Chemistry 121 Formulas and Constants (copy) (copy)

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if ΔE > 0

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if ΔE > 0

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if ΔE > 0

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Farenheit (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}}}