Neutrons
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Part of nucleus
Together with proton, makes up nucleon Approximately same size and mass as proton (1 amu) Slightly heavier than proton No charge, electrically neutral |
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Protons
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Part of nucleus
Together with neutron, makes up nucleon Same size and mass as neutron (1 amu) Slightly lighter than neutron Positive charge |
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Electrons
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Surround nucleus at distance of 1-3 angstroms
Over 1800 lighter than nucleon Electrons (negative charge) and protons have opposite charges of equal magnitude |
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Electron charge (e)
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Charge of one electron
e = 1.6e^-19 coulombs (C) |
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Atom
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Electrically neutral
Made up of neutrons, protons and electrons Same number of protons as electrons Composed mostly of empty space |
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Elements
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A single atom
Building blocks of all compounds Cannot be decomposed into simple substances by chemical means Characterized by: 1. Mass number (A) 2. Atomic number (Z) 3. Atomic weight (amu) |
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Mass number (A)
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Number of protons plus neutrons
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Atomic number (Z)
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number of protons
Identity number of any element |
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Isotopes
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Two or more atoms of the same element that contain different numbers of neutrons
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Atomic Weight
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Also known as molar mass (MM or M)
Given in atomic mass units (amu or u) or grams/mole (g/mol) Actually a mass (ratio) and not a weight |
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Atomic mass units (amu)
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An amu is defined by carbon-12
1 atom of C-12 has an atomic weight of 12 amu All other atomic weights are measure against this standard Also known as a dalton |
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Mole
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Defined by C-12
Also known as Avogadro's number = 6.022e^23 The number of C atoms in 12 grams of C-12 6.022e^23 amu = 1 gram Moles = g/amu |
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Periodic table
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Lists the elements from left to right in order of their atomic numbers
Can be divided into: 1. Nonmetals (right) 2. Metals (left) 3. Metalloids (diagonal seperation between metals & nonmetals) |
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Period
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Each horizontal row of periodic table
Elements in the same family on the periodic table tend to have similar chemical properties Tend to make the same number of bonds Tend to exist as similarly charges ions |
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Groups or Families
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Vertical columns of periodic table
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Metals
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Large atoms that tend to lose electrons to form positive ions or form positive oxidation states
Metallic character (easy movement of electrons) All metals (except mercury) exists as solids at room temperature Form ionic oxides (ie: BaO) |
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Metallic Character
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Property of metals
Increases from right to left and top to bottom of periodic table 1. Ductility (easily stretched) 2. Malleability (easily hammered into think strips) 3. Thermal and electrical conductivity 4. Characteristic luster |
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Nonmentals
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Diverse appearances and chemical behaviors
Lower melting points than metals Form negative ions Make up molecular substances Form covalent oxides (ie: SiO2 or CO2) |
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Alkali metals
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Group (family) 1A in periodic table
metals Soft metallic solids Low densities and melting points Easily form 1+ cations Highly reactive, reacting with most nonmetals to form ionic compounds |
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Alkaline earth metals
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Group (family) 2A in periodic table
metals Harder, more dense and melt at higher temperatures than alkali metals Form 2+ cations Less reactive than alkali metals The heavier, the more reactive |
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Halogens
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Group (family) 17 in periodic table
Nonmetals & metalloids |
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Noble gases
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Groups (family) 18 in periodic table
Nonmetals Also known as rare gases Nonreactive, inert gases at room temperature |
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Metalloids
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Characteristics that resemble metals and nonmetals
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Representative or main-group elements
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Section A groups in periodic table
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Transition metals
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Section B groups in periodic table
When form ions, they lose electons from s-subshell first and then from d-subshell |
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Hydrogen
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Is unique and chemical/physical properties to do conform to any family
Nonmetal Colorless Odorless Diatomic gas |
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Group 4A
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Elements can form 4 covalent bonds with nonmetals
All (except carbon) can form 2 additional bonds with lewis bases Only carbon forms strong pi bonds to make strong double and triple bonds |
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Group 5A
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Elements can form 3 covalent bonds
All (except nitrogen) can form 5 covalent bonds by using d-orbitals Can bond with lewis base to form 6th covalent bond Nitrogen forms strong pi bonds to make double and triple bonds |
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Group 6A
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Elements called chalcogen
Oxygen and sulfur are most important Oxygen is second most electronegative element, divalent, can form strong pi bonds to make double bonds, reacts with metals to form oxides Sulfur can form 2, 3, 4 or 6 bonds and can pi bond to make double bonds |
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Group 7A
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Radioactively stable called halogens
1.Fluorine 2. Chlorine 3. Bromine 4. Iodine Highly reactive, like to gain electrons Fluorine makes only 1 bond, while other halogens can make more than 1 bond Bind to hydrogen to form hydrogen halides (soluble in water) Reacts with metals to form ionic halides |
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Diatomic molecules
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1. Hydrogen
2. Oxygen 3. Nitrogen 4. Halogens |
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Small atoms
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Make strong pi bonds due to overlap of p-orbitals
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Large atoms
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Make weak or are unable to make pi bonds due to lack of overlap of p-orbitals
Have d-orbitals allowing for more than 4 bonds |
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Pi bonds
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Allow for double and triple bonds
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Ion
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When element has more or fewer electrons than protons
Representative elements make ions by forming the closest noble gas electron configuration Made from metals and nonmetals |
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Cation
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Positive ion
Formed by metals Significantly smaller than neutral atom counterparts |
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Anion
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Negative ion
Formed by nonmetals Much larger than neutral atom counterpart |
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Predict ion charge based on:
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1. Atoms lose electrons from higher energy shell first
In transition metals, this means electrons are lost from s-subshell first and then from d-subshell 2. Ions are looking for symmetry Representative elements form noble gas electron configurations when they may ions Transition metals try to "even-out" their d-orbitals, so each orbital has the same number of electrons |
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Electron shielding
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1st electron shields nuclear charge from 2nd electron, so that 2nd electron doesn't feel entire nuclear charge
Instead, 2nd electron feels an effective nuclear charge |
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Effective nuclear charge (Zeff)
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Amount of charge felt by 2nd electron due to 1st electron shielding of nuclear charge
Zeff = nuclear charge (Z) - average # of electrons between nucleus and electron in question What should be plugged in to: F = Kqq/r^2 Increasing going left to right and top to bottom on periodic table |
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Periodic trends
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1. Atomic radius
2. Ionization energy 3. Electronegativity 4. Electron affinity 5. Metallic character |
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Atomic radius
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Since Zeff increases when moving left to right, each additional electron is pulled more strongly toward nucleus, resulting in a small atomic radius
Increases from top to bottom and right to left |
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Ionization energy
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Energy necessary to detach an electron from a nucleus
1st ionization energy = energy required to detach an electron from a neutral atom 2nd ionization energy = energy required to detach a 2nd electron from same atom 2nd ionization energy > 1st ionization energy because when electron is removed, Zeff on other electrons increases Increases from left to right and bottom to top of periodic table (explained by Zeff) |
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Electronegativity
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Tendency of an atom to attract an electron in a bond that it shared with another atom
Increases from left to right and bottom to top of periodic table Related to Zeff in similar way as ionization energy Undefined for noble gases |
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Electron affinity
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Willingness of an atom to accept an additional electron
Energy released when an electron is added to a gaseous atom Increases left to right and bottom to top of periodic table Related to Zeff Electron affinity is more exothermic to right and up on periodic table Endothermic for noble gases |
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SI units
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Mass = kg
Length = m Time = s Electric current = A Temperature = K Luminous intensity = cd Amount of substance = mol Mega (M) = 10^6 Kilo (k) = 10^3 Deci (d) = 10^-1 Centi (c) = 10^-2 Milli (m) = 10^-3 Micro (u) = 10^-6 Nano (n) = 10^-9 Pico (p) = 10^-12 Femto (f) = 10^-15 |
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Bonds
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What holds atoms together
2 types: 1. Covalent bonds 2. Ionic bonds 2 atoms will form a bond if they can lower their overall energy level by doing so Nature seeks lowest energy state Energy is always required to break a bond, no energy is every released by breaking a bond |
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Covalent bonds
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2 electrons are shared by 1 nuclei
Negatively charges electrons are pulled toward both positively charged nuclei by electrostatic forces |
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Bond length
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Point where energy level is lowest
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Bond dissociation energy (bond energy)
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Energy necessary to achieve a complete separation of atoms
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Compound
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Substance made from 2 or more elements
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Empirical formula
|
In pure compounds, relative number of atoms of 1 element to another can be represented by a ratio
Glucose = CH2O To find empirical formula from percent mass: Compound = 6% H & 94% O by mass Assume 100g of sample (6g H)/(1g/mol) = 6mol (94g O)/(16g/mol) = 5.9mol = 6 (must be whole #s) 6/6 = 1 Empirical formula = HO |
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Molecules
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Separate and distinct units, in molecular compounds, formed from repeated groups of atoms
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Molecular formula
|
Exact number of elemental atoms in each molecule in a molecular compound
Glucose = C6H12O6 |
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Percent mass
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Calculated from empirical formula and atomic weight of each atom
Ex: Percent mass of Carbon in Glucose (CH2O) (molecular weight C)/(molecular weight of CH2O) = 12/30 = 0.4 0.4 x 100 = 40% Glucose is 40% carbon by mass |
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Ionic compounds
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Named after their cation and anion
Put cation name in from on anion name (barium sulfate, BaSO4; sodium hydride, NaH) |
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Cation nomenclature
|
Metal cation:
1. Roman number in parentheses indicating charge [copper(I) = +1 or copper (II) = +2] 2. -ic greater charge (cupric, Cu2+) or -ous smaller charge (cuprous, Cu+) Nonmetal cation: 1. cation name ends in -ium (ammonium, NH4+) |
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Anion nomenclature
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1. -ide after anion (hydride ion, H-; hydroxide, OH-)
2. polyatomic anions with multiple oxygens end with -ite (less oxygenated) or -ate (more oxygenated) depending on relative # of oxygens (nitrite ion, NO2-; nitrate ion, NO3-) 3. More oxygens represented by hypo- (fewest oxygens) or per- (most oxygens) prefixes (hypochloride, ClO-; chlorite, ClO2-; chlorate, ClO3-; perchlorate, ClO4-) 4. If oxyanion has a hydrogen, word hydrogen is added (hydrogen carbonate ion, HCO3-) |
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Acid nomenclature
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Named based on their anions
1. If ends in -ide, name starts with hydro- and ends in -ic (hydrosulfuric acid, H2S) 2. If an oxyacid, ending -ic (more oxygen) and -ous (less oxygens) (sulfuric acid, H2SO4; sulfurous acid, H2SO3) |
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Binary molecular compounds
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Compounds with only 2 elements
Beings with name of element farthest to left and lowest in periodic table Name of 2nd element is given suffix -ide and greek # prefix is used on 1st element if necessary Ex: dinitrogen teroxide, N2O4 |
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Physical reaction
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When compound undergoes a reaction and maintains its molecular structure and this its identity
Ex: Melting, evaporation, dissolution and rotation of polarized light |
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Chemical reaction
|
When a compound undergoes a reaction and changes its molecular structure to form a new compound
Ex: Combustion, metathesis and redox Can be represented by a chemical equation with the molecular formulae of the reactants on the left and products on the right Ex: CH4 + 2O2 --> CO2 + 2H2O Coefficients indicate the relative number of molecules The atoms are always conserved, the equation is balanced |
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Reaction runs to completion
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Reaction from to the right until at least one of the reactants is depleted
Often reactions do not run to completion because they reach equilibrium first |
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Limiting reagent
|
Reactant which is depleted first if reaction runs to completion
Not necessarily the reactant of which there is least |
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Theoretical yield
|
Amount of product produced when a reaction runs to completion
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Actual yield
|
Amount of actual product after a real experiment
Reactions often don't run to completion or there are competing reactants that reduce the actual yield |
|
Percent yield
|
[(Actual yield)/(Theoretical yield)] x100
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|
Reaction types
|
1. Combination
2. Decomposition 3. Single displacement 4. Double displacement 5. Redox 6. Combustion 7. Bronsted-Lowry Acid-base 8. Lewis Acid-base |
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Combination reaction
|
A + B --> C
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Decomposition reaction
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C --> A + B
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Single Displacement reaction
|
A + BC --> B + AC
Also called single replacement |
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Double Displacement reaction
|
AB + CD --> AC + BD
Also called double replacement or metathesis |
|
Quantum mechanics
|
Elementary particles can only gain or lose energy and certain other quantities in discrete units
|
|
Principal quantum number (n)
|
First quantum number
Shell level Larger n the greater the size and energy of the electron orbital Representative elements: n for electrons in the outer most shell is given by the period in the periodic table Transition metals: n lags 1 shell behind the period Lanthanides & actinides: n lags 2 shells behind the period |
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Valence electrons
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Electrons which contribute most to an element's chemical properties
Located in outermost shell of atom Only electrons from s & p subshells are considered valence electrons |
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Azimuthal quantum number (l)
|
Second quantum number
Designate subshells: Orbital shapes such as s, p, d & f l = 0 = s subshell l = 1 = p subshell l = n-1; for each new shell (n) there exists an additional subshell |
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Magnetic quantum number (ml)
|
3rd quantum number
Designates the precise orbital of a given subshell Each subshell will have orbitals with ml from -l to +l 1st shell, n = 1, l = 0, ml = 0 n = 3, l = 2, ml = 5 (-2, -1, 0, +1, +2) |
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Electron spin quantum number (ms)
|
4th quantum number
Can values of -1/2 or +1/2 Any orbital can hold up to 2 electrons If 2 electrons occupy the same orbital, they have the same first 3 quantum numbers |
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Pauli Exclusion Principle
|
No 2 electrons in same atom can have same 4 quantum numbers
2 electrons in same orbital have identical 1st, 2nd and 3rd quantum numbers but must have opposite electron spin quantum numbers |
|
Number of orbitals within a shell
|
n^2
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|
Heisenberg Uncertainty Principle
|
Arise from dual nature of matter (wave & particle)
There exists an inherent uncertainty in the product of the position of a particle and its momemtum The uncertainty is plank's constant (Change in position) x (change in momemtum) = h The more we know about a particle's position, the less we know about its momemtum |
|
Aufbau Principle
|
With each new proton added to create a new element, a new electron is added as well
Electrons look for an orbital with the lowest energy state |
|
Electron configuration
|
For a given atom, list the shells and the subshells in order from lowest to highest energy level and add a subscript to show the number of electrons in each subshell
Na: 1s^2 2s^2 2p^6 3s^1 Ar: 1s^2 2s^2 2p^6 3s^2 3p^6 Fe: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6 Br: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^5 |
|
Abbreviated electron configuration
|
Using configuration of next smallest noble gas
Na: [Ne] 3s^1 Ar: [Ar] Fe: [Ar] 4s^2 3d^6 Br: [Ar] 4s^2 3d^10 4p^5 Cu: [Ar] 4s^1 3d^10 |
|
Ground state
|
atom whose electrons are all their lowest energy levels
|
|
Electron configuration for ions
|
Na+: 1s^2 2s^2 2p^6 or [Ne]
Fe 3+: [Ar] 3d^5 Br-: [Ar] 4s^2 3d^10 4p^6 or [Kr] Be with excited electron: 1s^2 2s^1 2p^1 |
|
Hund's Rule
|
Electons will not fill any orbital in the same subshell until all other orbitals in that subshell contain at least 1 electron
The unpaired electrons will have parallel spins |
|
Planck's Quantum Theory
|
Electromagnetic energy is quantized
Comes only in discrete units related to the wave frequency Change in E = hf h = planck's constant = 6.6e^-34 J s f = frequency E = energy Also the equation for the energy of a single photon wavelength = h/(mv) When electron falls from higher energy rung to lower energy rung, energy is released from atom in the form of a photon Opposite is also true: if electron collides with photon, they can be bumped up in energy Frequency of photon corresponds to change in energy of electron |
|
Photoelectric effect
|
One to one photon-electron collisions
Proved light was made up of particles Kinetic energy of electrons increases only when intensity is increased by increasing frequency of each photon |
|
Work function (ϕ)
|
Minimum amount of energy needed to eject an electron
KE = hf - (ϕ) KE = kinetic energy of ejected electron hf = energy of photon ϕ = work function |
|
Standard temperature and pressure (STP)
|
Temp: 0 degrees Celsius = 273 K
Pressure: 1 atm (atmosphere) moles: 1 mole |
|
Mean free path
|
Distance traveled by a gas molecule between collisions
|
|
Ideal gas
|
Based on kinetic molecular theory
1. zero volume 2. no forces other than repulsive forces due to collisions 3. elastic collisions 4. kinetic energy is proportional to temperature of gas |
|
Kinetic molecular theory
|
Ideal gas lack certain real gas characteristics
|
|
Ideal gas law
|
PV = nRT
P: pressure V: volume n: number of moles R: universal gas constant = 8.314 J/Kmol T: temperature in kelvin |
|
Charles' Law
|
Volume of gas is proportional to temperature at constant pressure
|
|
Boyle's Law
|
Volume of gas is inversely proportional to pressure a constant temperature
|
|
Avogadro's Law
|
Volume of gas is proportional to number of moles at constant temperature and pressure
|
|
Standard molar volume at STP
|
22.4 Liters
|
|
Partial pressure
|
Total pressure of gaseous mixture times mole fraction of particular gas
Pa = (Xa) (Ptotal) Pa: partial pressure of gas "a" Xa: mole fraction of gas "a" Ptotal: total pressure of gaseous mixture |
|
Mole fraction (X)
|
number of mole of gas "a" divided by total number of moles of gas in sample
|
|
Dalton's Law
|
Total pressure exerted by a gaseous mixture is the sum of the partial pressures of each of its gases
Ptotal = P1 + P2 + P3... |
|
Average translational energy of ideal gas
|
KE = (3/2) RT
KE: kinetic energy, found from RMS velocity R: universal gas constant T: temperature |
|
Graham's Law
|
Ratio of RMS velocities of 2 gases in homogeneous mixture
V1/V2 = square root (m2)/square root (m1) V: RMS velocities m: mass of gas molecules Describes effusion and diffusion |
|
Effusion
|
Spreading of gas from high pressure to low pressure through a pinhole
effusion rate 1/effusion rate 2 = square root (M2)/square root (M1) M: molecular weights |
|
Pinhole
|
opening much smaller than average distance between gas molecules
|
|
Diffusion
|
Spreading of one gas into another gas or empty space
diffusion rate 1/diffusion rate 2 = square root (M2)/square root (M1) |
|
How real gases deviate from ideal gases
|
Vreal > Videal
Intermolecular forces exist Preal < Pideal |
|
Collision model
|
In order for a reaction to occur, reacting molecules must collide
Rate of reaction is much lower than frequency of collisions Most collisions do not result in reaction |
|
Activation Energy
|
Threshold energy required for collisions to create new molecules in a reaction (Ea) - independent of temperature
Rate of reaction (K) increases with temperature |
|
Factors affecting rate of reaction
|
1. temperature
2. pressure (negligible effect) 3. concentration of reactants |
|
Elementary reaction
|
Coefficient tells you how many molecules participate in a reaction producing collision
|
|
Intermediates
|
Species that are products of one reaction and reactants of a later reaction in a reaction chain
Concentration is very low because they are often unstable and react as quickly as they are formed |
|
Rate Law
|
Rate (forward) = Kf [A]^a [B]^b
Kf: rate constant (not rate of reaction) a: order of each respective reactant b: order of each respective reactant a + b: overall order of reaction Assume no reverse reaction |
|
1st Order reaction
|
A --> product
rate = Kf[A] A decreases exponentially ln[A] vs. time graph: straight line slope: -Kf Constant half-life independent of concentration [A] No collisions take place |
|
2nd Order Reaction (1 reactant)
|
2A --> products
rate = Kf[A]^2 1/[A] vs time graph: straight line slope: Kf Half-life dependent upon concentration [A] |
|
2nd Order Reaction (2 reactants)
|
A + B --> products
Rate = Kf[A][B] no predictable half-life |
|
3rd Order Reaction
|
3A --> products
Rate = Kf[A]^3 1/2[A]^2 vs. time graph: straight line slope: Kf |
|
Zero Order Reaction
|
[A] vs. time graph: straight line
slope: -Kf |
|
Rate determining step
|
Rate of slowest elementary step determines the rate of the overall reaction
|
|
Catalyst
|
Substance that increases the rate of reaction without being consumed or permanently altered
Enhance product selectivity Reduce energy consumption Lower activation energy (Ea) Increase steric factor (p) Creates a new reaction pathway which typically includes an intermediate Cannot alter equilibrium constant (K) |
|
Heterogeneous catalyst
|
Different phase than reactants and products
Usually solids while reactants and products are liquids or gases Reaction rates can be enhanced by increasing surface area of catalyst |
|
Homogeneous catalyst
|
same phase as reactants and products, usually in gas or liquid phase
Aqueous acid or base solutions Autocatalysis = generate catalyst as product |
|
Chemical equilibrium
|
Rate of forward reaction equals rate of reverse reaction
No change in concentration of products or reactants Point of greatest entropy Reactions always move toward equilibrium, therefore Q will always change towards K |
|
Equilibrium constant (K)
|
Described by the law of mass action
Relationship between a chemical equation and the equilibrium constant aA + bB --> cC + dD K = ([C]^c [D]^d)/([A]^a [B]^b) = (products^coefficients)/(reactants^coefficients) Only depends upon temperature Equilibrium constant for a series of reaction is equal to the product of the equilibrium constants for each of its elementary steps Pure solid or liquid is K=1, therefore no included in the law of mass action equation |
|
Reaction quotient (Q)
|
For reactions not at equilibrium
Q = (products^coefficients)/(reactants^coefficients) Q is not constant, it can be any positive value Use Q to predict the direction in which a reaction will proceed Since reactions always move toward equilibrium, Q will always change towards K |
|
Comparison of Q & K
|
Q = K: reaction at equilibrium
Q > K: leftward shift because there is more product than reactant Q < K: rightward shift because there is more reactant than product |
|
Le Chatelier's Principle
|
When a system at equilibrium is stressed, the system will shift in the direction that will reduce that stress
Types of stress: 1. addition or removal of a product or reactant 2. changing the pressure of the system 3. heating or cooling the system |
|
Thermodynamics
|
Study of energy and its relationship to macroscopic properties of chemical systems
divides universe into 2: 1. surroundings 2. system |
|
System
|
macroscopic body under study
3 types: 1. open 2. closed 3. isolated |
|
Surroundings
|
everything other than macroscopic body under study (system)
|
|
Open system
|
exchange both mass and energy with surroundings
|
|
Close system
|
exchange energy but do not exchange mass with surroundings
|
|
Isolated system
|
doesn't exchange mass or energy with surroundings
|
|
Extensive properties
|
chance with amount
Examples: volume (V) and # of moles (n) describes macroscopic state of a system proportional to size of system |
|
Intensive properties
|
describes macroscopic state of a system
independent of size of system examples: Pressure (P) and temperature (T) |
|
State functions
|
properties that describe the state of a system
pathway independent state properties describe the state of a system change in a state property going from one state to another is the same regardless of the process via which the system was changed 3 state properties describe the state of a system (one being extensive property) unambiguously |
|
2 ways to transfer energy between systems
|
1. Heat (Q)
2. Work (W) |
|
Heat (Q)
|
Natural transfer of energy from a warmer body to a cooler body
3 forms: 1. conduction 2. convection 3. radiation movement of energy via conduction, convection or radiation; always from hot to cold |
|
Work (W)
|
any energy transfer between systems that is not heat
|
|
Conduction
|
thermal energy transfer via molecular collisions
requires physical contact higher energy molecules of one system transfer some of their energy to lower energy molecules of other system via molecular collisions |
|
Energy flow
|
heat flow:
change in T = IR T: temperature I: heat current (Q/t) R: resistance to heat flow Current flow: V = IR V: voltage I: electrical current R: resistance to electrical flow Fluid flow: change in P = QR P: pressure Q: heat R: resistance to flow thicker conduits = greater flow longer conduits = less flow |
|
Convection
|
thermal energy transfer via fluid movements
differences in pressure or density drive warm fluid in direction of cooler fluid examples: ocean and air currents |
|
Radiation
|
thermal energy transfer via electromagnetic waves
only type of heat that transfer though a vacuum |
|
PV work
|
system at rest, no gravitational potential energy and no kinetic energy may still be able to do PV work
at constant pressure, work is equal to pressure multiplied by change in volume W = P(change in V) [at constant pressure] No PV work is done if volume is constant PV work takes place when a gas expands against a force regardless of whether or not the pressure is constant |
|
1st Law of Thermodynamics
|
energy of system and surroundings is always conserved
any energy change to a system must equal heat flow into system + work done by system Change in E = Q + W Work done on the system is positive and work done by the system is negative |
|
2nd Law of Thermodynamics
|
Heat cannot be changed completely into work in a cyclical process
Qh = W + Qc Qh: heat entering engine W: work Qc: heat leaving engine |
|
Carnot's efficiency
|
e = 1 - (Tc/Th)
e: efficiency Tc: temperature of cold reservoir Th: temperature of hot reservoir |
|
7 state functions
|
1. internal energy (U)
2. temperature (T) 3. pressure (P) 4. volume (V) 5. enthalpy (H) 6. entropy (S) 7. Gibbs free energy (G) |
|
Internal energy (U)
|
all possible forms of energy imaginable on a molecular scale
examples: vibrational, rotational, translational, electronic, intermolecular potential and rest mass energy heat energy, thermal energy and or heat |
|
Temperature
|
thermodynamic property described by zeroth law:
2 systems in thermal equilibrium with 3rd system are in thermal equilibrium with each other any increase in thermal energy, increases temperature KE = (3/2)kT KE: kinetic energy T: temperature k: boltzmann constant (1.38e-23 J/K) measurement of how fast molecules are moving or vibrating |
|
Enthalpy (H)
|
H = U + PV
H: enthalpy (joules) U: internal energy PV: Pressure x Volume not conserved, constantly changing Enthalpy is a state function and en extensive property Depends only on temperature Change in H = (change in U) + P(change in V) [constant pressure] |
|
Standard State
|
do not confuse with STP (standard temperature and pressure)
an element at its standard state at 25 degrees Celsius is arbitrarily assigned an enthalpy value of 0 J/mol Any chosen temperature 1 bar of pressure = 750 torr = 10^5 Pascals |
|
Reference form
|
standard state for a pure solid or liquid
any chosen temperature 1 bar of pressure = 750 torr = 10^5 Pascals form that is most stable at the values |
|
Standard enthalpy of formation (change in Hf)
|
Change in enthalpy for a reaction that creates 1 mole of that compound from its raw elements in their standard state
|
|
Change in enthalpy at constant pressure
|
Change in H = Q
constant pressure, closed system at rest, PV work only H: enthalpy Q: heat |
|
Heat of reaction
|
change in enthalpy from reactants to products
Change in Hreaction = (Change in Hfproducts) - (change in Hfreactants) |
|
Hess' Law
|
Sum of enthalpy changes for each step is equal to total enthalpy change regardless of path chosen
when you add reactions, you can add their enthalpies because enthalpy is a state function |
|
Endothermic
|
if enthalpy change is positive
absorbs heat, making reaction system cold at constant P, where change in H = Q, endothermic reaction produces heat flow to system |
|
Exothermic
|
if enthalpy change is negative
release heat, making reaction system hot at constant P, where change in H = Q, exothermic reaction produces heat flow to surroundings |
|
Activation energy
|
initial increase in energy
|
|
Transition state
|
peal of energy hill represents molecules in transition state
old bonds are breaking and new bonds are forming occurs during reaction collision do not confuse with intermediates (products of 1st step in 2 step reaction) |
|
Catalyst
|
lowers activation energy of forward and reverse reactions
equilibrium and enthalpy change is unaffected Affects the rate of reaction |
|
Entropy (S)
|
Nature's tendency to create the most probably situation that can occur within a system
nature's tendency toward disorder state function, which means that entropy change of forward reaction is equal to negative entropy change of reverse reaction |
|
2nd law of thermodynamics
|
entropy of an isolated system will never decrease without some outside intervention (work)
|
|
Entropy change of universe
|
Change Ssystem + change Ssurroundings = Change Suniverse > or = 0
Sum of entropy changes of any system and its surroundings equals entropy change of universe, which is always equal to or greater than zero (positive) entropy of system can decrease, only if, at the same time, entropy of surroundings increases by greater or equal magnitude |
|
Reversible
|
Only ideal reactions are reversible because only ideal reactions create zero change in entropy of universe
on a microscopic scale, all real chemical reactions are reversible |
|
Irreversible
|
all reactions are irreversible on a macroscopic scale
|
|
what drives the direction of a reaction
|
it is entropy, not energy, that drives the direction of a reaction
entropy is nature's effort to spread energy evenly between system, from high to low entropy increases with number, volume and temperature reaction must increase the entropy of the universe in order to proceed |
|
Equilibrium
|
point in reaction where universe has achieved maximum entropy
|
|
3rd law of thermodynamics
|
assigns by convention a zero entropy value to any pure substance (either element or compound) at absolute zero and in internal equilibrium
|
|
Gibbs Free energy (G)
|
change G = (change H) = T(change S)
G: gibbs free energy H: enthalpy T: temperature (constant) S: entropy All variables refer to the system and not surroundings only good for constant T & P reactions negative change G = spontaneous reactions extensive property and state function not conserved, an isolated system can change its gibbs free energy Maximum non-PV work available from a reaction positive change in H and negative change in S, can never equal negative change in G, which means can never be spontaneous negative change in H and positive change in S = negative change in G = spontaneous if both change in H and S are same sign, than change in G depends on T |
|
Solution
|
homogeneous mixture of 2 or more compounds in a single phase (solid, liquid, gas)
|
|
Solvent
|
compound of which there is more in a solution
compound that predominates |
|
Solute
|
compound of which there is less in a solution
|
|
Ideal solution
|
solutions made from compounds that have similar properties
compounds can be interchanged within solution without changing spatial arrangements of molecules or intermolecular attractions ex: benzene in toluene (similar bonding properties and size) |
|
Ideally dilute solution
|
solute molecules are completely separated by solvent molecules so that they have no interaction with each other
mole fraction of solvent is ~1 |
|
Nonideal solution
|
violate both ideal solution and ideally dilute solution conditions
|
|
Colloid
|
larger particles that form mixture with solvents
as long as gravity doesn't cause particles to settle out of mixture over time larger than solute particles can be any combination of phases, except gas in gas ex: aerosol, foam, emulsion or sol too small to be extracted by filtration but large or charged enough to be separated by semipermeable membrane |
|
Colloid solution
|
scatter light, unlike true solution
can be attracted or repelled by dispersion medium (solvent) |
|
Disolution
|
when solute is mixed in a solvent
like dissolve like because of london dispersion forces polar solvents dissolve polar solutes nonpolar solvents dissolve nonpolar solutes |
|
Solvation
|
ionic compounds are dissolved by polar substances
Ionic compounds dissolve into anions and cations and are surrounded by oppositely charged ends of polar solvent water is a good solvent for ionic compounds (process called hydration) |
|
Aqueous phase
|
something that has been hydrated (solvated in water)
water is a poor conductor of electricity unless it contains electrolytes (compounds that form ions in aqueous solution) |
|
Polyatomic ions
|
1. nitrite, NO2 -
2. nitrate, NO3 - 3. sulfite, SO3 -2 4. sulfate, SO4 -2 5. hypochlorite, ClO - 6. chlorite, ClO2 - 7. chlorate, ClO3 - 8. perchlorate, ClO4 - 9. carbonate, CO3 -2 10. bicarbonate, HCO3 - 11. phosphate, PO4 -3 |
|
Electrolyte
|
compound which forms ions in aqueous solution and therefore solution can conduct electricity
|
|
5 ways to measure concentration in solution
|
1. molarity (M)
2. molality (m) 3. mole fraction (X) 4. mass percentage 5. parts per million (ppm) |
|
Molarity (M)
|
moles of a compound divided by volume of solution
mol/L M = moles solute/volume solution |
|
Molality (m)
|
moles of solute divided by kilograms of solvent
mol/kg m = moles solute/kg solvent |
|
mole fraction (X)
|
moles of compound divided by total moles of all species in solution
no units, since it is a ratio X = moles solute/total moles of all solutes and solvents |
|
mass percentage
|
100 times ratio of mass of solute to total mass of solution
mass % = (mass solute/total mass solution) x 100 |
|
Part per million (ppm)
|
10^6 times ratio of mass solute to total mass of solution
ppm = (mass solute)/(total mass solution) x 10^6 |
|
3 steps of solution formation
|
1. breaking of intermolecular bonds between solute molecules (endothermic)
2. breaking of intermolecular bonds between solvent molecules (endothermic) 3. formation of intermolecular bonds between solvent and solute molecules (exothermic) if overall reaction releases energy (heat) than new bonds are more stable (stronger) than old bonds energy is required to break bonds increase in entropy (S) because more disorder |
|
heat of solution
|
overall change in energy of solution (enthalpy)
Change H(sol) = change H1 + change H2 + change H3 neg = stronger, more stable bonds pos = weaker, less stable bonds |
|
vapor pressure
|
pressure created by molecules in open space
pressure necessary to bring liquid and gas phases of a compound to equilibrium increases with pure liquids, decreases with solutions increases with temperature because related to KE |
|
Boiling point
|
temperature at which vapor pressure of liquid equals Patm
|
|
melting point
|
temperature at which vapor pressure of solid equals vapor pressure of liquid
|
|
nonvolatile solute
|
solute with no vapor pressure
when added to liquid, decreases vapor pressure |
|
Raoult's law
|
vapor pressure of solution is proportional to mole fraction of liquid and vapor pressure of pure liquid
Pv = (Xa)(Pa) Pv: vapor pressure of solution Xa: mole fraction of liquid a Pa: vapor pressure of pure liquid a |
|
Raoult's law for nonvolatile solutes
|
if 97% of solution is solvent, then vapor pressure will be 97% of vapor pressure of pure solvent
|
|
Raoult's law of volatile solutes
|
if 97% of solution is solvent, the vapor pressure will be 97% of vapor pressure of pure solvent + 3% of vapor pressure of pure solute
|
|
Volatile solute
|
solute with vapor pressure
total pressure of solution is the sum of partial pressures Pv = XaPa + XbPb Pv: total vapor pressure XP: partial pressure of solvent |
|
vapor pressure
|
negative heats of solution = stronger bonds = lower vapor pressure
positive heats of solution = weaker bonds = higher vapor pressure |
|
solubility
|
solute's tendency to dissolve in solvent
solute: salt solvent: water not the same as solubility product max # of moles of solute that can dissolve in solution depends on temperature and ions in solution |
|
precipitation
|
reverse reaction of dissolution
takes place initially at a slower rate than dissolution, as salt dissolves and concentration of dissolved salt builds, rate of dissolution and precipitation equilibrate |
|
Saturate
|
rate of dissolution and precipitation are equal
concentration of dissolved salt has reached max in saturated solution |
|
Solubility product (Ksp)
|
equilibrium constant of equilibrium of solvation reaction
pure solids and liquids are excluded from equilibrium expression because have mole fraction of 1 Ksp = [products]^coefficient/ [reactants]^coefficient changes only with temperature |
|
Spectator ions
|
ions that have no effect on equilibrium because not in equilibrium expression
|
|
Common ion effect
|
ions that affect equilibrium because in equilibrium equation
ion involved is common to ion in equilibrium expression will push equilibrium in direction which tends to reduce concentration of that ion Common ion added to saturated solution will shift equilibrium increasing precipitate, but does not affect Ksp common ion added to unsaturated solution will not shift equilibrium because in unsaturated solution there is no equilibrium to shift |
|
Solubility guidelines
|
compounds with water solubilities of less than 0.01 mol/L = insoluble
ionic compounds containing: nitrate (NO3 -), ammonium (NH4 +) and alkali metals (Li+, Na+, K+) = soluble ionic compounds containing halogens (Cl-, Br-, I-) = soluble except: Ag+, Hg2 +2, Pb +2 = insoluble Sulfate compounds = soluble except: Hg2 +2, Pb +2, Ca +2, Sr +2, Ba +2 = insoluble When paired with sulfides (S -2) and hydroxides (OH-), compounds containing heavier alkaline metals = soluble: Ca +2, Sr +2, Ba +2 Carbonates (CO3 -2), phosphates (PO4 -3), sulfides (S -2) and hydroxides (OH-) = insoluble solubility of most salts increases with temperature solubility of gases decreases with temperature |
|
Henry's Law
|
demonstrates that solubility of gas is proportional to its vapor partial pressure
as pressure decreases, solubility of gas decreases (ex: can of soda) C = ka1Pv or Pv = Xaka2 C: solubility of gas a (mol/L) ka1: henry's law constant Pv: vapor partial pressure of gas a above the solution |
|
Heat Capacity (C)
|
measure of energy change needed to change temperature of substance
C = Q/(change T) Q = C(change T) constant pressure heat capacities are greater than constant volume heat capacities always positive, temperature will always increase when added to a substance at constant volume or pressure heat capacity does not change with temperature as amount of energy a substance can absorb per unit of temperature change |
|
Specific heat capacity (c)
|
heat capacity per unit mass
Q = mc(change in T) m: mass c: specific heat capacity T: temperature c(water) = 1 cal/gC (definition of 1 calorie) |
|
Coffee Cup calorimeter
|
measures energy change
constant pressure calorimeter because it measures energy change at Patm 1. 2 coffee cups are used to insulate the solution 2. a stirrer maintains equal distribution of energy 3. thermometer measures change in temperature cannot contain expanding gases reactions occur at constant pressure of local atmosphere used to measure heats of reaction at constant P, Q = change H |
|
Bomb calorimeter
|
measures energy change at constant volume
tells us the internal energy change in a reaction at constant volume, Q = change U 1. steel container full of reactants is placed inside another rigid, thermally insulated container 2. when reaction occurs, heat is transferred to surrounding water |
|
Normal melting point
|
melting point at constant pressure of 1atm, which is 0 degrees Celsius
|
|
Normal boiling point
|
boiling point at constant pressure of 1atm, which is 100 degree Celsius
|
|
Heat of fusion
|
enthalpy (H) change associated with melting
since pressure is constant, heat equals enthalpy change (Q = change H) exactly same amount of heat is absorbed during melting as is released during freezing |
|
Heat of vaporization
|
enthalpy (H) change associated with boiling
exactly same amount of heat is absorbed during vaporization as is released during condensation and sublimation/deposition |
|
phase change
|
no change in temperature until all molecules have undergone phase change
represented by flat line segments of heating curve 1. melting-freezing 2. vaporization-condensation 3. sublimation-deposition each phase of a substance has its own specific heat |
|
evaporation
|
occurs when partial pressure above liquid is less than liquid's vapor pressure, but atmospheric pressure is greater than vapor pressure
under these conditions, liquid evaporates rather than boils |
|
phase diagram
|
indicates phases of substance at different pressures and temperatures
1. each section represents a different phase 2. boundary lines represent temperatures and pressures where corresponding phases are in equilibrium |
|
triple point
|
point at which substance exists in equilibrium as solid, liquid and gas
|
|
critical temperature
|
temperature above which substance cannot be liquefied regardless of pressure applied
|
|
critical pressure
|
pressure required to produce liquefaction while substance is at critical temperature
|
|
critical point
|
where critical pressure and critical temperature intersect
fluid beyond this point has characteristics of both gas and liquid (supercritical fluid) |
|
colligative properties
|
depend solely on # of particles, irrespective of type of particle
4 of solution: 1. vapor pressure 2. boiling point 3. freezing point 4. osmotic pressure depend on number, not kind |
|
boiling point elevation
|
equation for ideally dilute solutions
solute addition increases boiling point due to addition of nonvolatile solute, cannot be applied to volatile solutes substance boils when its vapor pressure equals the local atmospheric pressure change T = kbmi T: temperature kb: specific constant of substance being boiled m: molality (mol/L) because molality doesn't change with temperature while molarity does i: van't Hoff factor consider # particles after dissociation |
|
Van't Hoff Factor (i)
|
number of particles into which a single solute particle will dissociate when added to solution
|
|
freezing point depression
|
equation for ideally dilute solution
impurities (solute) interrupt crystal lattice and lower freezing point only nonvolatile solutes change T = kfmi T: temperature kf: specific constant substance being frozen m: molality (mol/L) i: van't hoff factor (# particles dissociated) |
|
osmotic pressure
|
measure of tendency of water (or some solvent) to move into solution via osmosis
osmotic pressure = iMRT M: molarity R: resistance T: temperature i: van't hoff factor (# dissociated particles) only relevant when comparing 1 solution with another pressure pulling into a solution |
|
hydrostatic pressure
|
pressure pushing out of a solution
|
|
Arrhenius acid
|
anything that produces H ions in aqueous solution
|
|
Arrhenius base
|
anything that produces OH ions in aqueous solution
|
|
Bronsted and lowry acids
|
anything that donates a proton
|
|
Bronsted and lowry bases
|
anything that accepts a proton
|
|
Lewis acid
|
anything that accepts a pair or electrons
includes bronsted-lowry acids molecules that have incomplete octet of electrons around central atom (AlCl3 or BF3) simple cations expect alkali and heavier alkali earth metal cations smaller the cation and higher the charge, stronger the acid strength ex: Fe +3 |
|
lewis base
|
anything that donates a pair of electrons
|
|
pH
|
measure of H ion concentration (mol/L)
pH = -log[H+] scale runs from 0 to 14, each point on pH scale corresponds to 10X difference in H ion concentration acid at pH 2 produces 10X as many H ions as acid at pH 3 and 100X as many H ions as acid at pH 4 at 25 degrees: pH of 7 is neutral, lower pH is acidic and higher pH is basic |
|
estimating pH
|
10^0 = 1
10^1 = 10 10^x = 3.16 x = log(3.16) x = between 0 & 1 = 0.5 log(10^-3) = -3 -log(10^-3) = 3 log(4e-3) = between pH 2 & pH 3 = 2.4 log(AB) = log(A) + log(B) |
|
acid-base reaction
|
HA + H2O --> A- + H3O+
HA: acid H2O: base A-: conjugate base H3O+: conjugate acid stronger the acid, weaker its conjugate base Kw = KaKb stronger the base, weaker its conjugated acid weak acid may have a strong or weak conjugate base acids taste sour or tart bases taste bitter and are slippery when wet rate of reactions in living cells involving transfer of proton depends upon concentration of H+ ions or pH |
|
amphoteric
|
some substances act as acid or base depending on environment
ex: water |
|
strong acids
|
acid that is stronger than H3O+
completely dissociates in water 1. hydroiodic acid, HI 2. hydrobromic acid, HBr 3. hydrochloric acid, HCl 4. nitric acid, HNO3 5. perchloric acid, HClO4 6. chloric acid, HClO3 7. sulfuric acid, H2SO4 |
|
Strong bases
|
base that is stronger than OH-
completely dissociates in water 1. sodium hydroxide, NaOH 2. potassium hydroxide, KOH 3. amide ion, NH2- 4. hydride ion, H- 5. calcium hydroxide, Ca(OH)2 6. sodium oxide, Na2O 7. calcium oxide, CaO |
|
hydronium ion
|
H3O+
simply hydrated proton in acid-base reactions, H3O+ and H+ are same thing |
|
polyprotic acid
|
acids that can donate more than one proton
2nd proton donated by acid is very weak and its effect on pH is negligible (can be ignored) |
|
diprotic acid
|
acids that can donate only 2 protons
is a polyprotic acid as well |
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dissociation
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acid dissociation decreases with acid concentration but acid strength increase with acid concentration
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Acid-Base strengths
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pH 14: NaOH (strong base)
pH 11-12: NH3 (base) pH 10-11: HCO3- (base) pH 8-9: F- (weak base) pH 7: H2O (neutral) pH 6-7: H2CO3 (weak acid) pH 4-5: NH4+ (acid) pH 2-3: HC2H3O2 (acid) pH 1-2: HSO4- (acid) pH 0-1: HF (acid) pH 0: HCl (strong acid) |
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3 factors determined whether a H containing molecule will act as acid (release H into solution)
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1. strength of bond holding H to molecule
2. polarity of bond 3. stability of conjugate base |
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hydrogen halide acid strength
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weakest to strongest:
1. HF 2. HCl 3. HBr 4. HI as acidity increases, polarity decreases and bond strength decreases |
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oxyacid acid strength
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weakest to strongest:
1. HClO 2. HClO2 3. HClO3 4. HClO4 more O, means stronger acid acidity increases with oxidation number of central atom |
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hydrides
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can be basic, neutral or acidic
binary compounds (only 2 elements) containing H basic: left periodic table (NaH) acidic: right periodic table (H2S) metal hydrides are either basic or neutral nonmetal hydrides are acidic or neutral (NH3, ammonia, is an exception) acidity of nonmetal hydrides increases down periodic table: H2O < H2S < H2Se < H2Te CH4: neutral NH3: weak base H2O: neutral HF: weak acid SiH4: neutral PH3: weak base H2S: weak acid HCl: strong acid |
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autoionization of water
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H2O + H2O --> H3O+ + OH-
Kw = equilibrium constant for this reaction Kw = [H+][OH-] at 25 degrees: Kw = 10^14 (lies far to left) [H+] = [OH-] = 10^-7 mol/L = pH 7 pH + pOH = pKw = 14 |
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acid dissociation constant (Ka)
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acid has its own equilibrium constant in water
HA + H2O --> H3O+ + A- Ka = [H+][A-]/[HA] |
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base dissociation constant (Kb)
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for every Ka there is a Kb
equilibrium constant for reaction of conjugate base (A-) with water A- + H2O --> OH- + HA not reverse reaction of Ka Kb = [OH-][HA]/[A-] |
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product of Ka & Kb
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Ka X Kb = [H+][A-]/[HA] X [OH-][HA]/[A-] = [H+][OH-] = Kw
KaKb = Kw pKa + pKb = pKw = 14 larger the Ka, smaller the pKa and stronger the acid Ka greater than 1 or pKa less than 0 indicates a strong acid |
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Finding pH of strong acids and bases
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strong acids and bases dissociate completely in water, which means their concentration will = 0, thus there is no Ka or Kb
therefore, concentration of H+ or OH- = concentration of acid or base strong acid ex: 0.01 M HCl solution = 0.01 mol/L H+ ions 0.01 = 10^-2 -log(10^-2) = 2 = pH 2 strong base ex: 0.01 M NaOH solution = 0.01 mol/L OH- ions pOH = 2 = pH 12 |
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Finding pH of weak acids and bases
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weak acid ex:
0.01 M HCN solution HCN + H2O --> H3O+ + CN- Ka = [H+][CN-]/[HCN] = 6.2e-10 [x][x]/[0.01 - x] = 6.2e-10 assume x is smaller than 0.01 [x][x]/[0.01] = 6.2e-10 [x] = 2.5e-6 pH = 5 - 6 = 5.6 weak base ex: process is same as weak acid, except Kb value is used and we arrive at pOH remember to subtract pOH from 14 to find pH |
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Salts
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ionic compounds that dissociate in water
create acidic or basic conditions when they dissociate pH can be predicted by comparing conjugates of respective ions strong acids = weak conjugate bases strong bases = weak conjugate acids NaOH base = conjugate acid Na+ HCl acid = conjugate base Cl- NaCl = neutral all cations, except alkali metals ad heavier alkaline earth metals cations, act as weak lewis acids in aq solutions |
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titration
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drop-by-drop mixing of acid and base
performed in order to find concentration of an unknown by comparing it with concentration of titrant |
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titrant
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known substance added drop-by-drop into unknown to find concentration of unknown
can be acidic or basic |
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titration curve
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graphic representation changing pH of unknown as acidic or basic titrant is added (volume)
sigmoidal curve |
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equivalence point
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also known as stoichiometric point
midpoint of portion of titration curve that most nearly approximates a vertical line point in titration when there are equal equivalents of acid and base in solution not necessarily when they are at equal volumes for equally strong acid-base titrations, equivalence point will be at pH 7 (only for monoprotic acids) |
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equivalent
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mass of acid or base necessary to produce or consume 1 mole of protons
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half-equivalence point
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midpoint of portion of titration curve that most nearly approximates a horizontal line
1/2 of acid has been neutralized by base concentration of acid is equal to concentration of conjugate base point at which most acid or base could be added with least change in pH, such a solution is considered buffered point in titration where solution is most well buffered pH = pKa of acid |
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Henderson-Hasselbalch equation
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pH = pKa + log([A-]/[HA])
form of equilibrium expression for Ka when [A-] = [HA], pH = pKa which is at the half-equivalence point |
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Finding pH at half-equivalence point
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Ka = [H+][A-]/[HA] = [H+] [A-]/[HA]
-log(Ka) = -log[H+] + -log([A-]/[HA]) pKa = pH - log[A-]/[HA] pH = pKa + log[A-]/[HA] |
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Finding pH at equivalence point
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Kb = Kw/Ka
Kb = [OH-][HA]/[A-] solve for OH- concentration Find pOH Subtract pOH from 14 to find pH |
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Titration curve of weak acid or base
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equivalence point not as predictable
still sigmoidal graph if base is stronger than acid, equivalence point will be above pH 7 if acid is stronger than base, equivalence point will be below pH 7 |
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indicator
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chemical used to find find equivalence point
usually a weak acid whose conjugate base is a different color new form of indicator must reach 1/10 concentration of original form pH of color change depends upon direction of titration |
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range
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pH values of 2 points of color change give range of indicator
can be predicted by using: pH = pKa + log[In-]/[HIn lower range color: pH = pKa + log(1/10) --> pH = pKa - 1 upper range color: pH = pKa + log(10/1) --> pH = pKa +1 |
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endpoint
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point where indicator changes color
equivalence point does not equal endpoint |
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Polyprotic titrations
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will have more than 1 equivalence point and more than 1 half-equivalence point
1st proton completely dissociates before 2nd proton begins to dissociate |
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Redox reaction
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oxidation-reduction reaction
electrons are transferred form one atom to another atom that loses electrons is oxidized atom that gains electrons is reduced |
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Oxidized
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atom loses electrons in redox reaction
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reduced
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atom gains electrons in redox reaction
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oxidation states
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possible charge values that an atom may hold within a molecule
necessary for redox reactions must add up to charge on molecule or ion |
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oxidation state = 0
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atoms in their elemental form
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oxidation state of fluorine (F)
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equals -1
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oxidation state of hydrogen (H)
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equals +1
except when bonded to a metal; then -1 |
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oxidation state of oxygen (O)
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equals -2
except when it is in a peroxide like H2O2 |
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Oxidation state = +1
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group 1 elements
alkali metals |
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oxidation state = +2
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group 2 elements
alkaline earth metals |
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oxidation state = -3
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group 15 elements
nitrogen family |
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oxidation state = -2
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group 16 elements
oxygen family |
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oxidation state = -1
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group 17 elements
halogens |
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LEO the lion says GER
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LEO: Lose Electrons Oxidation
GER: Gain Electrons Reduction |
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Reducing agent
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reductant
giving electrons to an atom losing electrons, is oxidized |
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Oxidizing agent
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compound containing atom being reduced
gains electrons, is reduced oxidizes other atom |
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Electric potential (E)
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associated with any redox reaction
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Half-reaction
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Each component of a redox reaction
oxidation half-reaction potential is opposite reduction half-reaction potential usually listed as reduction potentials (sign is reversed for oxidation potential) |
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Half-reaction of standard hydrogen electrode
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2H+ + 2e- --> H2
Half-reaction potential = 0.00V |
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Balancing redox reactions in acidic solutions
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1. divide reaction into its corresponding half reactions
2. balance the elements other than H and O 3. Add H20 to one side until O atoms are balanced 4. Add H+ to one side until H atoms are balanced 5. Add e- to one side until charge is balanced 6. multiply each half reaction by an integer so that an equal number of electrons are transferred in each reaction 7. add the 2 half reactions and simplify |
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Balancing redox reactions in basic solutions
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same steps as acidic solutions
neutralize H+ ions by adding same number of OH- ions to both sides of reaction |
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Galvanic cell (voltaic cell)
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uses electric potential between phases to generate a current of electrons from one phase to another in a conversion of chemical energy to electrical energy
turns chemical energy into electrical energy |
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Salt bridge
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ionic conducting phase
electrolyte solution phase impermeable to electrons type of liquid junction that minimizes potential difference between different solutions carries current in form of ions |
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Terminals
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electronic conductors such as metal wires (T)
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Electrodes
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electronic conductors (E)
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Ionic conductor
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salt bridge (I)
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Simple galvanic cell
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T-E-I-E'-T'
has 2 electrodes: anode (-) and cathode (+) oxidation half reaction takes place at anode reduction half reaction takes place at cathode 2 terminals of cell is made from same material |
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cell potential (E)
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electromotive force (emf)
potential difference between the terminals when they are not connected connecting the terminals reduces the potential difference due to internal resistance within the cell drop in emf increases, as current increases current flows in direction opposite electron flow electrons flow from anode to cathode always positive, always has chemical energy than can be converted to work |
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RED CAT
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REDuction CAThode
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AN OX
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ANode OXidation
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Positive cell potential
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equals a negative Gibbs free energy (deltaG), which equals a spontaneous reaction (work is being done by system and not on system)
deltaG = -nFEmax deltaG: Gibbs free energy n: number of moles of electrons that are transferred in balanced redox reaction F: Faraday's constant E: voltage Free energy (deltaG) represents the product of total charge (nF) times voltage (E) |
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Reactions that do not occur at standard state
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deltaG = deltaG(not) + RT[ln(Q)]
deltaG: Gibbs free energy deltaG(not): Gibbs free energy (standard conditions) T: temperature Q: reaction quotient |
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reactions that are at equilibrium conditions
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at equilibrium, there is no available free energy with which to do work; deltaG = 0
deltaG(not) = -RT[ln(K)] |
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relationship between K and deltaG(not)
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if K = 1, then deltaG(not) = 0
if K > 1, then deltaG(not) < 0 if K < 1, then deltaG(not) > 0 |
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Concentration cell
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limited form of a galvanic cell with a reduction half reaction taking place in 1 half cell and the exact reverse of that half reaction taking place in the other half cell
type of galvanic cell it is never at standard conditions, so Nerst equation is required to solve for cell potential if concentrations were equal on both sides, the concentration cell potential would be zero |
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galvanic cell
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positive cell potential
spontaneous |
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electrolytic cell
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negative cell potential
forced by outside power source to run backwards cathode is negative anode is positive RED CAT & AN OX still the same |