Chemistry: Unit 14-15

Chemical Equilibrium

A dynamic process in which concentrations of reactants and products remain constant over time

Dynamic Equilibrium

The state of a chemical reaction in which reactant and product concentrations do not change, but products and reactants are continually interconverting

Rate(forward) =

Rate(reverse)

Concentrations remain

constant

The Haber Process

1. Initially only N2 and H2 collide and react
2. As NH3 forms there is less N2 and H2 which leads to less collisions = less reactions
3. Equilibrium concentrations are constants;
Rate of dissociation and formation are the same

2NO2(g) ⇌ N2O4(g)

kf/kr = [N2O4]/[NO2]^2 = Keq

Keq

Equilibrium constant

Equilibrium Constant Expression

Ratio of concentrations (or partial pressures) of products to reactants, where each term is raised to a power equal to the coefficient in the balanced chemical equation.

Equilibrium Constant

Value of the ratio of concentration (or partial pressure) terms in the equilibrium constant expression at a specific temperature.

aA + bB ⇌ cC + dD

Kc = [C]^c[D]^d / [𝐴]𝑎[𝐵]𝑏

K >> 1

Equilibrium lies to right; products dominate

K << 1

Equilibrium lies to left; reactants dominate

K = 0

No reaction

K = infinity

Irreversible reaction

Law of Mass Action

The equilibrium constant expression at chemical
equilibrium has a characteristic value at a given
temperature

Mass Action Expression

Equivalent to equilibrium constant expression,
but applied to reaction mixtures that may or may
not be at equilibrium

2 H2(g) + O2(g) ⇌ 2 H2O(g) K = 3 × 1081

Very large K: favors formation of products

Very large K: favors formation of products

2 CO2(g) ⇌ 2 CO(g) + O2(g) K = 3 × 10-92

Very small K: favors reactants; not much product formed at equilibrium

Very small K

Favors reactants; not much product formed at equilibrium

H2O(g) + CO(g) ⇌ H2(g) + CO2(g) K = 24

Intermediate value of K (comparable amounts of products and reactants at equilibrium)

Intermediate value of K

Comparable amounts of products and reactants at equilibrium

Px =

Units of partial pressure

[X] =

Concentration units of moles/liter

P = (n/V)RT

n/V = molarity (moles per liters)

Kp = Kc(RT)Δn

Δn = # of moles of products - # of moles of reactants

Kc = constant

R (Gas Constant)

0.0821 (L * atm) / (mol * K)

K(forward) =

1/K(reverse)

Reaction 2 = ½(Reaction 1)

K2 = (K1)½

Reaction Quotient Q

Where concentrations are not necessarily equilibrium concentrations

Q < K

Reaction goes left to right (forward)

Q > K

Reaction goes right to left (reverse)

Q = K

Reaction is at equilibrium

Homogeneous equilibria

Equilibria involving reactants and products in the
same phase.

Heterogeneous equilibria

Equilibria involving reactants and products in more
than one phase.

I.C.E

Initial, Change, Equation

Common Structures of Acids

Binary Acid (HCl), Oxyacid (H2SO4), Carboxylic (H2C4H4O5)

Common Structures of Bases

Hydroxides (NaOH), Amines (NH3)

Hydronium Ion

H+ or H3O+

Hydroxide Ion

OH-

H+ in water

H3O+

Arrhenius acid

Will increase proton concentration in water

HCl(g) → H+(aq) + Cl-(aq)

HCl is an Arrhenius acid

Arrhenius base

Will increase hydroxide concentration in water

NaOH(s) → Na+(aq) + OH-(aq)

NaOH is an Arrhenius base

Lewis acid

electron pair acceptor

Lewis base

electron pair donor

Does H+ and OH- have to be involved?

H+ and OH− do not necessarily have to be involved

Brønsted-Lowry acid

H+ ion donor

Brønsted-Lowry base

H+ ion acceptor

HCl(aq) + NH3(aq) → NH4+(aq) + Cl-(aq)

HCl is acid, NH3 is base

NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH−(aq)

H2O is acid, NH3 is base

HCl(aq) + H2O(l) → Cl-(aq) + H3O+(aq)

HCl is acid, H2O is base

Conjugate acid-base pairs

Differ from each other only by the presence or absence of a proton.

Acid Ionization Constant (Ka)

Acid strength is measured by the size of the equilibrium constant when it reacts with H2O

Ka >>1

Strong acids

Kb >>1

Strong bases

Ka < 1

Weak Acids

Kb < 1

Weak bases

Strong acid

Completely ionized in water

HBr (Hydrobromic Acid)

HBr(aq) + H2O(l) → Br-(aq) + H3O+(aq)

HCl (Hydrochloric Acid)

HCl(aq) + H2O(l) → Cl-(aq) + H3O+(aq)

HI (Hydroiodic Acid)

HI(aq) + H2O(l) → I-(aq) + H3O+(aq)

HNO3 (Nitric Acid)

HNO3(aq) + H2O(l) → NO3-(aq) + H3O+(aq)

HClO4 (Perchloric Acid)

HClO4(aq) + H2O(l) → ClO4-(aq) + H3O+(aq)

H2SO4 (Sulfuric Acid)

H2SO4(aq) + H2O(l) → HSO4-(aq) + H3O+(aq)

Weak acid

Partially ionized in water.

Acetic Acid

CH3COOH(aq) + H2O(l) → CH3COO-(aq) + H3O+(aq)

Formic Acid

HCOOH(aq) + H2O(l) → HCOO-(aq) + H3O+(aq)

Hydrofluoric Acid

HF(aq) + H2O(l) → F-(aq) + H3O+(aq)

Hypochlorous Acid

HClO(aq) + H2O(l) → ClO-(aq) + H3O+(aq)

Nitrous Acid

HNO2(aq) + H2O(l) → NO2-(aq) + H3O+(aq)

Strong Bases

Oxides/hydroxides of Group 1, 2 metals

Lithium hydroxide

LiOH(aq) → Li+(aq) + OH-(aq)

Sodium hydroxide

NaOH(aq) → Na+(aq) + OH-(aq)

Potassium hydroxide

KOH(aq) → K+(aq) + OH-(aq)

Calcium hydroxide

Ca(OH)2(aq) → Ca2+(aq) + 2OH-(aq)

Barium hydroxide

Ba(OH)2(aq) → Ba2+(aq) + 2OH-(aq)

Strontium hydroxide

Sr(OH)2(aq) → Sr2+(aq) + 2OH-(aq)

Ammonia

NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq)

Aniline

C6H5NH2(aq) + H2O(l) → C6H5NH3+(aq) + OH-(aq)

Dimethylamine

(CH3)2NH(aq) + H2O(l) → (CH3)2NH2+(aq) + OH-(aq)

Methylamine

CH3NH2(aq) + H2O(l) → CH3NH3+(aq) + OH-(aq)

Pyridine

C5H5N(aq) + H2O(l) → C5H5NH+(aq) + OH-(aq)

HClO4

ClO4-

HI

I-

HBr

Br-

HCl

Cl-

H2SO4

HSO4-

HNO3

NO3-

H3O+

H2O

HClO3

ClO3-

HClO2

ClO2-

HSO4-

SO42-

H2SO3

HSO3-

H3PO4

H2PO4-

HF

F-

HNO2

NO2-

HCOOH

HCOO-

CH3COOH

CH3COO-