Chemistry 121 Formulas and Constants

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

U_rms 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

$$\dfrac{r_{1}}{r_{2}}=\sqrt{\dfrac{M_{2}}{M_{1}}}$$