Fahrenheit (F)
\dfrac{9}{5}\left( °C+32\right)
Celsius (C)
\dfrac{5}{9}\left( °F-32\right)
Kelvin (K)
°C+273
Density (d)
\dfrac{m}{V}
1 amu =
1.66054\times 11^{-24}g
Atomic weight
\Sigma [(mass)(% fractional abundance)]
Percentage to decimal
Divide by 100
Percent composition
% element =
[(# of atoms)(molecular weight)/ molar mass of the compound] x 100
Avogadro’s number
6.02\times 10^{23} atoms or molecules
Percent yield
(actual yield/theoretical yield) x 100
Empirical to molecular formula
given mass of molecular formula/ molar mass of empirical formula = multiplier
Molarity (M)
\dfrac{n}{V}
Dilution (before and after)
M_{1}V_{1}=M_{2}V_{2} or M_{c}V_{c}=M_{d}V_{d}
(c = concentrated, d = diluted)
If \Delta E >0
Endothermic
If \Delta E <0
Exothermic
Change in internal energy (\Delta E)
E_{f}-E_{0}
Exchange of energy betweem system and surroundings
\Delta E=q+\omega
q = heat
w = work
Work
PV or -PV
P = pressure
V = volume
Enthalpy (H)
E+PV
Change in enthalpy (\Delta H)
\Delta E+P\Delta V
Enthalpy of reaction
\Delta H_{rxn}=H_{products}+H_{reactants}
Calorimetry
q=mc\Delta T
Specific heat for H2O
4.184 J/g°C
Hess’s law
\Delta H=\sum products-\sum reactants
Bond enthalpy
\Sigma bonds broken - \Sigma bonds formed
Speed of light (constant)
c=3.00\times 10^{8}m/s
Speed of light (formula)
c=\lambda v
\lambda = wavelength
v = frequency
Energy is proportional to frequency
E=hv
h = plank’s constant
v = frequency
Plank’s constant (h)
6.626\times 10^{-34}J\cdot s
Dipole moment (\mu)
Qr
Q = charge
r = distance
Charge (Q)
1.60\times 10^{-19}C
C = Coulomb
1 D =
3.335\times 10^{-30}C\cdot m
1 A =
10^{-10}m
Formal charge
v.e. - lines - dots
6
octrahedral
Bond order
\dfrac{1}{2} (bonding e - antibonding e)
Pressure (P)
\dfrac{F}{A}
Pascals
1 Pa = 1 N/m²
Bar
1 bar = 105
Pa =
100 kPa
Atmosphere
1 atm = 760 torr = 760 mmHg
Hg =
101.325 kPa = 1.10325 bar
STP 1
1 atm
760 torr (mmHg)
101.325 kPa
Boyle’s law
P_{1}V_{1}=P_{2}V_{2}
Charles’s law
\dfrac{V_{1}}{T_{1}}=\dfrac{V_{2}}{T_{2}}
Gay-Lussac’s law
\dfrac{P_{1}}{T_{1}}=\dfrac{P_{2}}{T_{2}}
Combined gas law
\dfrac{P_{1}V_{1}}{T_{1}}=\dfrac{P_{2}V_{2}}{T_{2}}
Avogadro’s law
\dfrac{V_{1}}{n_{1}}=\dfrac{V_{2}}{n_{2}}
STP 2
V = 22.4 L
T = 0°C = 273 K
P = 1 atm
Ideal gas law
PV=nRT
R constant
0.08206 \dfrac{L\cdot atm}{mol\cdot K}
n
moles
M
molar mass
m
mass
Density of gases
d=\dfrac{MP}{RT}
Dalton’s law of partial pressures
P_{t}=P_{1}+P_{2}+P_{3}
Mole fraction
X_{1}=\dfrac{n_{1}}{n_{t}}
n_{1} = moles of compound (part)
n_{t} = total mass (whole)
Pressure and mole fraction
P_{t}=X_{2}P_{t}
X = mole fraction
Urms and molecular mass
u_{rms}=\sqrt{\dfrac{3RT}{M}}
R = 8.314 J/mol·K
R = 8.314 kg·m2/ s2·mol·K
Graham’s law describes diffusion and effusion
\dfrac{r_{1}}{r_{2}}=\sqrt{\dfrac{M_{2}}{M_{1}}}