DAT General Chemistry

Hydroxide

OH^-

Nitrate

NO₃^-

Carbonate

CO₃^2-

Sulfate

SO₄^2-

Phosphate

PO₄^3-

Ammonium

NH₄⁺

Methane

CH₄

Ammonia

NH₃

Hydrogen Peroxide

H₂O₂

accuracy

closeness of data value to literature value

precision

consistency of all data values

log 2

0.3

log 3

0.48

1 Liter = ? mL and ? cm³

1000 mL and 1000 cm³

t (K) =

t (C) + 273.15 (K)

Pressure: 1 atm = ? torr / mmHg, ? kPa, and ? bar

760 torr / mmHg, 101.3 kPa, and 1.013 bar

Boiling Point of water

100 C

Freezing Point of water

0 C

Amount of substance

mole (mol)

thermodynamic temperature

kelvin (K)

mass

kilogram (kg)

length

meter (m)

time

second (s)

giga-

G = 10⁹

mega

M = 10⁶

kilo

k = 10³

hecto

h = 10²

deca

da = 10¹

deci

d = 10^-1

centi

c = 10^-2

milli

m = 10^-3

micro

u = 10^-6

nano

n = 10^-9

Have No Fear Of Ice Cold Beer

Diatomic gases: H2, N2, F2, O2, I2, Cl2, Br2

Metal Properties: (6)

• Malleable and lustrous

• Form basic oxides

• High melting and boiling points

• Lose electrons to form cations

• Good conductors of electricity 

• Solid at room temperature (exception: mercury)

Non-metal Properties: (6)

• Brittle and dull

• Form acidic oxides

• Low melting and boiling points

• Gain electrons to form anions

• Poor conductors of electricity and heat 

• Gas or solid at room temperature (exception: bromine) 

Atomic radius

½ the bond length between identical atoms;

increases to the left and down the periodic table

Effective Nuclear Charge

Z_eff = Z (proton #) - S (# of shielding electrons)

= the amount of positive charge experience by an electron;

increasing to the right and down the periodic table

isoelectronic series

atoms/ions with the same number of electrons but different proton numbers; anions have larger atomic radius, cations have smaller atomic radius

ionization energy/potential

the energy needed to remove an electron from a gaseous, neutral atom… = change in enthalpy (delta H)

increases to the right and up (1st ionization energy)

• 2nd ionization energy: after 1st electron is removed, the 2nd ionization is harder and higher energy; easier for 1st group and smaller elements than group 2, but still a high energy;

• alkaline earth metals have a higher than usual 1st ionization energy because of p orbital valence electrons;

• group 15 higher energy than group 16 because of half filled orbitals

• group 2 have higher energy than group 13 because of group 2 filled orbital

electron affinity

the amount of energy released/required when an electron is added to a gas-phase atom

increase to the right and up; becomes more negative

• noble gases have low affinity (they have all of their valence electrons);

• if an element can perfectly half fill their orbital with the electron, it will have a higher affinity than an adjacent element with an already half-filled orbital (group 4 > group 5)

electronegativity

the ability of an atom to attract electrons to itself in a covalent bond

increase to the right and up

• noble gases are negligible

Soluble Salts (5)

• group 1 metal cations

• NO₃^-

• ClO₄^-

• C₂H₃O₂^-

• NH₄⁺

Insoluble Salts (7)

• Ag⁺

• Pb²⁺

• Hg²⁺

• OH^-

• S^2-

• CO₃^2-

• PO₄^3-

Strong Acids: (7)

• HCl

• HBr

• HI

• HClO₄

• HClO₃

• H₂SO₄

• HNO₃

Strong Bases (4)

• Group 1 Metal Hydroxides

• Ba(OH)₂

• Sr(OH)₂

• Ca(OH)₂

Weak Acids: (3)

• HF

• HCN

• CH₃COOH

Weak Base: (2)

• NH₃

• CH₃NH₂

Colligative properties

depend upon the concentration of solute molecules or ions, but not the identity of the solute

ex: salt makes water harder to freeze

Non-colligative properties

depend on the identity of the dissolved species and the solvent

ex: surface tension varies between substances, color depends on substances

Surface Tension

property of a liquid that allows it to resist an external force, due to cohesion

non-colligative

Viscosity

measure of a liquids resistance to flow

non-colligative

Solubility

ability of a solute to for a solution with a solvent

non-colligative

Color and density

do not depend on the number of particles

non-colligative

Boiling Point

when the vapor pressure = atmospheric pressure; liquid is converted into gas

colligative

BP at sea level

100 C

BP above sea level

decreases; as elevation increases the atmospheric pressure decreases and this BP is easier to reach

BP below sea level

increases; as elevation decreases the atmospheric pressure increases because the air is denser

Vapor-Pressure Depression

vapor pressure without solute is higher than VP with a non-volatile solute; adding solute increases boiling point by lowering VP and requiring higher temp to reach atmospheric pressure; Raoult’s Law

colligative

Raoult’s Law (Vapor Pressure)

VP of solution = mole fraction of solvent x VP of solvent

colligative

Boiling Point Elevation

BP of a liquid will be higher when a non-volatile solute is added; large van’t Hoff factor increases BP

Change in BP = solvent’s BP-elevation constant x i x molal

colligative

Freezing-Point Depression

a drop in the minimum temp at which a substance freezes when a non-volatile solute is added; solution has a lower FP due to the solute

ex: alcohol lowers the FP of water

Change in FP = (-) solvent’s FP-depression constant x i x molal

colligative

Osmotic Pressure

the pressure required to stop osmosis

Osmotic Pressure (atm) = i x M x universal gas constant x temp (K)

colligative

van’t Hoff Factor

“i”

the number of particles a solute breaks down into

ex: MgCl2 … i = 3 because of Mg²⁺, Cl^-, and Cl^-

Normality

N= number of equivalents x M

Combined Gas Law

(P1V1)/(n1T1) = (P2V2)/(n2T2)

Boyle’s Law

P1 V1 = P2 V2

Charles’s Law

V1/T1 = V2/T2

Avogado’s Law

V1/n1 = V2/n2 where n is the number of moles

Gay-Lussac’s Law

P1/T1 = P2/T2

Dalton’s Law of Partial Pressure

total pressure of the mixture is equal to the sum of the partial pressures of individual gases

• P_total= P1 + P2 + …

• P1 = X1 x P_total where X1 is the mole fraction of gas 1

Gas density

Density = m / v

Density = (P x M) / (R x T)

Effusion

A confined gas escaping through a small hole, which gas effused faster are ones with lower molecular weight

Graham’s Law of Effusion

r1/r2 = √(M2/M1) where r is the effusion rate and M is the molar mass

inversely proportional

Ideal Gas Law

PV = nRT

STP

Standard Temp and Pressure: 273K, 1 atm, 1 mol = 22.4 L

1st Order Half-life

t_1/2= 0.693/k

Properties of Work (2)

1. Delta E = q + w (change in internal energy = heat + work)

2. w = -P x delta V (work = -Pressure x change in volume)

Arrhenius Acid

produces H⁺ ions

Arrhenius Base

produces OH^- ions

Bronsted-Lowry Acid

a proton (H⁺) donor

Bronsted-Lowry Base

a proton (H⁺) acceptor

Lewis Acid

accepts a pair of electrons

Lewis Base

donates a pair of electrons

Henderson-Hasselbach Equation

pH = pKa + Log ([A^-]/[HA])

determine the ratio of acid to conjugate base in a buffer solution

Heat absorbed/released:

q= mass x C x delta T

Common Oxidizing agents: (4)

• O2

• O3

• H2SO4

• halogens

Common Reducing Agents: (4)

• H2

• Fe

• Zn

• Alkali metals (Group 1)