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92 Cards in this Set
- Front
- Back
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Atoms
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Tiny particles making up mass
Each atom is composed of a nucleus surrounded by one or more electrons. |
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Nucleus
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Contains protons & neutrons (collectively called: nucleons)
Held together by the strong nuclear force Surrounded by one or more electrons Radius = 10^-4 angstroms 1 angstrom = 10^-10 meters |
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Neutrons
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Part of nucleus
Together with proton, makes up the nucleons of the nucleus Approximately same size and mass as proton (1 amu) Slightly heavier than proton No charge, electrically neutral |
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Protons
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Part of nucleus
Together with neutron, makes up the nucleons of the nucleus Same size and mass as neutron (1 amu) Slightly lighter than neutron Positive charge (1+) |
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Electrons
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Surround nucleus at distance of 1 to 3 Angstroms
Over 1800 times lighter than mass of a nucleon Electrons (1-) and protons (1+) have opposite charges of equal magnitude |
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Electronic charge (e)
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Charge of one electron (1 e)
e = 1.6e^-19 coulombs (C) |
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Atom
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Electrically neutral
Made up of neutrons, protons and electrons Same number of protons as electrons Composed mostly of empty space |
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Elements
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A single atom
Building blocks of all compounds Cannot be decomposed into simple substances by chemical means Characterized by: 1. Mass number (A) 2. Atomic number (Z) 3. Atomic weight (amu) |
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Mass number (A)
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Number of protons plus neutrons
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Atomic number (Z)
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number of protons
Identity number of any element |
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Isotopes
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Two or more atoms of the same element that contain different numbers of neutrons
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Atomic Weight
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Also known as molar mass (MM or M)
Given in atomic mass units (amu or u) or grams/mole (g/mol) Actually a mass (ratio) and not a weight |
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Atomic mass units (amu)
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An amu is defined by carbon-12
1 atom of C-12 has an atomic weight of 12 amu All other atomic weights are measure against this standard Also known as a dalton |
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Mole
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Defined by C-12
Also known as Avogadro's number = 6.022e^23 The number of C atoms in 12 grams of C-12 6.022e^23 amu = 1 gram Moles = g/amu |
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Periodic table
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Lists the elements from left to right in order of their atomic numbers
Can be divided into: 1. Nonmetals (right) 2. Metals (left) 3. Metalloids (diagonal seperation between metals & nonmetals) |
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Period
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Each horizontal row of periodic table
Elements in the same family on the periodic table tend to have similar chemical properties Tend to make the same number of bonds Tend to exist as similarly charges ions |
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Groups or Families
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Vertical columns of periodic table
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Metals
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Large atoms that tend to lose electrons to form positive ions or form positive oxidation states
Metallic character (easy movement of electrons) All metals (except mercury) exists as solids at room temperature Form ionic oxides (ie: BaO) |
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Metallic Character
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Property of metals
Increases from right to left and top to bottom of periodic table 1. Ductility (easily stretched) 2. Malleability (easily hammered into think strips) 3. Thermal and electrical conductivity 4. Characteristic luster |
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Nonmentals
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Diverse appearances and chemical behaviors
Lower melting points than metals Form negative ions Make up molecular substances Form covalent oxides (ie: SiO2 or CO2) |
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Alkali metals
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Group (family) 1A in periodic table
metals Soft metallic solids Low densities and melting points Easily form 1+ cations Highly reactive, reacting with most nonmetals to form ionic compounds |
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Alkaline earth metals
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Group (family) 2A in periodic table
metals Harder, more dense and melt at higher temperatures than alkali metals Form 2+ cations Less reactive than alkali metals The heavier, the more reactive |
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Halogens
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Group (family) 17 in periodic table
Nonmetals & metalloids |
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Noble gases
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Groups (family) 18 in periodic table
Nonmetals Also known as rare gases Nonreactive, inert gases at room temperature |
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Metalloids
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Characteristics that resemble metals and nonmetals
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Representative or main-group elements
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Section A groups in periodic table
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Transition metals
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Section B groups in periodic table
When form ions, they lose electons from s-subshell first and then from d-subshell |
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Hydrogen
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Is unique and chemical/physical properties to do conform to any family
Nonmetal Colorless Odorless Diatomic gas |
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Group 4A
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Elements can form 4 covalent bonds with nonmetals
All (except carbon) can form 2 additional bonds with lewis bases Only carbon forms strong pi bonds to make strong double and triple bonds |
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Group 5A
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Elements can form 3 covalent bonds
All (except nitrogen) can form 5 covalent bonds by using d-orbitals Can bond with lewis base to form 6th covalent bond Nitrogen forms strong pi bonds to make double and triple bonds |
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Group 6A
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Elements called chalcogen
Oxygen and sulfur are most important Oxygen is second most electronegative element, divalent, can form strong pi bonds to make double bonds, reacts with metals to form oxides Sulfur can form 2, 3, 4 or 6 bonds and can pi bond to make double bonds |
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Group 7A
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Radioactively stable called halogens
1.Fluorine 2. Chlorine 3. Bromine 4. Iodine Highly reactive, like to gain electrons Fluorine makes only 1 bond, while other halogens can make more than 1 bond Bind to hydrogen to form hydrogen halides (soluble in water) Reacts with metals to form ionic halides |
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Diatomic molecules
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1. Hydrogen
2. Oxygen 3. Nitrogen 4. Halogens |
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Small atoms
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Make strong pi bonds due to overlap of p-orbitals
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Large atoms
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Make weak or are unable to make pi bonds due to lack of overlap of p-orbitals
Have d-orbitals allowing for more than 4 bonds |
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Pi bonds
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Allow for double and triple bonds
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Ion
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When element has more or fewer electrons than protons
Representative elements make ions by forming the closest noble gas electron configuration Made from metals and nonmetals |
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Cation
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Positive ion
Formed by metals Significantly smaller than neutral atom counterparts |
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Anion
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Negative ion
Formed by nonmetals Much larger than neutral atom counterpart |
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Predict ion charge based on:
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1. Atoms lose electrons from higher energy shell first
In transition metals, this means electrons are lost from s-subshell first and then from d-subshell 2. Ions are looking for symmetry Representative elements form noble gas electron configurations when they may ions Transition metals try to "even-out" their d-orbitals, so each orbital has the same number of electrons |
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Electron shielding
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1st electron shields nuclear charge from 2nd electron, so that 2nd electron doesn't feel entire nuclear charge
Instead, 2nd electron feels an effective nuclear charge |
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Effective nuclear charge (Zeff)
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Amount of charge felt by 2nd electron due to 1st electron shielding of nuclear charge
Zeff = nuclear charge (Z) - average # of electrons between nucleus and electron in question What should be plugged in to: F = Kqq/r^2 Increasing going left to right and top to bottom on periodic table |
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Periodic trends
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1. Atomic radius
2. Ionization energy 3. Electronegativity 4. Electron affinity 5. Metallic character |
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Atomic radius
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Since Zeff increases when moving left to right, each additional electron is pulled more strongly toward nucleus, resulting in a small atomic radius
Increases from top to bottom and right to left |
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Ionization energy
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Energy necessary to detach an electron from a nucleus
1st ionization energy = energy required to detach an electron from a neutral atom 2nd ionization energy = energy required to detach a 2nd electron from same atom 2nd ionization energy > 1st ionization energy because when electron is removed, Zeff on other electrons increases Increases from left to right and bottom to top of periodic table (explained by Zeff) |
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Electronegativity
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Tendency of an atom to attract an electron in a bond that it shared with another atom
Increases from left to right and bottom to top of periodic table Related to Zeff in similar way as ionization energy Undefined for noble gases |
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Electron affinity
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Willingness of an atom to accept an additional electron
Energy released when an electron is added to a gaseous atom Increases left to right and bottom to top of periodic table Related to Zeff Electron affinity is more exothermic to right and up on periodic table Endothermic for noble gases |
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SI units
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Mass = kg
Length = m Time = s Electric current = A Temperature = K Luminous intensity = cd Amount of substance = mol Mega (M) = 10^6 Kilo (k) = 10^3 Deci (d) = 10^-1 Centi (c) = 10^-2 Milli (m) = 10^-3 Micro (u) = 10^-6 Nano (n) = 10^-9 Pico (p) = 10^-12 Femto (f) = 10^-15 |
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Bonds
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What holds atoms together
2 types: 1. Covalent bonds 2. Ionic bonds 2 atoms will form a bond if they can lower their overall energy level by doing so Nature seeks lowest energy state Energy is always required to break a bond, no energy is every released by breaking a bond |
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Covalent bonds
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2 electrons are shared by 1 nuclei
Negatively charges electrons are pulled toward both positively charged nuclei by electrostatic forces |
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Bond length
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Point where energy level is lowest
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Bond dissociation energy (bond energy)
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Energy necessary to achieve a complete separation of atoms
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Compound
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Substance made from 2 or more elements
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Empirical formula
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In pure compounds, relative number of atoms of 1 element to another can be represented by a ratio
Glucose = CH2O To find empirical formula from percent mass: Compound = 6% H & 94% O by mass Assume 100g of sample (6g H)/(1g/mol) = 6mol (94g O)/(16g/mol) = 5.9mol = 6 (must be whole #s) 6/6 = 1 Empirical formula = HO |
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Molecules
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Separate and distinct units, in molecular compounds, formed from repeated groups of atoms
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Molecular formula
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Exact number of elemental atoms in each molecule in a molecular compound
Glucose = C6H12O6 |
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Percent mass
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Calculated from empirical formula and atomic weight of each atom
Ex: Percent mass of Carbon in Glucose (CH2O) (molecular weight C)/(molecular weight of CH2O) = 12/30 = 0.4 0.4 x 100 = 40% Glucose is 40% carbon by mass |
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Ionic compounds
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Named after their cation and anion
Put cation name in from on anion name (barium sulfate, BaSO4; sodium hydride, NaH) |
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Cation nomenclature
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Metal cation:
1. Roman number in parentheses indicating charge [copper(I) = +1 or copper (II) = +2] 2. -ic greater charge (cupric, Cu2+) or -ous smaller charge (cuprous, Cu+) Nonmetal cation: 1. cation name ends in -ium (ammonium, NH4+) |
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Anion nomenclature
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1. -ide after anion (hydride ion, H-; hydroxide, OH-)
2. polyatomic anions with multiple oxygens end with -ite (less oxygenated) or -ate (more oxygenated) depending on relative # of oxygens (nitrite ion, NO2-; nitrate ion, NO3-) 3. More oxygens represented by hypo- (fewest oxygens) or per- (most oxygens) prefixes (hypochloride, ClO-; chlorite, ClO2-; chlorate, ClO3-; perchlorate, ClO4-) 4. If oxyanion has a hydrogen, word hydrogen is added (hydrogen carbonate ion, HCO3-) |
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Acid nomenclature
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Named based on their anions
1. If ends in -ide, name starts with hydro- and ends in -ic (hydrosulfuric acid, H2S) 2. If an oxyacid, ending -ic (more oxygen) and -ous (less oxygens) (sulfuric acid, H2SO4; sulfurous acid, H2SO3) |
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Binary molecular compounds
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Compounds with only 2 elements
Beings with name of element farthest to left and lowest in periodic table Name of 2nd element is given suffix -ide and greek # prefix is used on 1st element if necessary Ex: dinitrogen teroxide, N2O4 |
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Physical reaction
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When compound undergoes a reaction and maintains its molecular structure and this its identity
Ex: Melting, evaporation, dissolution and rotation of polarized light |
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Chemical reaction
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When a compound undergoes a reaction and changes its molecular structure to form a new compound
Ex: Combustion, metathesis and redox Can be represented by a chemical equation with the molecular formulae of the reactants on the left and products on the right Ex: CH4 + 2O2 --> CO2 + 2H2O Coefficients indicate the relative number of molecules The atoms are always conserved, the equation is balanced |
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Reaction runs to completion
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Reaction from to the right until at least one of the reactants is depleted
Often reactions do not run to completion because they reach equilibrium first |
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Limiting reagent
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Reactant which is depleted first if reaction runs to completion
Not necessarily the reactant of which there is least |
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Theoretical yield
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Amount of product produced when a reaction runs to completion
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Actual yield
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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 |
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Percent yield
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[(Actual yield)/(Theoretical yield)] x100
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Reaction types
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1. Combination
2. Decomposition 3. Single displacement 4. Double displacement 5. Redox 6. Combustion 7. Bronsted-Lowry Acid-base 8. Lewis Acid-base |
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Combination reaction
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A + B --> C
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Decomposition reaction
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C --> A + B
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Single Displacement reaction
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A + BC --> B + AC
Also called single replacement |
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Double Displacement reaction
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AB + CD --> AC + BD
Also called double replacement or metathesis |
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Quantum mechanics
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Elementary particles can only gain or lose energy and certain other quantities in discrete units
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Principal quantum number (n)
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First quantum number
Shell level Larger n the greater the size and energy of the electron orbital Representative elements: n for electrons in the outer most shell is given by the period in the periodic table Transition metals: n lags 1 shell behind the period Lanthanides & actinides: n lags 2 shells behind the period |
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Valence electrons
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Electrons which contribute most to an element's chemical properties
Located in outermost shell of atom Only electrons from s & p subshells are considered valence electrons |
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Azimuthal quantum number (l)
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Second quantum number
Designate subshells: Orbital shapes such as s, p, d & f l = 0 = s subshell l = 1 = p subshell l = n-1; for each new shell (n) there exists an additional subshell |
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Magnetic quantum number (ml)
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3rd quantum number
Designates the precise orbital of a given subshell Each subshell will have orbitals with ml from -l to +l 1st shell, n = 1, l = 0, ml = 0 n = 3, l = 2, ml = 5 (-2, -1, 0, +1, +2) |
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Electron spin quantum number (ms)
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4th quantum number
Can values of -1/2 or +1/2 Any orbital can hold up to 2 electrons If 2 electrons occupy the same orbital, they have the same first 3 quantum numbers |
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Pauli Exclusion Principle
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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 |
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Number of orbitals within a shell
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n^2
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Heisenberg Uncertainty Principle
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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 |
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Aufbau Principle
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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 |
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Electron configuration
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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 |
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Abbreviated electron configuration
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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 |
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Ground state
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atom whose electrons are all their lowest energy levels
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Electron configuration for ions
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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 |
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Hund's Rule
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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 |
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Planck's Quantum Theory
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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 |
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Photoelectric effect
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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 |
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Work function (ϕ)
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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 |
