The Properties of Acids and Bases | General Chemistry 3

The properties of acids and bases are studied in this chapter: ion-product constant, definition of pH and pOH, pH calculations, strong and weak acids and bases, dissociation constant, conjugate acid-base pairs, polyprotic acids

Acids and Bases

Arrhenius: acid donates H+; base donates HO-
Bronsted: acid donates H+; base accepts H+
Lewis: acid accepts electron pair; base donates electron pair

Arrhenius acid: HCl; Arrhenius base: NaOH
Bronsted acid: HCl; Bronsted base: NH3
Lewis acid: BF3; Lewis base: NH3

Proton-transfer reaction (or protonation reaction): reaction involving the transfer of a proton from one molecule to another. Acid-base reactions are proton transfer reactions

HCl (aq) + H2O (l) → Cl(aq) + H3O(aq) is a proton-transfer reaction

Ion-Product Constant

Self-ionization of water:   2 H2O (l)  H3O(aq) + HO(aq)

Ion-product constant of water Kw:

Equilibrium-constant expression for the self-ionization of water:
Kw = [H3O+][HO-]
Kw = 1.0 x 10-14 2  at 25°C

Neutral solution: [H3O+] = [HO-] = 1.0 x 10-7 M
Acidic solution: [H3O+] > [HO-]
Basic solution: [H3O+] < [HO-]

pH and pOH

pH scale: numeric scale used to quantify the acidity or basicity of an aqueous solution
pH and pOH are unitless


pH = - log [H3O+]

[H3O+] = concentration of H3O+ (in M)
⇒ [H3O+] = 10-pH​​​​​​


pOH = - log [HO-]

[HO-] = concentration of HO- (in M)
⇒ [HO-] = 10-pOH


Neutral solution: [H3O+] = 1.0 x 10-7
⇒ pH = -log [H3O+] = 7

Acidic solution: [H3O+] > 1.0 x 10-7
- log [H3O+] < - log ( 1.0 x 10-7 )   ⇒   pH < = 7


Neutral solution: pH = pOH = 7
Acidic solution: pH < 7
Basic solution: pH > 7

Relationship between pH and pOH:

pH + pOH = 14

Kw = [H3O+] [HO-] = 1.0 x 10-14 2  
⇒ - log ( [H3O+] [HO-] ) = -log ( 1.0 x 10-14 )
⇒ - log [H3O+] - log [HO-] = 14
⇒ pH + pOH = 14


Strong and Weak Acids / Bases

Strong acids / strong bases: acids / bases completely dissociated in aqueous solutions
⇒ strong acids are able to donate protons to water (H3O+ formation)
⇒ strong bases are able to accept protons from water (HO- formation)

HBr is a strong acid:
HBr (aq) + H2O (l) → Br- (aq) + H3O+ (aq)


There are only a few strong acids / bases in water
Most acids / bases are weak in water
⇒ weak acids are partially dissociated in water into H3O+ and its conjugate base
⇒ weak bases are partially dissociated in water into HO- and its conjugate acid

NH3 is a weak base:
NH(aq) + H2O (l)  NH4(aq) + HO(aq)


% dissociation of a weak acid = H3O+acid0 × 100


Dissociation Constant and Strength

Ka and Acid Strength:

Acid dissociation reaction of an acid HA: HA (aq) + H2O (l)  H3O(aq) + A(aq)
Acid dissociation constant Ka: equilibrium constant of the acid dissociation reaction (considering [H2O] = 1M)
Acids with large values of Ka are stronger than acids with smaller values of Ka

Ka = A-H3O+HA 

pKa = - log Ka


Kb and Base Strength:

Base dissociation reaction of a base A-: A(aq) + H2O (l)  HA (aq) + HO(aq)
Base dissociation constant Kb: equilibrium constant of the base dissociation reaction (considering [H2O] = 1M)


pKb = - log Kb


Conjugate Acid-Base Pairs

Conjugate acid-base pair HA/A-: a pair of acid and base which differ only by the presence or absence of a proton

CH3COOH / CH3COO- is a conjugate acid-base pair
CH3COOH is the conjugate acid of the base CH3COO-


Relation between Ka and Kb for a conjugate acid-base pair:

Ka Kb = Kw

Ka = Acid dissociation constant
Kb = Base dissociation constant
Kw = equilibrium-constant expression for the self-ionization of water

Relation between pKa and pKb for a conjugate acid-base pair (at 25°C):

pKa + pKb = pKw = 14

⇒ the stronger the acid, the weaker its conjugate base
⇒ the weaker the acid, the stronger its conjugate base


pH Calculations

1) Determine what ions will be formed when the salt is dissolved in water
2) Determine whether the salt contains a conjugate weak acid or weak base
3) Determine the equilibrium constant for the reaction
4) Set up an equilibrium table
5) Calculate [H3O+] or [HO-]

The method of successive approximations is often used to solve acid-base equilibrium problems or to determine the pH of a solution

Polyprotic Acids

Amphoteric compound: compound that can act as both a Bronsted acid and base
Polyprotic acid: acid that donate more than one acidic proton

H2CO3 is a polyprotic acid:

H2CO(aq) + H2O (l)  HCO3(aq) + H3O(aq)     [Ka1]
HCO3(aq) + H2O (l)  CO32- (aq) + H3O(aq)      [Ka2]