Polyatomic Ions

Cations

Aluminum: Al(3+)

Ammonium: NH(4+)

Barium: Ba(2+)

Calcium: Ca(2+)

Copper(I): Cu(+)

Copper(II): Cu(2+)

Hydrogen: H(+)

Hydronium: H3O(+)

Iron(II): Fe(2+)

Iron(III): Fe(3+)

Lead: Pb(2+)

Lithium: Li(+)

Magnesium: Mg(2+)

Mercury(I): Hg2(2+) (diatomic cation)

Tin(II): Sn(2+)

Tin(IV): SN(4+)

Potassium: K(+)

Silver: Ag(+)

Sodium: Na(+)

Anions

Acetate: C2H3O2(-) or CH3COO(-)

Carbonate: CO3(2-)

Hydrogen: carbonate HCO(3-)

Chloride: Cl(-)

Hypochlorite: ClO(-)

Chlorite: ClO2 (-)

Chlorate: ClO3 (-)

Perchlorate: ClO4 (-)

Chromate: CrO4 (2-) (Cr Oxidation number +6)

Dichromate: Cr2O7 (2-) (Cr Oxidation number +6)

Cyanide: CN(-)

Thiocyanate: SCN(-)

Hydride: H(-)

Hydroxide: OH(-)

Nitride: N(3-)

Nitrite: NO2(-)

Nitrate: NO3(-) (N oxidation number +5)

Oxalate: C2O2(2-)

Oxide: O(2-)

Peroxide: O2(2-)

Permanganate: MnO4(-) (Mn oxidation number +7)

Phosphide: P(3-)

Phosphite: PO3(3-)

Phosphate: PO4(3-)

Sulfide: S(2-)

Sulfite: SO3(2-)

Sulfate: SO4(2-)

Hydrogen sulfate: HSO4(-)

EDTA: EDTA(4-)

Acids and Bases

Strong Acids and Bases

Hydrochloric acid: HCl

Nitric acid: HNO3

Sulfuric acid: H2SO4

Hydrobromic acid: HBr

Hydroiodic acid: HI

Perchloric acid: HClO4

Hydroxide: OH(-)

Acids and Bases

Hydronium: (acid) H3O(+)

Ammonia: (weak base) NH3

Ammonium: (weak conjugate acid) NH4(+)

Acetate: (weak conjugate base) C2H3O2(-) or CH3COO(-)

Elements

Alkali Metals

Lithium: Li, atomic number 3, average atomic mass 7, period 2

Sodium: Na, atomic number 11, average atomic mass 23.0, period 3

Potassium: K, atomic number 19, average atomic mass 39.1, period 4

Alkaline Earth Metals

Magnesium: Mg, atomic number 12, average atomic mass 24.3, period 3

Calcium: Ca, atomic number 20, average atomic mass 40.1, period 4

Metals

Copper: Cu, atomic number 29, average atomic mass 64, period 4

Zinc: Zn, atomic number 30, average atomic mass 65, period 4

Silver: Ag, atomic number 47, average atomic mass 108, period 5

Nonmetals

Carbon: C, atomic number 6, average atomic mass 12.0, period 2

Nitrogen: N, atomic number 7, average atomic mass 14.0, period 2

Oxygen: O, atomic number 8, average atomic mass 16.0, period 2

Fluorine: F, atomic number 9, average atomic mass 19.0, period 2

Aluminium: Al, atomic number 13, average atomic mass 27.0, period 3

Silicon: Si, atomic number 14, average atomic mass 28, period 3

Phosphorus: Mg, atomic number 15, average atomic mass 31, period 3

Sulfur: S, atomic number 16, average atomic mass 32.1, period 3

Chlorine: Cl, atomic number 17, average atomic mass 35.5, period 3

Noble Gases

Hydrogen: H, atomic number 1, average atomic mass 1.01, period 1

Helium: He, atomic number 2, average atomic mass 4, period 1

Neon: Ne, atomic number 10, average atomic mass 20, period 2

Argon: Ar, atomic number 18, average atomic mass 40, period 3

Diatomic Elements

H2: single covalent bond B.E. 450 kJ/mp;

F2, Cl2, Br2, I2: single covalent bond B.E 150-250 kJ/mol

O2: double covalent bond B.E. 500 kJ/mol

N2: triple covalent bond B.E. 900 kJ/mol

Molecular Mass of Molecules

H2 = 2.0 g/mol

N2 = 28 g/mol

O2 = 32 g/mol

H2O = 18 g/mol

CO2 = 44 g/mol

NaOH = 40.0 g/mol

AP02

from least to most accurate: beaker, erlenmeyer flask, graduate cylinder, balance, volumetric flask, burette, pipette

Mass spectrometer: A mass spectrometer sorts isotopes of elements by mass and can shows the relative abundance of each isotope.

AP03

Avogadro’s number = 6.02 x 10^23 particles/mole

AP04

soluble ions: Na(+), K(+), NH4(+), NO3(-)

Molecular equations: full equation

Ionic equations: soluble substances shown as ions

Net ionic equations: spectator ions omitted (weak acids and bases shown as molecules)

AP05

Gas Law: PV=nRT

molar gas volume at STP = 22.4 L per mol

molas mass = DRT/P (DiRTy over P)

Molecular speed of gases: lighter gases are faster at a given temp

non-ideal conditions: high pressure and low temperature (intermolecular attractions interfere)

higher than expected pressure: large molecular volume

lower than expected pressure: high intermolecular attractions

AP06

specific heat: change in temp with addition of heat (water= high, metal=low)

Calorimeter: q = mass x ∆T x specific heat

enthalpy from calorimeter: ∆H = -q/mol

endothermic rxn: positive ∆H, decrease in BE, increase in PE, stronger bonds reactants, weaker bonds products, thermodynamically unfavored, heat is reactant

exothermic rxn: negative ∆H, increase in BE, decrease in PE, weaker bonds reactants, stronger bonds products, thermodynamically favored, heat is product

AP07

shorter wavelengths = higher frequencies and greater photon energies

Energy of a photon: E = hc/(λ x 10^-9)

X-ray wavelength = 10 nm (ionize inner electrons)

UV wavelength = 10-400 nm (ionize outer electrons)

Visible light wavelength = 400 (violet) - 700 (red) nm (excite outer electrons)

Infrared Radiation (IR) wavelength = 700 nm - 1 mm (molecular rotations/vibrations)

order of orbital filling = 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2

order of ionization = 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2

Electron pairing, electrons will half fill orbitals in a subshell before pairing, must have opposite spins, easier to ionize paired atoms

quantum numbers = provide more detailed description of electrons in an atom

valence electrons of transition metals = 2, 1 for Cr and Cu

hydrogen bonding: F, O, H

Ionization energy = always endothermic

AP08

lattice energy = KE release when ions form crystal lattice (more KE = stronger bond)

morse diagram: well = BE and bond distance

∆H = -∆BE

Formal charge = difference between normal number of valence e- and number of electrons controlled by atom in a molecule (lowest is most likely, or most electronegative will have - charge)

electron domain = unshared pairs of electrons + sigma bonds

6 pairs of electrons = Boron and Aluminum

more than 4 domains = electrons in d orbitals (4th, 5th, 6th, 7th periods)

Symmetrical molecular shapes: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, square planar

unsymmetrical molecular shapes: nonlinear, trigonal pyramidal, t-shaped, distorted tetrahedron, square pyramidal

2 domain hybridization = sp

3 domain hybridization = sp2

4 domain hybridization = sp3

Lewis Dot structure of H2O: 4 electron domains, 2 bonded, nonlinear, 105˚, sp3

Lewis Dot structure of CO2, 2 electron domains, 2 bonded, 2 sigma bonds, 2 pi bonds, linear, 180˚, sp

Lewis Dot structure of NH3: 4 electron domains, 3 bonded, trigonal pyramidal, 105˚, sp3

AP09

intermolecular attractions: LDF, dipole-dipole, dipole-ion, h bonding

high intermolecular attractions mean: higher freezing points, lower vapor pressures (higher boiling temps), higher enthalpies of fusion and vaporization

molecular solid: low melting temp and no electrical conductivity

ionic solid: high melting temp and electrical conductivity only when molten

metallic solid: malleable, electrical conductivity solid and molten (decreases when heated), mobile valence e-s

interstitial alloys: in between lattice of metals, smaller atoms, decrease malleability and increase strength and hardness

substitutional alloys: replacing atoms in metal lattice, same size atom, lower malleability

network solid: high melting temp,

3-D covalent networks: rigid and hard

2-D covalent networks: flexible and soft

fusion: s to l, ∆H and ∆S = +

vaporization: l to g, ∆H and ∆S = +

sublimation: s to g, ∆H and ∆S = +

freezing: l to s, ∆H and ∆S = -

condensation: g to l, ∆H and ∆S = -

AP10

alkanes, alkenes, and alkynes: only C and H, no dipoles, only LDF, not soluble in water

alkane bonds: C-C, single bond, sp3

alkene bonds: C=C, double bond, sp2

alkane bonds: C≡C, triple bond, sp

alcohols: -COH hydrogen bonding and LDFs, smaller chains soluble in water

carboxylic acids: -COOH, h bonding and LDFs

amines: -CNH2, h-bonding and LDFs

aldehydes, esters, ethers, and ketones: −CH=O, -COO-, -COC-, -C=O; dipole-dipole and LDFs, small chains dissolve in water

polymer: (plastics) long-chain carbon chains with repeating units (mer)

one carbon prefix: meth-(form)

two carbon prefix: eth-(acet)

three carbon prefix: prop-

four carbon prefix: but-

five carbon prefix: pent-

six carbon prefix: hex-

seven carbon prefix: hept-

eight carbon prefix: oct-

nine carbon prefix: non-

ten carbon prefix: dec-

AP11

what types of molecular substances will dissolve in water: strong dipoles/-OH and low LDFs

what types of ionic substances will dissolve in water: sodium, potassium, ammonium, and nitrate ions, or if ion-dipole > cation-anion attractions

volatile: higher vapor pressure, lower intermolecular attraction

distillation: used to separate solutions by evaporating most volatile components

distillate: condensed vapor

chromatography: method of separating small quantities of components of a mixture using differences in intermolecular attractions

retention factor: component distance ÷ solvent distance

colorimeter/spectrophotometer: uses absorption of light to determine concentration of solution

%transmittance: % of a specific wavelength that makes it through a solution

absorbance: how much of a specific wavelength of light is absorbed

wavelength of light used in spectrophotometry: set of wavelengths absorbed most strongly by solution

concentration vs absorbance: directly proportional

AP12

rate units: M/s

rate law expression for A+B→C: rate = k[A]^x[B]^y

zero order reactions: rate = k[A]^0, half life increases with time, [A] vs t is linear, k = -∆[A]/∆t

first order reactions: rate = k[A]^1, half life is constant, [A] vs t is linear is a slight curve, ln[A] vs t is linear, k = 0.693/half life = -∆ln[A]/∆t

second order reactions: rate = k[A]^2, half life decreases with time, [A] vs t is linear is a sharper curve, 1/[A] vs t is linear, k = ∆(1/[A])/∆t

two factors required for successful activated complex collision: 1. collision must have sufficient energy 2. collision must have proper orientation

what does a catalyst do: lowers activation E of forward and reverse rxn, reaches equilibrium faster

what doesn’t a catalyst do: doesn’t change K, ∆G, or ∆H

catalyst: present as a reactant and produced as a product, may be included in rate law

slow step: determines reactants in rate law, highest bump (Ea)

how does equilibrium rxn affect rxn rate: preceding slow step results in reactants included in rate law expression

two reasons that rxn rates increase with temp: 1. more frequent collisions 2. greater percentage of collisions have energy to reach Ea and form activated complex

AP13

equilibrium constant: [products]/[reactants]

Kc = Kp when: mole gas of reactant = mol gas product

K in exothermic rxn: K decreases with increases in temp

K in endothermic rxn: K increases with increases in temp

Q>K: forward rxn

Q<K: reverse rxn

Q=K: equilibrium reached

K for added equations: K1 x K2

Equilibrium for reverse rxn: Kreverse = 1/Kforward

solubility product expression: Ksp = [A(y+)]^x[B(x-)]^y

common ion: lowers molar solubility

AP14

polyprotic acid: can donate >1 proton, Ka = product of each proton’s Ka

amphiprotic subtance: can donate or accept a proton, act as acid or base

neutral water: [H+] = [OH-] and pH = pOH

Kw at 25˚C = 1E-14 = [H+][OH-]

Kw and pKw as temp increases: Kw increase and pKw decreases

brønsted acid: proton donator, turns into brønsted base

brønsted base: proton acceptor, turns into brønsted acid

Ka and Kb relationship between acid and conjugate base: 1E-14 = Ka x Kb

molecular rxn for strong acid and base: HA(aq) + BOH(aq) → H2O(l) + AB(aq)

ionic rxn for strong acid and base: H+ + A- + OH- → H2O(l) + A- + B+

net ionic rxn for strong acid and base: H+ + OH- → H2O(l)

net ionic rxn for strong acid and weak base: H+ + B → HB+

net ionic rxn for weak acid and strong base: HA + OH- → A- + H2O(l)

%ionization: [H+]/[HA] x 100, increases with dilution

approximation [H+] of weak acid from [HA]: [H+] ≈ √[HA] x Ka

approximation [OH-] of weak acid from [B]: [OH-] ≈ √[B] x Kb

Equivalence point of strong acid and base: pH = 7

Equivalence point of strong acid and weak base: pH < 7 because of weak conjugate acid

Equivalence point of weak acid and strong base: pH > 7 because of weak conjugate base

Half equivalence point: [HA] = [A-], Ka = [H+], pH = pKa

when [HA] > [A]: more acidic buffer [H+] > Ka

when [HA] < [A]: less acidic buffer [H+] < Ka

strong acid titrated with strong base: start at low pH, end at high pH

weak acids/bases titration: buffer softens curve

diprotic acid titration curve: two equivalence and half titration points

AP15

S: entropy, measure of dispersion of matter and energy

S of phases: s < l < g and usually s < aq

rxns that increase S: rxns that break up compounds/produce gases

units of S: J/K

∆G˚: standard Gibbs free energy, determines if rxn is thermodynamically favored based on 1 molar 1 atm, indicates how much work can be done as rxn progresses to equilibrium

∆G˚ units: kJ/mol

if ∆G˚ = -20kJ/mol then K will be very large and rxn goes to completion

if ∆G˚ = +20kJ/mol then K will be very small and no significant amt of product

AP16

Oxidation number of F: -1

oxidation number of halogens: -1, unless bonded to atom with > electronegativity

oxidation number of O: -2, unless in peroxide, -1

oxidation number of H: +1, unless as a hydride, -1

oxidation number of alkali metals: +1

oxidation number of alkaline earth metals: +2

oxidation number of Al: +3

oxidation: loss of electrons at anode

reduction: gain of electrons at cathode

flow of electrons: towards cathode

cathode: site of reduction, any nonreactive conductive substance, surface for Y+ to Y

anode: site of oxidation, metal erodes into solution

salt bridge: allows movement of ions to balance charges, A- goes to anode, B+ goes to cathode

standard voltage for hydrogen reduction: 0

electrolytic cell mol X = amps x seconds / 96485|ion charge|