Aaa Mcat Z Comprehensive Mcat General Chemistry

Front 1

Metric Abbreviations: Mega, Kilo, Hecto, Deka, Deci, Centi, Milli, Nano

Back 1

Mega=M; Kilo=k; Hecto=h; Deka =da; Deci =d; Centi =c; Milli=m; Nano=n

Front 2

Conversion Factor

Back 2

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)

Front 3

Use conversion factors to solve: 1) How many kilograms in 2000 g? 2) How many feet in 60"?

Back 3

1) (2000g)(1kg/1000g) = 2kg; 2) (60")(1'/12") = 5 ft

Front 4

Significant digits

Back 4

Digits which are measured. All non-zero digits are significant. Zeros are significant unless they are placeholders

Front 5

Scientific notation

Back 5

A number expressed as Ax10B. "A" is between 1.00 and 9.99 and "B" is an integer.

Front 6

Write conversion forumlas: 1) Celsius to Fahrenheit; 2) Celsius to Kelvin

Back 6

1) ˚f = (1.8 x ˚celsius) + 32; 2) Kelvin = ˚celsius +273

Front 7

Energy

Back 7

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)

Front 8

Matter

Back 8

A substance that occupies space has mass.

Front 9

Compare Weight vs. Mass

Back 9

Mass is the amount of matter. Weight measures gravitational force. Mass never varies. Weight can vary.

Front 10

Compare potential energy and kinetic energy

Back 10

Potential energy is energy due to position. KE is energy of motion.

Front 11

Explain how to convert a number greater than 1 to scientific notation

Back 11

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

Front 12

Explain how to convert a number less than 1 to scientific notation

Back 12

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

Front 13

Explain the rules for multiplying the numbers in scientific notation

Back 13

Multiply the first numbers and add the exponents. (3x10^5)(2x10^3)= 6x10^8

Front 14

Explain the rules for dividing the numbers in scientific notation

Back 14

Divide the first numbers and subtract the exponents. (8x10^6)/(2x10^10) = 4x10^-4

Front 15

Physical properties

Back 15

Properties which can be observed without changing the substance into something different. Color, odor, hardness, density, luster, state, conductivity, solubility, boiling and melting points

Front 16

Chemical properties

Back 16

A chemical property is observed when a substance changes into a new substance. Iron forms rust in air & water; gasoline burns in oxygen

Front 17

Physical change

Back 17

Substance changes form or state only. Boiling, melting, freezing, dissolving, grinding, cutting

Front 18

Chemical change

Back 18

Where new substances are formed with new chemical and physical properties. Oxygen & hydrogen form water; sodium & chlorine form sale (sodium chloride)

Front 19

Density: Write the general equation and three standard units

Back 19

The mass of a unit volume of a substance. Density = mass/volume = g/mL; g/L; kg/L

Front 20

Element. List some examples

Back 20

A substance is composed of identical atoms. Gold, silver, oxygen, hydrogen, lead, chlorine, helium, iron, copper, fluorine, arsenic

Front 21

Compound. List some examples

Back 21

substance composed of two or more elements chemically combined. Water - H2O; Salt - NaCl; Sugar - C6H12O6; Ammonia - NH3

Front 22

Mixture. List some examples.

Back 22

A combination of substances held together by physical means (dirt, milk, soup, saltwater, granite)

Front 23

Homogeneous and Heterogeneous mixtures. Provide examples.

Back 23

Homogeneous mixtures are uniform in composition (air, metal alloy, salt water). Heterogeneous mixtures are not uniform in composition (dirt, spaghetti sauce)

Front 24

Three postulates of Dalton's Atomic Theory

Back 24

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

Front 25

The Law of Conservation of Mass

Back 25

During a chemical reaction, matter is neither created nor destroyed

Front 26

The Law of Constant Composition

Back 26

A compound always contains the same elements combined in the same proportions by mass (H2O) is 88% oxygen no matter where it is found)

Front 27

Law of Multiple Proportions

Back 27

The same elements may combine to form more than one compound. The ratios of atoms are in small whole numbers (H2O and H2O2)

Front 28

Atomic Mass Unit

Back 28

the mass of a proton or neutron is equal to 1 atomic mass unit. Symbol - "amu"; 1 amu=1.66x10^-24

Front 29

Atomic Number: What are the atomic numbers of helium, hydrogen, carbon, oxygen?

Back 29

The number of protons in the nucleus of an atom of an element. Helium-2; carbon-6; hydrogen-1; oxygen-8

Front 30

Mass number

Back 30

The sum of protons plus neutrons n the nucleus of an atom

Front 31

Isotope

Back 31

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)

Front 32

Molecule

Back 32

A group of two or more atoms held together by chemical bonds

Front 33

Ion (provide examples)

Back 33

An atom or group of atoms which contains a positive or negative electrical charge (ex. Na+; Cl-; SO4^2-)

Front 34

Cation and Anion (provide examples)

Back 34

cation - positively charged ion (Na+; Fe+2; NH4+; Ag+); anion - negatively charged ion (Cl-; SO4-2; OH-; P-3)

Front 35

Valence electroncs

Back 35

The electrons found in the outermost energy level of an atom

Front 36

Oxidation number

Back 36

A number (positive or negative) representing the charge on an ion or atom involved in a chemical bond

Front 37

Three general rules for determining oxidation numbers

Back 37

1) Atoms of uncombined elements equal 0; 2) Hydrogen = +1 (in metallic hydrides =-1); 3) Oxygen = -2 (in peroxides =-1); (bonded with fluorine =+2)

Front 38

Empirical Formula. What are the molecular and empirical formulas of peroxide?

Back 38

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)

Front 39

Molecular Formula. What is the molecular formula for ammonia?

Back 39

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)

Front 40

Positive Ion: which elements tend to form them?

Back 40

Metals tend to form positive ions by losing electrons (Na → Na+ e-)

Front 41

Negative Ion: which elements tend to form them?

Back 41

Non-metals form negative ions by gaining electrons (Cl +e- →Cl-)

Front 42

Write oxidation numbers for ions of: Group IA & IIA; Group VIA & VIIA

Back 42

IA→+1; IIA→+2; VIA→-2; VIIA→-1

Front 43

Write formula for compounds of: 1) Sodium & Sulfate; 2) Magnesium & Nitrite; 3) Aluminum & Phosphate

Back 43

1) Na2SO4; 2) Mg(NO2)2; 3) AlPO4

Front 44

Names of the ionic compounds: 1) FeCL3; 2) FeO; 3) Cu(OH); 3) Cu3PO4

Back 44

1) iron (III) chloride; 2) iron (II) oxide; 3) copper (II) hydroxide; 4) copper (I) phosphate

Front 45

10 prefixes used to name covalent compounds

Back 45

Mono=1; Di=2; Tri=3; Tetra=4; Penta=5; Hexa=6; Hepta=7; Octa=8; Nona=9; Deca=10

Front 46

Forumlas and names of acids formed from: 1) F; 2) Cl; 3) Br; 4) I

Back 46

1) HF - hydrofluoric acid; 2) HCl - hydrochloric acid; 3) HBr - hydrobromic acid; 4) HI - hydriodic acid

Front 47

Names of acids: 1)H2SO4; 2) HNO2; 3) H3PO4; 4)HClO

Back 47

1) sulfuric acid; 2) nitrous acid; 3) phosphoric acid; 4) hypochlorous acid

Front 48

Write the equation for Percent Composition. What is the percent composition of Ca in CA(OH)2?

Back 48

%=(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

Front 49

Atomic Mass

Back 49

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)

Front 50

Write the equation to calculate the atomic mass of an element

Back 50

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

Front 51

Molecular Mass

Back 51

Found by adding all the atomic masses of an element. Ex. H2O. Molecular mass = 18; H2 = 2(1) O = 16

Front 52

Avogadro's Number

Back 52

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)

Front 53

Mole

Back 53

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)

Front 54

For any substance, write a general formula to convert from Moles to Grams

Back 54

g=(n)(MM) where, n = moles MM = molecular mass g = grams

Front 55

For any substance, write a general formula to convert from Grams to Moles

Back 55

n = g/MM; n = moles; MM = molecular mass; g = grams

Front 56

For any substance, write a general formula to convert from Moles to Number of Particles

Back 56

(n)(6.02x10^23) = P; n = number of moles; P = number of particles

Front 57

For any substance, write a general formula to convert from Number of Particles to Moles

Back 57

n = P/6.02x1023; n = number of moles; P = number of particles

Front 58

Reactants; Products

Back 58

1) The starting materials in a chemical reaction; 2) The substances formed in a chemical reaction

Front 59

Balance Equations: H2 + N2 ↔ NH3; NaCL + Br2 ↔ NaBr + Cl2

Back 59

1) 3H2 + N2 ↔2NH3; 2) 2NaCl + Br2 ↔2NaBr + Cl2

Front 60

Diatomic; list 7 diatomic elements found in nature

Back 60

A molecule composed of two atoms. H2, N2, O2, Cl2, F2, Br2, F2, I2

Front 61

Synthesis or combination

Back 61

A reaction where 2 or more elements form a compound; N2 + 3H2 ↔2NH3; 2H2 + O2 ↔ 2H2O

Front 62

Decomposition

Back 62

A reaction where a compound breaks down into elements; CO2 ↔ C + O2; 2CaO ↔2Ca + O2

Front 63

Single Replacement

Back 63

A reaction involving the replacement ina compound of an element by another element. Zn + CuCl2 ↔ Cu + ZnCl2

Front 64

Double replacement

Back 64

A reaction where two ionic substances "trade" anions; NaCl + AgNO3 ↔ NaNO3 + AgCl

Front 65

Period

Back 65

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.

Front 66

Column or Family

Back 66

A verticle group of elements on the periodic table in the same column. They have similar properties and the same number of valence electrons.

Front 67

Metalloid. List 5 examples.

Back 67

An element with the properties of both metals and non-metals. Ex. Si, As, Ge, Sb, Te

Front 68

Metals. List 6 characteristics

Back 68

Solid (except Hg); malleable & ductile; conduct heat & electricity; shiny reflective & lustrous; lose e- to form cations; good reducing agents

Front 69

Non-metals. List 7 characteristics

Back 69

1/2 are gases; solids are brittle; poor conductors of heat and electricity; dull & non-reflective; gain e- to form anions; good oxidizing agents

Front 70

Transition element

Back 70

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

Front 71

Describe the change in atomic radius across the periodic table

Back 71

The atomic radius tends to decrease from left to right across the table and increase down the columns

Front 72

Ionization energy

Back 72

The energy change required for the removal of the outermost electron from the gaseous atom to form a +1 ion

Front 73

Describe the change in Ionization Energy across the periodic table

Back 73

Ionization energy tends to increase from left to right across the table and decrease down the columns

Front 74

Electronegativity

Back 74

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)

Front 75

Describe the change in electronegativity across the periodic table

Back 75

Electronegativity tends to increase across the table left to right and decrease down the columns

Front 76

Describe the change in metallic activity across the periodic table

Back 76

Metallic activity decreases across the table (left to right) and increases down the columns

Front 77

Describe the change in nonmetallic activity across the periodic table

Back 77

Non-metallic activity increases across the table (L to R) and decreases down the columns

Front 78

Compare the ionic radii of 1) Cl, S, P; 2) Na, Mg, Al; What is the reason for their different size?

Back 78

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-)

Front 79

Quantum

Back 79

A packet of energy associated with a specific wavelength of electromagnetic radiation

Front 80

Quantum Number

Back 80

A number used to describe the energy levels available to an electron. Each electron in an atom has a unique set of four.

Front 81

Emission Spectrum

Back 81

A bright line spectrum formed when energy absorbed by an element is emitted at specific wavelengths. Each element has a unique spectrum.

Front 82

Absorption Spectrum

Back 82

A dark-line spectrum formed when white light is passed through a vaporized element and a few specific wavelengths are absorbed.

Front 83

List three basic postulates of the bohr model for the hydrogen atom

Back 83

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

Front 84

Ground State

Back 84

Electron is at its lowest energy level as close to the nucleus as possible

Front 85

Excited State

Back 85

An electron absorbs energy and moves to a higher energy level above the ground state

Front 86

List four quantum numbers and their symbols

Back 86

1) Principal energy level "n"; 2) sublevel "l"; 3) orbital "m1"; 4) spin "ms"

Front 87

List the names of the first four energy levels (or shells)

Back 87

1 or K; 2 or L; 3 or M; 4 or N

Front 88

List the names of the four sublevels and their electron capacities

Back 88

sublevel s holds two electrons; p holds 6; d holds 10; f holds 14

Front 89

Which sublevels are present in energy levels 1, 2, 3, and 4-7?

Back 89

1: s; 2: s, p; 3: s, p, d; 4-7: s, p, d, f

Front 90

Describe the shapes of the orbitals: 1) s; 2) p; 3) d; 4) f

Back 90

1) s- sphere; 2) p - dumbbell shape with 2 lobes; 3) d - double dumbbell; most have 4 lobes; 4) f - most have 8 lobes

Front 91

What is the order for filling sublevels (aufbau process) from lowest to highest energy

Back 91

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d

Front 92

Electron configuration

Back 92

The distribution of electrons into shells and sublevels for an atom of an element. Each element has a unique electron configuration.

Front 93

Write the electron configuration for 1) Lithium; 2) Iron

Back 93

1) 1s2, 2s1; 2) 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d6

Front 94

Valence

Back 94

The number of electrons in the atom's highest numbered shell.

Front 95

What are the valences of the elements of families IA through VIIIA?

Back 95

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

Front 96

Electronegativity difference

Back 96

A number found by taking the difference between the electronegativities of two atoms in a bond. Its value determines the type of bond.

Front 97

Describe an ionic bond in terms of electronegativity difference

Back 97

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.

Front 98

Nonpolar covalent bond

Back 98

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.

Front 99

Polar covalent bond

Back 99

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.

Front 100

Coordinate covalent bond

Back 100

When both electrons in a covalent bond are supplied by one atom

Front 101

Metallic Bonds

Back 101

A sea of electrons surrounding positive metal ions

Front 102

Hydrogen bonding

Back 102

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

Front 103

Octet Rule

Back 103

Atoms tend to gain or lose outer shell electrons in order to achieve a noble gas configuration of 8 electrons

Front 104

Double and Triple covalent bond

Back 104

In a double bond, wo pairs of electrons are shared. In a triple bond, three pairs of electrons are shared.

Front 105

Resonance structures

Back 105

Where there is more than one possible bonding structure in a molecule

Front 106

Hybrid orbitals (list three types)

Back 106

Where 2 or more pure atomic orbitals are mixed to form identical hybrid orbitals (ex. Sp, sp2, sp3)

Front 107

Describe hybrid bonding in water, ammonia, methane

Back 107

sp3 bonding results in a tetrahedron shape with bond angles of 109.5˚ in methane and slightly less in water and ammonia.

Front 108

Sigma bonds & Pi bonds

Back 108

A sigma bond is present between any 2 orbitals except when 2 p orbitals share electrons; then this is a pi bond

Front 109

List properties of ionic substances

Back 109

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

Front 110

List properties of molecular substances

Back 110

Non-conducting as liquids and solids; Volatile liquids & solids; Many are gases at 25C; Low melting and boiling points; Soft and waxy solids

Front 111

Exothermic

Back 111

A chemical reaction which evolves heat

Front 112

Endothermic

Back 112

A chemical reaction which absorbs heat

Front 113

Enthalpy

Back 113

The heat content of a system

Front 114

Enthalpy change (state the equation)

Back 114

The difference in heat content between the products and the reactants (ΔH = ΣHproducts - Σhreactants)

Front 115

Exothermic reaction: ΔH is (positive or negative); Enthalphy is (increased or decreased)

Back 115

Negative; Decreased

Front 116

Endothermic Reaction: ΔH is (positive or negative); Enthalphy is (increased or decreased)

Back 116

Positive; Increased

Front 117

List in the order of increasing enthalpy: solid, gas, liquid

Back 117

Solid (least enthalpy) -> Liquid -> Gas (most enthalpy)

Front 118

Thermochemical equation

Back 118

A balanced chemical equation which includes the enthalpy change. (H2(g) +1/2O2(g) →H2O(l) ; ΔH = -285kJ

Front 119

In a thermochemical equation, what happens to ΔH when the moles of reactants double?

Back 119

Enthalpy is directly proportional to mass. Therefore when the moles double, so does ΔH.

Front 120

How are ΔH for a forward and ΔH for a reverse reaction related?

Back 120

forward is equal in magnitude but opposite in sign to ΔH reverse.

Front 121

Hess' Law

Back 121

for a reaction is the same regardless of the path travelled from reactants to products.

Front 122

Molar Heat of Formation

Back 122

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

Front 123

Write the general equation to calculate ΔH for a chemical reaction

Back 123

ΔH = Σ(ΔHF-products) - Σ(ΔHF-reactants); where ΔHf = Heat of formation of reactants or products

Front 124

Specific Heat. Give value for liquid water in calories and joules.

Back 124

The amount of heat required to raise one gram of substance 1C. Water: 1cal/g-C or 4.18 J/g-C

Front 125

Write an equation to calculate energy change when a fixed mass of substance changes temperature

Back 125

ΔH = mCpΔT, where ΔT = temperature change, Cp = Specific Heat; ΔH = heat absorbed or given off

Front 126

Charles Law

Back 126

V1/T1 = V2/T2 (pressure and amount of gas are constant; V=volume; T=Kelvin)

Front 127

Boyles Law

Back 127

P1V1 = P2V2 (Temperature and amount of gas are constant, P=pressure, V=vol)

Front 128

Combined Gas Law

Back 128

P1V1/T1 = P2V2/T2 (amount of gas is contant. P=pressure, V=vol; T=Kelvin)

Front 129

Dalton's Law of Partial Pressures

Back 129

In a gas mixture, the total pressure equals the sum of the partial pressures of each component. Ptotal = P1 + P2 + P3…

Front 130

Ideal Gas Law

Back 130

PV = nRT, (P=pressure in atm; V = volume in L; n = # moles; R = 0.0820 Latm/Mol-K; T = Temp in K)

Front 131

STP

Back 131

Standard temperature (0C or 273 K) and Standard Pressure (1 atm or 760 torr)

Front 132

Gay-Lussac's Law of Combining Gas Volumes

Back 132

When only gases are involved in a reaction, the volumes of reactants and products are in a small, whole number ratio.

Front 133

Avogadro's Law re. gases

Back 133

Under the same conditions of temperature and pressure; equal volumes of gases contain equal numbers of moles

Front 134

What is the volume of 1 mole of any gas at STP?

Back 134

22.4 L

Front 135

Phase equilibrium

Back 135

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.

Front 136

Dynamic equilibrium

Back 136

In a closed container where opposing changes are taking place at equal rates; the concentration of all components remains constant.

Front 137

Normal Boiling Point

Back 137

The temperature at which a liquid phase becomes a gas phase at a pressure of 1 atm.

Front 138

Heat of Fusion (value for water)

Back 138

The amount of energy required to change a gram of substance from solid to liquid at its melting point (water = 80 cal/g)

Front 139

Heat of Vaporization (value for water?)

Back 139

The amount of energy required per gram to change a liquid to a gas at its boiling point (water = 540 cal/g)

Front 140

Triple Point

Back 140

The only temperature and pressure combination at which the 3 phases of a substance (solid, liquid, gas) can co-exist in equilibrium

Front 141

Vapor Pressure

Back 141

The pressure the gas phase exerts on its liquid phase in a closed container. This pressure varies with temperature

Front 142

Molarity - general equation

Back 142

The number of moles of solute it a liter of solution; M = n/L

Front 143

Solute (provide example)

Back 143

The substance dissolved in another (solvent). Salt is the solute in salt water.

Front 144

Solvent (provide example

Back 144

A substance, usually a liquid, into which another substance (solute) is dissolved. Water is the solven in iced tea.

Front 145

Solution (provide example)

Back 145

A liquid, gas or solid phase containing 2 or more components uniformly dispersed (air, coffee, saltwater)

Front 146

Solubility curves

Back 146

A curve for a given substance which shows how much dissolves in a given amount of solvent at different temperatures.

Front 147

How do temperatures and pressure affect the solubility of a solid?

Back 147

Solubility usually increases with increasing temperature. Pressure has little effect.

Front 148

How do temperature and pressure affect the solubility of a gas?

Back 148

Solubility usually decreases with increasing temperature. Solubility increases in direct proportion to an increase in pressure.

Front 149

Which three factors affect the rate of solubility?

Back 149

Pulverizing; stirring; heating

Front 150

What is a general rule for solubilities of polar and nonpolar compounds?

Back 150

"Like dissolves like"; Ionic and polar solvents dissolve ionic, polar solutes (water dissolves salt). Non polar solvents dissolve nonpolar solutes (acetone dissolves gasoline)

Front 151

List some basic facts about solutions

Back 151

Particle size less than 1 mmicron; Clear (may be colored); Particles don't settle; Can pass through membranes; Particles not visible

Front 152

List some basic facts about colloids

Back 152

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

Front 153

List some basic facts about suspensions

Back 153

No brownian motion; Don't pass through filter paper or a membrane; Cloudy but particles settle on standing; Particles visible with microscope or eye

Front 154

How many grams of NaCl are required to prepare 500 grams of a 5% solution?

Back 154

%Concentration = (gNaCl/gsolution)(100%); 5% =(x/500g)(100%);x=25 g NaCl

Front 155

Calculate the molality of 10 moles of H2SO4 dissolved in a 4 kg of water

Back 155

Molality = Moles Solute / kg solvent = 10/4 = 2/5 Molal

Front 156

Gram-equivalent weight

Back 156

The amount of substance which reacts with or displaces 1 mole of H+ ions.

Front 157

Normality

Back 157

The number of gram-equivalent weights in a liter of solution

Front 158

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?

Back 158

1) 1.86C for each molal of particles of solute; 2) 0.51C for each molal of particles of solute

Front 159

List five factors that control reaction rate

Back 159

nature of reactants; exposed surface area; concentrations; temperatures; presence of catalyst

Front 160

State the collision theory of reaction rates

Back 160

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)

Front 161

How is the reaction rate related to concentration?

Back 161

Reaction rate is directly proportional to the concentrations of reactants

Front 162

Activation energy

Back 162

The energy necessary for a reaction to begin. Obtained from the kinetic energy released during collision

Front 163

Catalyst

Back 163

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.

Front 164

Law of Mass Action

Back 164

The rate of a chemical reaction is proportional to the product of the concentrations of the reactants

Front 165

Specific Rate Constant

Back 165

Symbol is "k" in a rate equation. A constant specific to temperature and reaction which is part of every rate equation

Front 166

Reversible reaction

Back 166

A system where the following opposite reactions are taking place: reactant becoming product; product becoming reactant

Front 167

Equilibrium

Back 167

The point in a reversible reaction where the forward and reverse reactions are taking place at the same rate.

Front 168

Are concentrations of product and reactant equal at equilibrium?

Back 168

No. The are constant but not equal. Their relative concentrations are determined by the value of the equilibrium constant at that temperature.

Front 169

Write the equilibrium expression for aA + bB ↔ cC + dD

Back 169

Keq = [C]^c[D]^d / [A]^a[B]^b, where Keq = Equilibrium constant

Front 170

How are reactant & product concentrations related to the magnitude fo Keq?

Back 170

Keq is large: [reactant] is small and [product] is large; when Keq is small: [reactant] is large and [product] is small

Front 171

Le Chatelier's Principle

Back 171

If stress is placed on a system at equilibrium, the equilibrium shifts in order the counteract the effects of the stress and regain equilibrium

Front 172

How does a concentration change affect equilibrium?

Back 172

If one of the substances is added or removed, all the concentrations of substances adjust to a new equilibrium with the same Keq

Front 173

How does a change in temperature affect equilibrium

Back 173

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.

Front 174

How does a pressure change affect equilibrium?

Back 174

Only in reactions where gases are involved. The reaction will shift to oppose pressure change, resulting in fewer moles of gas particles

Front 175

Ionization constant. Write the expression for the ionization of acid "HA"

Back 175

For substances in solution that partially ionize. An equilibrium expression may be written with Ki; Ki = [H+][A-] / [HA]

Front 176

What are the two driving forces that control reactions?

Back 176

A drive towards increased entropy (disorder). A drive towards decreased enthalpy (lower heat content)

Front 177

Second Law of Thermodynamics

Back 177

The entropy of the universe increases for any spontaneous process.

Front 178

Free Energy Change (ΔG). Write the free energy equation.

Back 178

A property which reflects a system's capacity to do useful work. ΔG = ΔH - TΔS; G=free energy; S=entropy; H=enthalpy; T=kelvin

Front 179

How can ΔG be used to predict if a reaction is spontaneous?

Back 179

When ΔG is positive it is not spontaneous; negative, it is spontaneous; equals 0 it is at equilibrium.

Front 180

Ionization Constant

Back 180

Kw = 1 x 10-14 at 25C

Front 181

pH. For what values is a solution acidic, basic and neutral?

Back 181

pH = -log[H+] ; The degree of acidity of a solution. <7 =acid; 7 = neutral; >7 = basic.

Front 182

pOH. For what values is a solution acidic, basic, neutral?

Back 182

pOH = -log[OH-]; The degree of basicity of a solution. <7 = basic; 7 = neutral; >7=acid.

Front 183

How are pH and pOH of a solution related?

Back 183

The sum of the two values = 14; (pH +pOH = 14)

Front 184

How are the concentrations of [H+] and [OH-] related in a solution?

Back 184

The product of the concentrations = 1x10-14; [H+][OH-] = 1x10-14

Front 185

Solubility Product Constant (Ksp)

Back 185

An equilibrium exists in a saturated solution between dissolved and undissolved solute. Ksp is the equilibrium constant for this reaction.

Front 186

Write the solubility product expression for AgCl ↔ Ag+ + Cl-

Back 186

Ksp = [Ag+][Cl-]

Front 187

Six common characteristics of acids

Back 187

Form H2O solutions; Conduct electricity; React with active metals; Turn blue litmus red; Neutralize bases; Sour taste

Front 188

Seven common characteristics of bases

Back 188

Form H2O solutions; Conducts electricity; Turns red litmus blue; Feels slippery; Caustic; Neutralizes acids; Bases + fats form SOAP

Front 189

Neutralization Reaction - Write equation for hydrochloric acid and sodium hydroxide

Back 189

acid + base → salt + water; HCl + NaOH →NaCl +H2O

Front 190

Arhennius Theory

Back 190

An acid yields protons in solution (H+ ions); A base yields hydroxide ions in solution (OH- ions)

Front 191

Bronsted Theory

Back 191

An acid is a proton donor. A base is a proton acceptor.

Front 192

Lewis Theory

Back 192

An acid is an electron pair acceptor; a base is an electron pair donor.

Front 193

Conjugate Base - Write conjugate base of HCl

Back 193

When a Bronsted acid donates a proton, it becomes its conjugate base (conjugate base of HCl is Cl-)

Front 194

Conjugate Acid - Write conjugate acid of I-

Back 194

When a Bronsted base accepts a proton it becomes its conjugate acid (conjugate acid of I- is HI)

Front 195

Calculate the volume of 10M NaOH needed to titrate 5L of 2M HCl

Back 195

MA x VA = MB x VB; 2M x 5L = 10M x VB; 1L = volume of base

Front 196

Electrolyte (give 3 examples)

Back 196

A solute whose aqueous solution contains ions and conducts electricity (acids, bases, salts)

Front 197

Nonelectrolyte (give 2 examples)

Back 197

A solute whose aqueous solution does not conduct electricity (sugar, benzene, most organic compounds)

Front 198

Anode

Back 198

A positively charged electrode which attracts anions. Where oxidation takes place.

Front 199

Cathode

Back 199

A negatively charged electrode which attracts cations. Where reduction takes place.

Front 200

Oxidation

Back 200

The loss of electrons (ex. Cumetal →Cu+2 +2e-

Front 201

Reduction

Back 201

The gain of electrons (Zn+2 +2e- → Znmetal)

Front 202

Electrode Potential

Back 202

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)

Front 203

Half Reaction

Back 203

One of the two parts, either the reduction or the oxidation, of an oxidation-reduction reaction

Front 204

How can you determine if a redox reaction will take place spontaneously

Back 204

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.

Front 205

Faraday

Back 205

A unit of electric charge which deposits by electrolysis one equivalent weight of an element. Equals 96,500 coulombs

Front 206

Equivalent weight

Back 206

The number of grams of an element which will accept or donate 1 mole of electrons

Front 207

Oxidizing agent

Back 207

A substance which causes another substance to be oxidized (oxidizing agent is simultaneously reduced)

Front 208

Reducing agent

Back 208

A substance which causes another substance to be reduced (reducing agent is simultaneously oxidized)

Front 209

List some basic facts about carbon bonding

Back 209

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

Front 210

Alkane

Back 210

A series of hydrocarbons with only single covalent bonds (CnH2n+2)

Front 211

Alkene

Back 211

A series of hydrocarbons containing at least one double covalent bond (CnH2n)

Front 212

Alkyne

Back 212

A series of hydrocarbons containing at least one triple covalent bond (CnH2n-2)

Front 213

List prefixes for naming hydrocarbons for 1-10 carbons in a molecule

Back 213

1 meth; 2 eth; 3 prop; 4 but; 5 pent; 6 hex; 7 hept; 8 oct; 9 non; 10 dec

Front 214

List names for alkanes with 1 to 10 carbons in a molecular chain

Back 214

1 methane; 2 ethane; 3 propane; 4 butane; 5 pentane; 6 hexane; 7 heptane; 8 octane; 9 nonane; 10 decane

Front 215

Substitution reaction. Which elements commonly substitute in alkanes?

Back 215

A reaction where one of the hydrogen atoms in a hydrocarbon is replaced by another. Usually a halogen. (CH4 + Br2→Ch3Br + HBr

Front 216

Cycloalkane

Back 216

An alkane which has a ring structure instead of a chain

Front 217

Saturated and Unsaturated. Which are sturated? Alkanes, alkenes, alkynes.

Back 217

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.

Front 218

Addition reaction

Back 218

In an unsaturated hydrocarbon, two atoms may be added to the structure across a double or triple bond (C2H2 + Br2 → Ch2Br2)

Front 219

Aromatic compounds

Back 219

Unsaturated ring structures with six carbon atoms. Benzene is the simplest aromatic. (CnH2n-6)

Front 220

Isomer

Back 220

Compounds with the same molecular formula but different structural formulas (different connectivity)

Front 221

Polymerization

Back 221

The combination of two or more unsaturated molecules to form a larger chain molecule. This is how plastics are made.

Front 222

Hydrogenation

Back 222

The process of adding hydrogen to an unsaturated hydrocarbon

Front 223

Dehydrogenation

Back 223

The process of removing hydrogen from a hydrocarbon

Front 224

Alcohol

Back 224

hydrocarbons which contain the hydroxyl functional group (OH-) attached to a saturated carbon (R-O-H)

Front 225

Aldehyde

Back 225

A hydrocarbon containing the aldehyde functional group (R-C(=O)-H)

Front 226

Carboxylic Acid

Back 226

A hydrocarbon containing the carboxyl functional group. (R-C(=O)-O-H)

Front 227

Ketone

Back 227

A hydrocarbon containing a ketone functional group (R-C(=O)-R')

Front 228

Ether

Back 228

A hydrocarbon containing an ether functional group. (R-O-R')

Front 229

Amine

Back 229

A hydrocarbon containing an amine functional group. (R-NH2)

Front 230

Ester

Back 230

A hydrocarbon containing an ester functional group. (R-O-C(=O)-R')

Front 231

Amino Acid

Back 231

Organic compounds which contain an amine and a carboxyl group. (H2N-CH(-R)-COOH)

Front 232

Compare primary and secondary alcohols

Back 232

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.

Front 233

Alpha particle

Back 233

a helium nucleus; charge =+2; High energy; Low velocity; Ejection reduces atomic number by 2amu and atomic weight by 4amu

Front 234

Beta particle

Back 234

An electron ejected from the nucleus when a neutron decays to a proton; Increases atomic number by one; High velocity; Low energy

Front 235

Gamma radiation

Back 235

Usually emitted with beta radiation; Has neither charge nor mass; High energy; Travels at the speed of light

Front 236

Half-life

Back 236

The time it takes for half of a radioactive sample to decay. It can range from a fraction of a second to many years.

Front 237

Transmutation

Back 237

The conversion of an element to a new element due to a change in number of protons. Ex. Alpha or beta decay

Front 238

Fission

Back 238

The breakdown of heavy nuclei into lighter nuclei. The source of nuclear power

Front 239

Fusion

Back 239

The joining of lighter nuclei to form heavier nuclei. Source of the sun's energy

Front 240

Write the nuclear equation: 1) U238 loses an alpha particle; 2) Th234 loses a beta particle.

Back 240

92U238 →90Th234 +2He4; 90Th234→ 91Pa234 + -1e0

Front 241

Atoms

Back 241

Tiny particles making up mass

Front 242

Nucleus

Back 242

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

Front 243

Neutrons

Back 243

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

Front 244

Protons

Back 244

Part of nucleus

Together with neutron, makes up nucleon

Same size and mass as neutron (1 amu)

Slightly lighter than neutron

Positive charge

Front 245

Electrons

Back 245

Surround nucleus at distance of 1-3 angstroms

Over 1800 lighter than nucleon

Electrons (negative charge) and protons have opposite charges of equal magnitude

Front 246

Electron charge (e)

Back 246

Charge of one electron

e = 1.6e^-19 coulombs (C)

Front 247

Atom

Back 247

Electrically neutral

Made up of neutrons, protons and electrons

Same number of protons as electrons

Composed mostly of empty space

Front 248

Elements

Back 248

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)

Front 249

Mass number (A)

Back 249

Number of protons plus neutrons

Front 250

Atomic number (Z)

Back 250

number of protons

Identity number of any element

Front 251

Isotopes

Back 251

Two or more atoms of the same element that contain different numbers of neutrons

Front 252

Atomic Weight

Back 252

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

Front 253

Atomic mass units (amu)

Back 253

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

Front 254

Mole

Back 254

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

Front 255

Periodic table

Back 255

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)

Front 256

Period

Back 256

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

Front 257

Groups or Families

Back 257

Vertical columns of periodic table

Front 258

Metals

Back 258

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)

Front 259

Metallic Character

Back 259

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

Front 260

Nonmentals

Back 260

Diverse appearances and chemical behaviors

Lower melting points than metals

Form negative ions

Make up molecular substances

Form covalent oxides (ie: SiO2 or CO2)

Front 261

Alkali metals

Back 261

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

Front 262

Alkaline earth metals

Back 262

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

Front 263

Halogens

Back 263

Group (family) 17 in periodic table

Nonmetals & metalloids

Front 264

Noble gases

Back 264

Groups (family) 18 in periodic table

Nonmetals

Also known as rare gases

Nonreactive, inert gases at room temperature

Front 265

Metalloids

Back 265

Characteristics that resemble metals and nonmetals

Front 266

Representative or main-group elements

Back 266

Section A groups in periodic table

Front 267

Transition metals

Back 267

Section B groups in periodic table

When form ions, they lose electons from s-subshell first and then from d-subshell

Front 268

Hydrogen

Back 268

Is unique and chemical/physical properties to do conform to any family

Nonmetal

Colorless

Odorless

Diatomic gas

Front 269

Group 4A

Back 269

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

Front 270

Group 5A

Back 270

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

Front 271

Group 6A

Back 271

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

Front 272

Group 7A

Back 272

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

Front 273

Diatomic molecules

Back 273

1. Hydrogen
2. Oxygen
3. Nitrogen
4. Halogens

Front 274

Small atoms

Back 274

Make strong pi bonds due to overlap of p-orbitals

Front 275

Large atoms

Back 275

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

Front 276

Pi bonds

Back 276

Allow for double and triple bonds

Front 277

Ion

Back 277

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

Front 278

Cation

Back 278

Positive ion

Formed by metals

Significantly smaller than neutral atom counterparts

Front 279

Anion

Back 279

Negative ion

Formed by nonmetals

Much larger than neutral atom counterpart

Front 280

Predict ion charge based on:

Back 280

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

Front 281

Electron shielding

Back 281

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

Front 282

Effective nuclear charge (Zeff)

Back 282

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

Front 283

Periodic trends

Back 283

1. Atomic radius
2. Ionization energy
3. Electronegativity
4. Electron affinity
5. Metallic character

Front 284

Atomic radius

Back 284

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

Front 285

Ionization energy

Back 285

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)

Front 286

Electronegativity

Back 286

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

Front 287

Electron affinity

Back 287

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

Front 288

SI units

Back 288

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

Front 289

Bonds

Back 289

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

Front 290

Covalent bonds

Back 290

2 electrons are shared by 1 nuclei

Negatively charges electrons are pulled toward both positively charged nuclei by electrostatic forces

Front 291

Bond length

Back 291

Point where energy level is lowest

Front 292

Bond dissociation energy (bond energy)

Back 292

Energy necessary to achieve a complete separation of atoms

Front 293

Compound

Back 293

Substance made from 2 or more elements

Front 294

Empirical formula

Back 294

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

Front 295

Molecules

Back 295

Separate and distinct units, in molecular compounds, formed from repeated groups of atoms

Front 296

Molecular formula

Back 296

Exact number of elemental atoms in each molecule in a molecular compound

Glucose = C6H12O6

Front 297

Percent mass

Back 297

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

Front 298

Ionic compounds

Back 298

Named after their cation and anion

Put cation name in from on anion name (barium sulfate, BaSO4; sodium hydride, NaH)

Front 299

Cation nomenclature

Back 299

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+)

Front 300

Anion nomenclature

Back 300

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-)

Front 301

Acid nomenclature

Back 301

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)

Front 302

Binary molecular compounds

Back 302

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

Front 303

Physical reaction

Back 303

When compound undergoes a reaction and maintains its molecular structure and this its identity

Ex:
Melting, evaporation, dissolution and rotation of polarized light

Front 304

Chemical reaction

Back 304

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

Front 305

Reaction runs to completion

Back 305

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

Front 306

Limiting reagent

Back 306

Reactant which is depleted first if reaction runs to completion

Not necessarily the reactant of which there is least

Front 307

Theoretical yield

Back 307

Amount of product produced when a reaction runs to completion

Front 308

Actual yield

Back 308

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

Front 309

Percent yield

Back 309

[(Actual yield)/(Theoretical yield)] x100

Front 310

Reaction types

Back 310

1. Combination
2. Decomposition
3. Single displacement
4. Double displacement
5. Redox
6. Combustion
7. Bronsted-Lowry Acid-base
8. Lewis Acid-base

Front 311

Combination reaction

Back 311

A + B --> C

Front 312

Decomposition reaction

Back 312

C --> A + B

Front 313

Single Displacement reaction

Back 313

A + BC --> B + AC

Also called single replacement

Front 314

Double Displacement reaction

Back 314

AB + CD --> AC + BD

Also called double replacement or metathesis

Front 315

Quantum mechanics

Back 315

Elementary particles can only gain or lose energy and certain other quantities in discrete units

Front 316

Principal quantum number (n)

Back 316

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

Front 317

Valence electrons

Back 317

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

Front 318

Azimuthal quantum number (l)

Back 318

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

Front 319

Magnetic quantum number (ml)

Back 319

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)

Front 320

Electron spin quantum number (ms)

Back 320

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

Front 321

Pauli Exclusion Principle

Back 321

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

Front 322

Number of orbitals within a shell

Back 322

n^2

Front 323

Heisenberg Uncertainty Principle

Back 323

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

Front 324

Aufbau Principle

Back 324

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

Front 325

Electron configuration

Back 325

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

Front 326

Abbreviated electron configuration

Back 326

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

Front 327

Ground state

Back 327

atom whose electrons are all their lowest energy levels

Front 328

Electron configuration for ions

Back 328

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

Front 329

Hund's Rule

Back 329

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

Front 330

Planck's Quantum Theory

Back 330

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

Front 331

Photoelectric effect

Back 331

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

Front 332

Work function (ϕ)

Back 332

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

Front 333

Standard temperature and pressure (STP)

Back 333

Temp: 0 degrees Celsius = 273 K
Pressure: 1 atm (atmosphere)
moles: 1 mole

Front 334

Mean free path

Back 334

Distance traveled by a gas molecule between collisions

Front 335

Ideal gas

Back 335

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

Front 336

Kinetic molecular theory

Back 336

Ideal gas lack certain real gas characteristics

Front 337

Ideal gas law

Back 337

PV = nRT
P: pressure
V: volume
n: number of moles
R: universal gas constant = 8.314 J/Kmol
T: temperature in kelvin

Front 338

Charles' Law

Back 338

Volume of gas is proportional to temperature at constant pressure

Front 339

Boyle's Law

Back 339

Volume of gas is inversely proportional to pressure a constant temperature

Front 340

Avogadro's Law

Back 340

Volume of gas is proportional to number of moles at constant temperature and pressure

Front 341

Standard molar volume at STP

Back 341

22.4 Liters

Front 342

Partial pressure

Back 342

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

Front 343

Mole fraction (X)

Back 343

number of mole of gas "a" divided by total number of moles of gas in sample

Front 344

Dalton's Law

Back 344

Total pressure exerted by a gaseous mixture is the sum of the partial pressures of each of its gases

Ptotal = P1 + P2 + P3...

Front 345

Average translational energy of ideal gas

Back 345

KE = (3/2) RT
KE: kinetic energy, found from RMS velocity
R: universal gas constant
T: temperature

Front 346

Graham's Law

Back 346

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

Front 347

Effusion

Back 347

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

Front 348

Pinhole

Back 348

opening much smaller than average distance between gas molecules

Front 349

Diffusion

Back 349

Spreading of one gas into another gas or empty space

diffusion rate 1/diffusion rate 2 = square root (M2)/square root (M1)

Front 350

How real gases deviate from ideal gases

Back 350

Vreal > Videal

Intermolecular forces exist

Preal < Pideal

Front 351

Collision model

Back 351

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

Front 352

Activation Energy

Back 352

Threshold energy required for collisions to create new molecules in a reaction (Ea) - independent of temperature

Rate of reaction (K) increases with temperature

Front 353

Factors affecting rate of reaction

Back 353

1. temperature
2. pressure (negligible effect)
3. concentration of reactants

Front 354

Elementary reaction

Back 354

Coefficient tells you how many molecules participate in a reaction producing collision

Front 355

Intermediates

Back 355

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

Front 356

Rate Law

Back 356

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

Front 357

1st Order reaction

Back 357

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

Front 358

2nd Order Reaction (1 reactant)

Back 358

2A --> products

rate = Kf[A]^2

1/[A] vs time graph: straight line
slope: Kf

Half-life dependent upon concentration [A]

Front 359

2nd Order Reaction (2 reactants)

Back 359

A + B --> products

Rate = Kf[A][B]

no predictable half-life

Front 360

3rd Order Reaction

Back 360

3A --> products

Rate = Kf[A]^3

1/2[A]^2 vs. time graph: straight line
slope: Kf

Front 361

Zero Order Reaction

Back 361

[A] vs. time graph: straight line
slope: -Kf

Front 362

Rate determining step

Back 362

Rate of slowest elementary step determines the rate of the overall reaction

Front 363

Catalyst

Back 363

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)

Front 364

Heterogeneous catalyst

Back 364

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

Front 365

Homogeneous catalyst

Back 365

same phase as reactants and products, usually in gas or liquid phase

Aqueous acid or base solutions

Autocatalysis = generate catalyst as product

Front 366

Chemical equilibrium

Back 366

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

Front 367

Equilibrium constant (K)

Back 367

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

Front 368

Reaction quotient (Q)

Back 368

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

Front 369

Comparison of Q & K

Back 369

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

Front 370

Le Chatelier's Principle

Back 370

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

Front 371

Thermodynamics

Back 371

Study of energy and its relationship to macroscopic properties of chemical systems

divides universe into 2:
1. surroundings
2. system

Front 372

System

Back 372

macroscopic body under study

3 types:
1. open
2. closed
3. isolated

Front 373

Surroundings

Back 373

everything other than macroscopic body under study (system)

Front 374

Open system

Back 374

exchange both mass and energy with surroundings

Front 375

Close system

Back 375

exchange energy but do not exchange mass with surroundings

Front 376

Isolated system

Back 376

doesn't exchange mass or energy with surroundings

Front 377

Extensive properties

Back 377

chance with amount

Examples:
volume (V) and # of moles (n)

describes macroscopic state of a system

proportional to size of system

Front 378

Intensive properties

Back 378

describes macroscopic state of a system

independent of size of system

examples:
Pressure (P) and temperature (T)

Front 379

State functions

Back 379

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

Front 380

2 ways to transfer energy between systems

Back 380

1. Heat (Q)
2. Work (W)

Front 381

Heat (Q)

Back 381

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

Front 382

Work (W)

Back 382

any energy transfer between systems that is not heat

Front 383

Conduction

Back 383

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

Front 384

Energy flow

Back 384

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

Front 385

Convection

Back 385

thermal energy transfer via fluid movements

differences in pressure or density drive warm fluid in direction of cooler fluid

examples:
ocean and air currents

Front 386

Radiation

Back 386

thermal energy transfer via electromagnetic waves

only type of heat that transfer though a vacuum

Front 387

PV work

Back 387

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

Front 388

1st Law of Thermodynamics

Back 388

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

Front 389

2nd Law of Thermodynamics

Back 389

Heat cannot be changed completely into work in a cyclical process

Qh = W + Qc
Qh: heat entering engine
W: work
Qc: heat leaving engine

Front 390

Carnot's efficiency

Back 390

e = 1 - (Tc/Th)

e: efficiency
Tc: temperature of cold reservoir
Th: temperature of hot reservoir

Front 391

7 state functions

Back 391

1. internal energy (U)
2. temperature (T)
3. pressure (P)
4. volume (V)
5. enthalpy (H)
6. entropy (S)
7. Gibbs free energy (G)

Front 392

Internal energy (U)

Back 392

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

Front 393

Temperature

Back 393

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

Front 394

Enthalpy (H)

Back 394

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]

Front 395

Standard State

Back 395

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

Front 396

Reference form

Back 396

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

Front 397

Standard enthalpy of formation (change in Hf)

Back 397

Change in enthalpy for a reaction that creates 1 mole of that compound from its raw elements in their standard state

Front 398

Change in enthalpy at constant pressure

Back 398

Change in H = Q
constant pressure, closed system at rest, PV work only
H: enthalpy
Q: heat

Front 399

Heat of reaction

Back 399

change in enthalpy from reactants to products

Change in Hreaction = (Change in Hfproducts) - (change in Hfreactants)

Front 400

Hess' Law

Back 400

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

Front 401

Endothermic

Back 401

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

Front 402

Exothermic

Back 402

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

Front 403

Activation energy

Back 403

initial increase in energy

Front 404

Transition state

Back 404

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)

Front 405

Catalyst

Back 405

lowers activation energy of forward and reverse reactions

equilibrium and enthalpy change is unaffected

Affects the rate of reaction

Front 406

Entropy (S)

Back 406

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

Front 407

2nd law of thermodynamics

Back 407

entropy of an isolated system will never decrease without some outside intervention (work)

Front 408

Entropy change of universe

Back 408

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

Front 409

Reversible

Back 409

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

Front 410

Irreversible

Back 410

all reactions are irreversible on a macroscopic scale

Front 411

what drives the direction of a reaction

Back 411

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

Front 412

Equilibrium

Back 412

point in reaction where universe has achieved maximum entropy

Front 413

3rd law of thermodynamics

Back 413

assigns by convention a zero entropy value to any pure substance (either element or compound) at absolute zero and in internal equilibrium

Front 414

Gibbs Free energy (G)

Back 414

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

Front 415

Solution

Back 415

homogeneous mixture of 2 or more compounds in a single phase (solid, liquid, gas)

Front 416

Solvent

Back 416

compound of which there is more in a solution

compound that predominates

Front 417

Solute

Back 417

compound of which there is less in a solution

Front 418

Ideal solution

Back 418

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)

Front 419

Ideally dilute solution

Back 419

solute molecules are completely separated by solvent molecules so that they have no interaction with each other

mole fraction of solvent is ~1

Front 420

Nonideal solution

Back 420

violate both ideal solution and ideally dilute solution conditions

Front 421

Colloid

Back 421

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

Front 422

Colloid solution

Back 422

scatter light, unlike true solution

can be attracted or repelled by dispersion medium (solvent)

Front 423

Disolution

Back 423

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

Front 424

Solvation

Back 424

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)

Front 425

Aqueous phase

Back 425

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)

Front 426

Polyatomic ions

Back 426

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

Front 427

Electrolyte

Back 427

compound which forms ions in aqueous solution and therefore solution can conduct electricity

Front 428

5 ways to measure concentration in solution

Back 428

1. molarity (M)
2. molality (m)
3. mole fraction (X)
4. mass percentage
5. parts per million (ppm)

Front 429

Molarity (M)

Back 429

moles of a compound divided by volume of solution

mol/L

M = moles solute/volume solution

Front 430

Molality (m)

Back 430

moles of solute divided by kilograms of solvent

mol/kg

m = moles solute/kg solvent

Front 431

mole fraction (X)

Back 431

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

Front 432

mass percentage

Back 432

100 times ratio of mass of solute to total mass of solution

mass % = (mass solute/total mass solution) x 100

Front 433

Part per million (ppm)

Back 433

10^6 times ratio of mass solute to total mass of solution

ppm = (mass solute)/(total mass solution) x 10^6

Front 434

3 steps of solution formation

Back 434

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

Front 435

heat of solution

Back 435

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

Front 436

vapor pressure

Back 436

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

Front 437

Boiling point

Back 437

temperature at which vapor pressure of liquid equals Patm

Front 438

melting point

Back 438

temperature at which vapor pressure of solid equals vapor pressure of liquid

Front 439

nonvolatile solute

Back 439

solute with no vapor pressure

when added to liquid, decreases vapor pressure

Front 440

Raoult's law

Back 440

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

Front 441

Raoult's law for nonvolatile solutes

Back 441

if 97% of solution is solvent, then vapor pressure will be 97% of vapor pressure of pure solvent

Front 442

Raoult's law of volatile solutes

Back 442

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

Front 443

Volatile solute

Back 443

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

Front 444

vapor pressure

Back 444

negative heats of solution = stronger bonds = lower vapor pressure

positive heats of solution = weaker bonds = higher vapor pressure

Front 445

solubility

Back 445

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

Front 446

precipitation

Back 446

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

Front 447

Saturate

Back 447

rate of dissolution and precipitation are equal

concentration of dissolved salt has reached max in saturated solution

Front 448

Solubility product (Ksp)

Back 448

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

Front 449

Spectator ions

Back 449

ions that have no effect on equilibrium because not in equilibrium expression

Front 450

Common ion effect

Back 450

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

Front 451

Solubility guidelines

Back 451

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

Front 452

Henry's Law

Back 452

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

Front 453

Heat Capacity (C)

Back 453

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

Front 454

Specific heat capacity (c)

Back 454

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)

Front 455

Coffee Cup calorimeter

Back 455

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

Front 456

Bomb calorimeter

Back 456

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

Front 457

Normal melting point

Back 457

melting point at constant pressure of 1atm, which is 0 degrees Celsius

Front 458

Normal boiling point

Back 458

boiling point at constant pressure of 1atm, which is 100 degree Celsius

Front 459

Heat of fusion

Back 459

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

Front 460

Heat of vaporization

Back 460

enthalpy (H) change associated with boiling

exactly same amount of heat is absorbed during vaporization as is released during condensation and sublimation/deposition

Front 461

phase change

Back 461

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

Front 462

evaporation

Back 462

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

Front 463

phase diagram

Back 463

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

Front 464

triple point

Back 464

point at which substance exists in equilibrium as solid, liquid and gas

Front 465

critical temperature

Back 465

temperature above which substance cannot be liquefied regardless of pressure applied

Front 466

critical pressure

Back 466

pressure required to produce liquefaction while substance is at critical temperature

Front 467

critical point

Back 467

where critical pressure and critical temperature intersect

fluid beyond this point has characteristics of both gas and liquid (supercritical fluid)

Front 468

colligative properties

Back 468

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

Front 469

boiling point elevation

Back 469

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

Front 470

Van't Hoff Factor (i)

Back 470

number of particles into which a single solute particle will dissociate when added to solution

Front 471

freezing point depression

Back 471

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)

Front 472

osmotic pressure

Back 472

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

Front 473

hydrostatic pressure

Back 473

pressure pushing out of a solution

Front 474

Arrhenius acid

Back 474

anything that produces H ions in aqueous solution

Front 475

Arrhenius base

Back 475

anything that produces OH ions in aqueous solution

Front 476

Bronsted and lowry acids

Back 476

anything that donates a proton

Front 477

Bronsted and lowry bases

Back 477

anything that accepts a proton

Front 478

Lewis acid

Back 478

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

Front 479

lewis base

Back 479

anything that donates a pair of electrons

Front 480

pH

Back 480

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

Front 481

estimating pH

Back 481

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)

Front 482

acid-base reaction

Back 482

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

Front 483

amphoteric

Back 483

some substances act as acid or base depending on environment

ex: water

Front 484

strong acids

Back 484

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

Front 485

Strong bases

Back 485

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

Front 486

hydronium ion

Back 486

H3O+

simply hydrated proton

in acid-base reactions, H3O+ and H+ are same thing

Front 487

polyprotic acid

Back 487

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)

Front 488

diprotic acid

Back 488

acids that can donate only 2 protons

is a polyprotic acid as well

Front 489

dissociation

Back 489

acid dissociation decreases with acid concentration but acid strength increase with acid concentration

Front 490

Acid-Base strengths

Back 490

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)

Front 491

3 factors determined whether a H containing molecule will act as acid (release H into solution)

Back 491

1. strength of bond holding H to molecule
2. polarity of bond
3. stability of conjugate base

Front 492

hydrogen halide acid strength

Back 492

weakest to strongest:

1. HF
2. HCl
3. HBr
4. HI

as acidity increases, polarity decreases and bond strength decreases

Front 493

oxyacid acid strength

Back 493

weakest to strongest:

1. HClO
2. HClO2
3. HClO3
4. HClO4

more O, means stronger acid

acidity increases with oxidation number of central atom

Front 494

hydrides

Back 494

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

Front 495

autoionization of water

Back 495

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

Front 496

acid dissociation constant (Ka)

Back 496

acid has its own equilibrium constant in water

HA + H2O --> H3O+ + A-

Ka = [H+][A-]/[HA]

Front 497

base dissociation constant (Kb)

Back 497

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-]

Front 498

product of Ka & Kb

Back 498

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

Front 499

Finding pH of strong acids and bases

Back 499

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

Front 500

Finding pH of weak acids and bases

Back 500

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

Front 501

Salts

Back 501

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

Front 502

titration

Back 502

drop-by-drop mixing of acid and base

performed in order to find concentration of an unknown by comparing it with concentration of titrant

Front 503

titrant

Back 503

known substance added drop-by-drop into unknown to find concentration of unknown

can be acidic or basic

Front 504

titration curve

Back 504

graphic representation changing pH of unknown as acidic or basic titrant is added (volume)

sigmoidal curve

Front 505

equivalence point

Back 505

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)

Front 506

equivalent

Back 506

mass of acid or base necessary to produce or consume 1 mole of protons

Front 507

half-equivalence point

Back 507

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

Front 508

Henderson-Hasselbalch equation

Back 508

pH = pKa + log([A-]/[HA])

form of equilibrium expression for Ka

when [A-] = [HA], pH = pKa which is at the half-equivalence point

Front 509

Finding pH at half-equivalence point

Back 509

Ka = [H+][A-]/[HA] = [H+] [A-]/[HA]

-log(Ka) = -log[H+] + -log([A-]/[HA])

pKa = pH - log[A-]/[HA]

pH = pKa + log[A-]/[HA]

Front 510

Finding pH at equivalence point

Back 510

Kb = Kw/Ka

Kb = [OH-][HA]/[A-]

solve for OH- concentration

Find pOH

Subtract pOH from 14 to find pH

Front 511

Titration curve of weak acid or base

Back 511

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

Front 512

indicator

Back 512

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

Front 513

range

Back 513

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

Front 514

endpoint

Back 514

point where indicator changes color

equivalence point does not equal endpoint

Front 515

Polyprotic titrations

Back 515

will have more than 1 equivalence point and more than 1 half-equivalence point

1st proton completely dissociates before 2nd proton begins to dissociate

Front 516

Redox reaction

Back 516

oxidation-reduction reaction

electrons are transferred form one atom to another

atom that loses electrons is oxidized

atom that gains electrons is reduced

Front 517

Oxidized

Back 517

atom loses electrons in redox reaction

Front 518

reduced

Back 518

atom gains electrons in redox reaction

Front 519

oxidation states

Back 519

possible charge values that an atom may hold within a molecule

necessary for redox reactions

must add up to charge on molecule or ion

Front 520

oxidation state = 0

Back 520

atoms in their elemental form

Front 521

oxidation state of fluorine (F)

Back 521

equals -1

Front 522

oxidation state of hydrogen (H)

Back 522

equals +1

except when bonded to a metal; then -1

Front 523

oxidation state of oxygen (O)

Back 523

equals -2

except when it is in a peroxide like H2O2

Front 524

Oxidation state = +1

Back 524

group 1 elements

alkali metals

Front 525

oxidation state = +2

Back 525

group 2 elements

alkaline earth metals

Front 526

oxidation state = -3

Back 526

group 15 elements

nitrogen family

Front 527

oxidation state = -2

Back 527

group 16 elements

oxygen family

Front 528

oxidation state = -1

Back 528

group 17 elements

halogens

Front 529

LEO the lion says GER

Back 529

LEO: Lose Electrons Oxidation

GER: Gain Electrons Reduction

Front 530

Reducing agent

Back 530

reductant

giving electrons to an atom

losing electrons, is oxidized

Front 531

Oxidizing agent

Back 531

compound containing atom being reduced

gains electrons, is reduced

oxidizes other atom

Front 532

Electric potential (E)

Back 532

associated with any redox reaction

Front 533

Half-reaction

Back 533

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)

Front 534

Half-reaction of standard hydrogen electrode

Back 534

2H+ + 2e- --> H2
Half-reaction potential = 0.00V

Front 535

Balancing redox reactions in acidic solutions

Back 535

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

Front 536

Balancing redox reactions in basic solutions

Back 536

same steps as acidic solutions

neutralize H+ ions by adding same number of OH- ions to both sides of reaction

Front 537

Galvanic cell (voltaic cell)

Back 537

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

Front 538

Salt bridge

Back 538

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

Front 539

Terminals

Back 539

electronic conductors such as metal wires (T)

Front 540

Electrodes

Back 540

electronic conductors (E)

Front 541

Ionic conductor

Back 541

salt bridge (I)

Front 542

Simple galvanic cell

Back 542

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

Front 543

cell potential (E)

Back 543

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

Front 544

RED CAT

Back 544

REDuction CAThode

Front 545

AN OX

Back 545

ANode OXidation

Front 546

Positive cell potential

Back 546

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)

Front 547

Reactions that do not occur at standard state

Back 547

deltaG = deltaG(not) + RT[ln(Q)]
deltaG: Gibbs free energy
deltaG(not): Gibbs free energy (standard conditions)
T: temperature
Q: reaction quotient

Front 548

reactions that are at equilibrium conditions

Back 548

at equilibrium, there is no available free energy with which to do work; deltaG = 0

deltaG(not) = -RT[ln(K)]

Front 549

relationship between K and deltaG(not)

Back 549

if K = 1, then deltaG(not) = 0

if K > 1, then deltaG(not) < 0

if K < 1, then deltaG(not) > 0

Front 550

Concentration cell

Back 550

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

Front 551

galvanic cell

Back 551

positive cell potential

spontaneous

Front 552

electrolytic cell

Back 552

negative cell potential

forced by outside power source to run backwards

cathode is negative
anode is positive

RED CAT & AN OX still the same