equilibrium – acids and bases. review of acids and bases arrhenius theory of acids and bases ▫an...

Equilibrium – Acids and Bases

Review of Acids and Bases

•Arrhenius Theory of Acids and Bases▫An acid is a substance that dissociates in

water to produce one or more hydrogen ions (H+)

▫A base is a substance that dissociates in water to form one or more hydroxide ions. (OH-)

▫Examples: Acid: HCl(aq) H+

(aq) + Cl-(aq)

Base: LiOH Li+(aq) + OH-

(aq)

Limitations:•Classified based on chemical formula •Some substances do not have OH- in their

chemical formulas but still yield OH- when they react with water. E.g. NH3 (ammonia)

•Solution?

•Bronsted-Lowry Theory of Acids and Bases▫An acid is a proton (H+) donor and must

have H in its formula. ▫A base is a proton acceptor and must

have a lone pair of electrons to form a bond with H+

•Two molecules or ions that are related by the transfer of a proton are called a conjugate acid-base pair. ▫Conjugate acid of a base is the particle

that results when the base receives the proton from the acid.

▫Conjugate base of the acid is the particle that results when the acid donates a proton.

Practice

•Identify the conjugate acid/base pairs in the following:

  NH3(aq) + H2O(l) NH4

+(ag) + OH-

(aq)

•Amphiprotic: Can act as either an acid or a base i.e has both a lone pair and an H-atom▫Ex: H2O

 HCO3

-(aq) ) + H2O(l) H2CO3(aq) + OH-

(aq)

HCO3-(aq) + H2O(l) CO3

2-(aq) + H3O+

(aq)

Strong Acids and Bases

•Completely dissociate in water into their ions (quantitative reactions)

100%

HCl(aq) + H2O(aq) H3O+(aq) + Cl-

(aq)

100%

LiOH + H2O(aq) LiOH(aq) + OH-(aq)

•As a result the [H3O+] in a solution of a strong acid is equal to the concentration of the acid.

•Strong acids include HClO4 (perchloric), HI, HBr, HCl, H2SO4 (sulfuric), and HNO3 (nitric)

•Strong bases include all oxides and hydroxides of alkali metals as well as alkaline earth metal oxides and hydroxides below beryllium.

•The stronger the acid, the weaker it’s conjugate base and vice versa

Weak Acids and Bases• Do NOT completely dissociate in water into their ions 1%

CH3COOH(aq) + H2O(aq) ↔ H3O+(aq) + CH3COO-

(aq)

  1%

NH3(aq) + H2O(aq) ↔ NH4+

(aq) + OH-(aq)

• As a result, the concentration [H3O+] in a solution of a weak acid is always less than the concentration of the dissolved acid.

Percent Ionization

• % Ionization for strong acids is 100%

• % Ionization for weak acids is < 100%

Polyprotic Acids

•Monoprotic acids contain only a single hydrogen ion that can dissociate. ▫Example: HCl

 •Polyprotic acids contain more than one

hydrogen ions that can dissociate. ▫ Example H2SO4, H3PO4

Autoionization of Water

•Water dissociates: H2O(l) + H2O(l) <--> H3O+

(aq) + OH-(aq)

What is the equilibrium constant (K) of this reaction?

Kw = [H3O+][OH-]

Kw is the ion product constant of water

Kw = 1.0 x 10-14 @ SATP

•[H3O+] > [OH-] acidic

•[H3O+] < [OH-] basic

•[H3O+] = [OH-] neutral

Practice•There is a 0.25 mol/L solution of HBr(aq)

a) Calculate the hydrogen ion concentration

b) Calculate the hydroxide ion concentration

• Strong acid – ionizes completely

• Kw = [H3O+][OH-] = Kw = 1.0 x 10-14

Practice•In a 0.13 mol/L solution of NaOH, what is

the [H+] and [OH-]?

•NaOH is hydroxide of an alkali metal so it is a STRONG base meaning [OH-]= [base]

•Kw = [H3O+][OH-] = Kw = 1.0 x 10-14

The pH Scale

•Measures the acidity of a solution. •Measure [H+] in a solution. •Ranges from 0 to 14 •Distilled water is 7 (neutral) •Acids < 7 •Bases > 7•A logarithmic scale

▫A pH of 1 is ten times more acidic then a pH of 2

pH equations

• pH = -log[H3O+]

• [H3O+] = 10-pH

•pOH = -log[OH-]

•[OH-] = 10-pOH

•pH + pOH = 14

Practice•Calculate the pH of a solution of 1.24 x 10-4

M HCl

•pH = -log[H3O+]

•pH = -log[1.24 x 10-4 mol/L]

• pH = 3.91

Practice• If the normal pH of blood is 7.3, then find the pOH,

[H3O+] and [OH-]

• pH + pOH = 14• 7.3 + pOH = 14• pOH = 6.7

• [H3O+] = 10-pH

• [H3O+] = 10-7.3

• [H3O+] = 5 x 10-8

• [OH-] = 10-pOH

• [OH-] = 10-6.7

• [OH-] = 2 x 10-7

Acid- Base Strength & Dissociation•Recall: Strong acids and bases dissociate

quantitatively (>99.9%) in water•Weak acids and bases dissociate partially

in water•When a weak acid or base is added to water

dynamic equilibrium is established

The Acid-Dissociation Constant, Ka

HA(aq ) H2O(l ) H3O(aq ) A

(aq )

Ka [H3O

][A ]

[HA]

For Weak Acids:

All concentrations are those at equilibriumNote: the smaller the value of Ka, the weaker the acid

Determine the Ka of propanoic acid (C2H5COOH(aq)) given that a 0.10 mol/L solution has a pH of 2.96.

(Hint: use an ICE table)

[H3O+] = 10-pH

[H3O+] = 10-2.96

[H3O+] =0.00110

C2H5COOH(aq) CH3COO- + H3O+

I 0.10 mol/L o mol/L 0 mol / L

C -x +x+x

E

Ka [H3O

][A ]

[HA]

The Base-Ionization Constant, Kb

B(aq ) H2O(l ) BH (aq ) OH

(aq )

Kb [BH ][OH ]

[B]

For Weak Bases:

All concentrations are those at equilibriumNote: the smaller the value of Kb, the weaker the base

Calculate the pH of a 3.6 X 10-3 mol/L solution of quinine (C20H24N2O2(aq)). Kb = 3.3 X 10-6

C20H24N2O2(aq) + H2O HC20H24N2O2 + +

OH-

I CE

Kb [BH ][OH ]

[B]

Relationship between Ka, Kb, & KwExample: Consider acid HCN and conjugate base CN-

HCN(aq) + H2O(l) H3O+(aq) + CN-

(aq)

Ka = [H3O+] [CN -]

[HCN]

Kb = [HCN] [OH -] [CN-]

Ka Kb = [H3O+] [CN -] [HCN] [OH -] [HCN] [CN-]

Ka Kb = [H3O+] [OH -]

Ka Kb = Kw

Practice•The Kb for hydrazine, N2H4(g), a rocket

fuel, is 1.7 x 10-6. What is the Ka of its conjugate acid, N2H5 (aq)?

•Ka Kb = Kw

•Ka (1.7 x 10-6)= 1.0 x 10-14

•Ka = 6.0 x 10-9

Practice• Chloracetic acid, HC2H2O2Cl(aq) is a weak acid.

Determine the pH of a 0.0100 mol/L solution of chloracetic acid if the Kb of the conjugate base is Kb= 7.35 x 10-12 .

HC2H2O2Cl (aq) C2H2O2Cl- + H3O+

I CE

• Ka Kb = Kw

• Ka (7.35 x 10-12 )= 1.0 x 10-14

• Ka =0.00136

Ka [H3O

][A ]

[HA]

Neutralization Reactions

•A salt is an ionic compound that results from a neutralization reaction

•Acid + base salt + water•Salts are strong electrolytes that

completely ionize in water•Salts can affect the pH of a solution

Neutral Salt Solutions

•Strong acid + strong base•Both will dissociate completely•Therefore…•Salts containing an anion from a strong

acid and cation from a strong base will be neutral

•Ex: NaOH + HCl NaCl + H2O

Acidic Salt Solutions•Strong acid + weak base •The acid dissociates completely, but the

base only dissociates partially•Therefore…•Salts containing an anion from a strong

acid and a cation from a weak base will be acidic

•Ex: HCl + NH3 NH4Cl NH4+ + Cl-

•NH4+ will act as a weak acid

Basic Salt Solutions•Weak acid + strong base•The base will dissociate completely but

the acid will only dissociate partially•Therefore…•Salts containing an anion from a weak

acid and a cation from a strong base will be basic

•Ex: HC2H3O2 + NaOH NaC2H3O2 + H2O Na+ + C2H3O2-

•C2H3O2- will act as a weak base

Buffers•Resist changes in pH when a moderate

amount of acid or base is added

•The acid and base components must not react in a neutralization reaction

•Solutions of a weak acid and the salt of its conjugate base OR a weak base and the salt of its conjugate acid

Acetic Acid/Sodium Acetate Buffer• Consider a buffered solution made by adding similar

molar concentrations of acetic acid (CH3COOH) and its salt, sodium acetate (CH3COONa)

• Sodium acetate ionizes completely in water:

• When an acid is added to the buffer, the acetate ion reacts with the hydronium ion to neutralize the solution

• When a base is added to the buffer, the acetic acid reacts with the hydroxide ions to neutralize the solution

)()(3)(3 aqaqs NaCOOCHCOONaCH

)(2)(3)(3)(3 laqaqaq OHCOOHCHOHCOOCH

)(2)(3)()(3 laqaqaq OHCOOCHOHCOOHCH

Buffer Examples

•It is extremely important for blood to remain near it’s optimal pH of 7.4

•Any change greater than 0.2 is life-threatening

•If the blood were not buffered, the acid absorbed by consuming a glass of orange juice would probably kill you