General Chemistry

Front 1

Elements in Decreasing Electronegativity

Back 1

F > O > (N ≈ Cl) > Br > (I ≈ S ≈ C) > H

Front 2

Light Spectra, Longest wavelength to Shortest

Back 2

Radio, Micro, Infrared, Visible (ROYGBIV), Ultraviolet, X-ray, Gamma

Front 3

Energy Associated with a Photon:

Back 3

E(photon) = h*f = h*c/lambda

h = Planck's constant
f = Frequency
c = Speed of Light
lambda = wavelength

Front 4

Rules for Assigning Oxidation State

Back 4

1) ∑(oxidation states) = 0 in a neutral molecule
∑(oxidation states) = charge on ion

2) Group 1 metals: +1
Group 2 metals: +2

3) Oxidation state of F = -1

4) Oxidation state of H = +1

5) Oxidation state of O = -2

6) Oxidation state of Grp 6 = -2
Oxidation state of Grp 7 = -1

**Oxidation states do not have to be whole numbers!

Front 5

Radioactive Decay

Back 5

Alpha Decay: Decrease in neutrons and protons
Mass # - 4, Atomic # - 2

Beta- Decay: Decrease in neutrons, increase in protons
Atomic # + 1

Beta+ Decay: Decrease in protons, increase in neutrons
Atomic # - 1

Electron Capture: Decrease in protons, increase in neutrons
Atomic # - 1

Gamma decay: Brings excited nucleus to lower energy state
No change in both Atomic # and Mass #


**Number of protons in an atom does not determine what type of decay it undergoes, rather only the number of neutrons.

Front 6

Nuclear Binding Energy (NBE) equation:

Back 6

E(binding) = ∆m(in amu) * 931.5 MeV

Front 7

Atomic Radius Periodic Trend:

Back 7

Increases towards bottom left

**Increasing atomic radius facilitates increasing London Dispersion Forces

Front 8

Ionization Energy Periodic Trend:

Back 8

Increases towards top right

**Ionization energy: energy required to remove the least tightly bound electron from an isolated atom

Front 9

Electron Affinity Periodic Trend:

Back 9

Increases (more negative) towards top right

**Electron affinity is the energy associated with the addition of an electron to an isolated atom

Front 10

Electronegativity Periodic Trend:

Back 10

Increases toward the top right

(F > O > (N ≈ Cl) > Br > (I ≈ S ≈ C) > H)

Front 11

Acidity Periodic Trend:

Back 11

Increases towards the bottom left

Front 12

Formal Charge Equation:

Back 12

FC = V - (# of Bonds to other atoms) - (# Lone Pair electrons)

Front 13

Heat of Phase Transition Equation:

Back 13

q = n * ∆H(phase change)

q = amount of heat
n = number of moles
∆H(phase change) = heat of transition


**When a substances absorbs or releases heat, it can EITHER change temperature, OR undergo a phase change, BUT NOT BOTH AT THE SAME TIME

Front 14

Equation for Heat Absorption/Release based on Temperature:

Back 14

q = mc∆T

q = amount of heat
m = mass
c = specific heat of the substance (constant)
∆T = temperature change

**1 cal ≈ 4.2 J

Front 15

Celsius/Kelvin Conversion:

Back 15

temperature in K = temp in °C + 273.15

Front 16

Pressure Conversions:

Back 16

1 atm = 101.3 kPa = 760 torr = 760 mm Hg

Front 17

Ideal Gas Law Equations:

Back 17

PV = nRT

P = pressure (in atm)
V = volume (L)
n = # of moles
R = universal gas constant (0.0821 L-atm/K-mol)
T = absolute temperature (K)

(P1*V1) / T1 = (P2*V2) / T2
---> [at constant number of moles]

**At STP, 1 mol of (an ideal) gas ≈ 22.4 L
****High Pressure and Low temperature cause real gasses to deviate MOST from ideal-gas behavior

Front 18

Partial Pressures Equation:

Back 18

P(total) = X(a)*P + X(b)*P + X(c)*P

P = Total Pressure
X(a) = Mole fraction of gas A

Front 19

Graham's Law of Effusion:

Back 19

( [rate of effusion of Gas A] / [rate of effusion of Gas B]) = sqrt( [molar mass of Gas B] / [molar mass of Gas A] )

Front 20

Molarity:

Back 20

M = ( [number of moles of solute] / [liters of total solution] )

Front 21

Molality:

Back 21

m = ( [number of moles of solute] / [mass of solvent in kg] )

Front 22

Mole Fraction of Substance [ X(s) ]:

Back 22

X(s) = ( [number moles of substance S] / [total number of moles in solution] )

Front 23

Salt Solubility Rules:

Back 23

Soluble salts:
*Group I elements
*Ammonium (NH4+)
*Nitrate (NO3-)
*Perchlorate (ClO4-)
*Acetate (C2H3O2-)

Insoluble Salts:
*Silver (Ag+)
*Lead (Pb2+ / Pb4+)
*Mercury (Hg2 2+ / Hg 2+)
****These are insoluble, EXCEPT when as a Nitrate, Perchlorate, or Acetate (see above).

Front 24

Phase Solubility Trends:

Back 24

Solids in Liquids:
More soluble at higher temperature

Gases in Liquids:
Less soluble at higher temperature
More soluble at higher pressure

Front 25

Vapor Pressure Depression Equation:

Back 25

∆P(a) = -X(b) * P(a)

∆P(a) = Vapor Pressure Depression of Liquid A
X(b) = Mole fraction of Liquid B
P(a) = Vapor Pressure of Pure A

Front 26

Boiling-Point Elevation/Freezing-Point Depression Equation:

Back 26

∆T = k*i*m

∆T = Change in boiling/freezing point
k = boiling point elevation constant of the solvent
i = solute's van't Hoff factor
m = molality

**For water:
k(b) ≈ 0.5°C / m
k(f) ≈ 1.9°C / m

Front 27

Osmotic Pressure (Van't Hoff) Equation:

Back 27

∏ = M*i*R*T

∏ = Osmotic Pressure
M = Molarity
i = Van't Hoff factor
R = Universal gas constant (0.0821 L-atm/K-moL)
T = Absolute temperature (K)

Front 28

Solubility Product Constant Equation:

Back 28

K(sp) = [concentration of product A]^(# moles of A) * [concentration of product B]^(# moles of B)

Front 29

Reaction Rate:

Back 29

rate = k * [A]^x * [B]^y * [C]^z

**The overall order equals the sum of the reactant coefficients in the RDS of the reaction
****The rate law of an elementary step can be determined from the coefficients of the reactions in the elementary step.

Front 30

Arrhenius Equation:

Back 30

k = A*e^( -E[a] / R*T )

k = rate constant
A = Arrhenius factor (constant)
E[a] = Activation energy
R = 0.0821 L-atm/K-moL
T = Absolute temperature (K)

Front 31

Equilibrium Constant/Reaction Quotient:

Back 31

aA +bB --> cC + dD

= ( [C]^c * [D]^d ) / ( [A]^a * [B]^b )

**No pure solids or liquids are included
****This means that for the solubility product constant ( K[sp] ), there will be no denominator!

Front 32

Strong Acids:

Back 32

HI
HBr
HCl
HClO4
H2SO4
HNO3

Front 33

Strong Bases:

Back 33

Group I/Grp I oxides:
NaOH
LiOH/Li2O
KOH
RbOH

Group II:
Ca(OH)2
Sr(OH)2
Ba(OH)2

Metal amides:
NH2
H-

Front 34

General Acid/Base Equations:

Back 34

K(a) = ( [H+][A-] / [HA] )
K(b) = ( [OH-][HB+] / [B] )
K(w) = [H+][OH-] = K(a) * K(b) = 1 * 10^(-14) @ 25°C

pH = -log[H+] = -log[H3O+]
pOH = -log[OH-]
pH + pOH = 14 @ 25°C

Front 35

Henderson-Hasselbalch Equations:

Back 35

pH = pK(a) - log( [weak acid] / [conj. base] )
pOH = pK(b) - log( [weak base] / [conj. acid] )

Front 36

Acid/Base Neutralization Equation:

Back 36

N(a) * V(a) = N(b) * V(b)

Front 37

Laws of Thermodynamics:

Back 37

1) E(universe) is constant; E(system) = q + W

2) Spontaneous reaction --> ∆S(universe) > 0
**increase in entropy

3) S = O for pure crystal @ 0 K

Front 38

Gibbs Free Energy:

Back 38

∆G = ∆H - T∆S

∆G = ∆G° + R*T*ln(Q);
∆G° = -R*T*ln( K[eq] ) ≈ (-5.7 kJ/moL)*log( K[eq] )

∆G = -n*F*E(cell)

∆G = Gibbs Free Energy
∆H = Enthalpy
T = Temperature
∆S = Entropy
∆G° = G.F.E. for a rxn under standard conditions
R = Universal Gas Constant
F = Faraday ≈ 96,500 C/moL e-

Front 39

Nernst equation:

Back 39

E ≈ E° - ( [0.06] / n ) * log( Q )

E = Cell potential (electromotive force)
E° = Standard cell potential @ temp of interest
n = Number of moles of electrons
Q = reaction quotient