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30-3-2011

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pH of salt solutions1. Salts derived from strong acids and strong

basesThese consist of cations from strong bases

and the anions from strong acids Examples: NaCl, KNO3, NaClO4, KBr, etc.These have no effect on the [H+] when

dissolved in water, since cations of strong bases are very weak conjugate acids (do not react with water), and anions of strong acids are very weak bases that do not react with water, thus resulting in neutral solutions.

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2. Salts derived from weak acids and strong bases

These are salt with cations from strong bases (such as Na+ or K+, which are very weak conjugate acids that do not react with water) and anions which are the conjugate bases of a weak acid (strong conjugate bases that react with water to produce hydroxide ions).

The aqueous solution of these salts will be basic.

Examples: NaF, KCN, NaC2H3O2, Na3PO4, Na2CO3, K2S, Na2C2O4, etc.

2. Salts derived from weak acids and strong bases

These are salt with cations from strong bases (such as Na+ or K+, which are very weak conjugate acids that do not react with water) and anions which are the conjugate bases of a weak acid (strong conjugate bases that react with water to produce hydroxide ions).

The aqueous solution of these salts will be basic.

Examples: NaF, KCN, NaC2H3O2, Na3PO4, Na2CO3, K2S, Na2C2O4, etc.

2. Salts derived from weak acids and strong bases

These are salt with cations from strong bases (such as Na+ or K+, which are very weak conjugate acids that do not react with water) and anions which are the conjugate bases of a weak acid (strong conjugate bases that react with water to produce hydroxide ions).

The aqueous solution of these salts will be basic.

Examples: NaF, KCN, NaC2H3O2, Na3PO4, Na2CO3, K2S, Na2C2O4, etc.

2. Salts derived from weak acids and strong bases

These are salt with cations from strong bases (such as Na+ or K+, which are very weak conjugate acids that do not react with water) and anions which are the conjugate bases of a weak acid (strong conjugate bases that react with water to produce hydroxide ions).

The aqueous solution of these salts will be basic.

Examples: NaF, KCN, NaC2H3O2, Na3PO4, Na2CO3, K2S, Na2C2O4, etc.

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Examples

NO2-(aq) + H2O (l) HNO2(aq) + OH-

(aq)

CN- + H2O HCN + OH-

ClO- + H2O HClO + OH-

PO43- + H2O HPO4

2- + OH-

You should observe that OH- is produced which means the reaction should be basic and we need to calculate Kb for the reaction.

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Calculate the pH of a 0.45M NaCN solution. The Ka value for HCNis 6.2 x 10-10.Solution: Since HCN is a weak acid, the Cyanide ion must have

significant affinity for protons.

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

(aq)

Initial 0.45 0 0Change -X +x +xEquilibrium 0.45 – x x x

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

The value of Kb can be calculatedfrom Kw and the Ka value for HCN.

Kb = = = 1.61 x 10-5 Kw

Ka (for HCN)1.0 x 10-14

6.2 x 10-10

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Thus:Kb = 1.61 x 10-5 = = =

[HCN][OH-] [CN-]

(x)(x)0.45 - x

~ x2

0.45

x = 2.69 x 10-3 ~

RE = 2.69*10-3/0.45}*100% = 0.6%The assumption is valid, where 0.45 >> x

x = [OH-] = 2.69 x 10-3 M

pOH = -log[OH-] = 2.57

pH = 14.00 – 2.57 = 11.43

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What is the pH of a 1.0 M Na2CO3 solution? For H2CO3, Ka1 = 4.3*10-7 and Ka2 = 4.8*10-11

Solution: In the solution, [Na+] = 2.0 M and [CO32–] = 1.0

M.

H2O + CO32– HCO3

- + OH– Kb = Kw / Ka2 = 2.1*10-4 H2O H+ + OH- K

w = 10-14

Comparing the two equilibria, it is evident that the first gives about 1010 more OH- than that produced from water. Thus, CO3

2– is a much stronger base than water and the solution is basic.

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H2O + CO32– HCO3

- + OH– Initial 1.0 0 0Change -x +x +xEquilibrium 1.0 – x x x

Kb = Kw / Ka2 = 2.1*10-4

2.1*10-4 = {X2/(1.0 – X)}, assume 1.0 >> xX = 1.45*10-2

RE = {1.45*10-2/1.0} * 100% = 1.45%Assumption is OK, 1.0 >> xX = [OH-] = 1.45*10-2 MpOH = 1.84pH = 14 – 1.84 = 12.16

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3. Salts derived from weak bases and strong acidsSalt containing the conjugate acid of a weak base (strong conjugate acid that dissociates to produce H+), and the anion of a strong acid (very weak conjugate base that does not react with water and will not affect acidity, like Cl-, Br-, I-, NO3

-, ClO4-)

Examples: NH4Cl, ArNH3NO3, etc

NH4+ NH3 + H+

ArNH3+ ArNH2 + H+

The solution is therefore acidic when such salts are dissolved in water.

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Find the pH of a 0.1 M NH4Cl soln. Kb = 1.8*10-5.

NH4+ NH3 + H+

0.1 – x x x

Ka = {10-14/1.8*10-5} = 5.6*10-10

5.6*10-10 = x2/(0.1 – x), assume 0.1>>>x

X = 7.5*10-6,

RE = {7.5*10-6/0.1}*100= 7.5*10-3% (OK)

X = [H+] = 7.5*10-6 M, pH = 5.13

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4. Salts derived from a weak base and a weak acid

Here the conjugate acid of the base will be strong and will react with water to produce H+. At the same time, the conjugate base of the weak acid will also be strong and will react with water to produce OH-.

The question is: which is larger the amount of H+ or OH-?

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Look at the equilibrium constants of the conjugate acid and the conjugate base. The acidity of the salt solution will depend on the larger one.

Example: NH4HSAcid part: NH4

+ NH3 + H+ Ka = {10-14/1.8*10-5} =5.6*10-10

Base part: HS- + H2O H2S + OH- Kb = {10-14/ka1} = {10-14/9.5*10-8} = 1.1*10-7

It is evident that the equilibrium constant (Kb) responsible for OH- is much greater than that related to H+ production (Kb), thus a basic solution is obtained

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5. Salts derived from a cation of small size and large charge and an anion from a strong acid

The cations that have a small size but large charge are conjugate acids (Lewis acids) and will react with water to produce H+, while anions from the strong acids will not react with water as these are very weak conjugate bases.

Therefore, such salts will produce acidic solutions.

Examples: AlCl3, Fe(NO3)3, etc.

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Lewis Acids and Bases

-The most general acid-base definition is based on electrons and is called the Lewis definition

(same Lewis that gave us Lewis structures) .

Acid = Electron pair acceptor

Base = Electron pair donor

• Lewis acid: An electron pair acceptor (e.g. AlCl3)

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Lewis Acids and Bases

Example: AlCl3 + Cl- AlCl4-

AlCl3 accepts a pair of electrons = Acid

Cl- donates a pair of electrons = Base

The above is an acid base reaction in the Lewis sense; AlCl3 & Cl- are Lewis acids & bases but not Brønsted or Arrhenius acids & bases.

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Lewis Acids

• Lewis acids are defined as electron-pair acceptors.

• Atoms with an empty valence orbital can be Lewis acids.

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Lewis Bases

• Lewis bases are defined as electron-pair donors.

• Anything that could be a Brønsted–Lowry base is a Lewis base.

• Lewis bases can interact with things other than protons, however.

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Metal Cations as Lewis Acids

M2+ + 4 H2O(l) M(H2O)42+

Metal ions can accept electron pairs from water molecules to form complexes. An example is nickel which forms the hexa aqua complex:

Ni2+ + 6 H2O(l) Ni(H2O)62+

(aq)

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Ka Values of Some Hydrated Metal Ions at 25oC

Ion Ka

Fe3+ (aq) 6 x 10-3

Sn2+ (aq) 4 x 10-4

Cr3+ (aq) 1 x 10-4

Al3+ (aq) 1 x 10-5

Be2+ (aq) 4 x 10-6

Cu2+ (aq) 3 x 10-8

Pb2+ (aq) 3 x 10-8

Zn2+ (aq) 1 x 10-9

Co2+ (aq) 2 x 10-10

Ni2+ (aq) 1 x 10-10

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Identifying Lewis Acids and Bases

Problem: Identify the Lewis acids and bases in the following reactions:

(a) F- + BF3 BF4-

(b) Co2+ + 6 H2O Co(H2O)62+

(c) NH3 + H+ NH4+

(a) The BF3 accepted an electron pair from the fluoride ion. BF3 is the acid and F- is the base.(b) The Co2+ ion accepted the electron pairs from the water molecules. Co2+ is the acid and H2O is the base.(c) The H+ ion accepted the electron pair from the ammonia molecule. H+ is the acid and water is the base.

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Solution:

Since the Al(H2O)63+ ion is a stronger acid than water,

the dominate equilibrium will be:

1.4 x 10-5 = Ka = [Al(OH)(H2O)5

2+][H+] [Al(H2O)6

3+]

Calculate the pH of a 0.010 M AlCl3 solution. The Ka value for the Al(H2O)6

3+ ion is 1.4 x 10-5.

Al(H2O)63+

(aq) Al(OH)(H2O)52+

(aq) + H+(aq)

0.01 –x X X

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Thus:

Ka =[Al(OH)(H2O)5

2+][H+] [Al(H2O)6

3+]

1.4*10-5 = = (x) (x)0.010 - x

x2

0.010

x = 3.7 x 10-4

The approximation is valid

[H+] = x = 3.7 x 10-4 M and pH = 3.43

RE = {3.7*10-4 /0.01} * 100% = 3.7%

Assume 0.01 >> x

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6. Salts derived from a cation of small size and large charge and an anion from a weak acid

The cations that have a small size but large charge are conjugate acids (Lewis acids) and will react with water to produce H+, while anions from the weak acids will react with water to produce OH-.

Therefore, such acidity will depend on the values of Ka (for the hydrated cation) and the Kb for the anion.

Ka > Kb results in an acidic solutionKb > Ka results in a basic solutionKa = Kb results in a neutral solutionExamples: Al(CH3COO-)3, Fe(NO2)3, etc.

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Determine whether an aqueous solution of iron(III) nitrite, Fe(NO2)3, is acidic, basic, or neutral.

Plan: The formula consists of the small, highly charged, and therefore weakly acidic, Fe3+ cation and the weakly basic NO2

- anion of the weak acid HNO2, To determine the relative acidity of the solution, we write equations that show the reactions of the ions with water, and then find Ka and Kb of the ions to see which ion reacts to form to a greater extent.

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Fe(H2O)63+

(aq) + H2O(l) Fe(H2O)5OH2+(aq) + H3O+

(aq)

NO2-(aq) + H2O(l) HNO2(aq) + OH -

(aq)

Obtaining Ka and Kb of the ions: For Fe3+(aq)

Ka = 6 x 10-3. For NO2

-(aq), Kb must be determined:

Kb of NO2- = = = 1.4 x 10-11

Ka of HNO2

Kw 1.0 x 10-14

7.1 x 10-4

Since Ka of Fe3+ > Kb of NO2-, the solution is acidic.

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Are the solutions of the following salts neutral, basic or acidic

NaCl? NaAc? NH4Cl? NH4Ac? KClO4? CH3NH3ClO4?

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Predict whether an aqueous solution of each of the following salts will be acidic, basic, or neutral.a) NH4C2H3O2 b) NH4CN c) Al2(SO4)3

a) The ions are the ammonium and acetate ions, Ka for NH4

+ is 5.6 x 10-10, and Kb for C2H3O2

- is 5.6 x 10-10 . Since they are equal, the solution will be neutral and the pH close to 7.

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b) The solution will contain the ammonium and cyanide ions, the Ka value for NH4

+ is 5.6 x 10-

10, and

Since Kb for CN- is much larger than Ka for NH4

+, this solution will be basic.c) This solution contains the hydrated Aluminum

ion and the sulfate ion. Ka for Al(H2O)63+ = 1.4

x 10-5, for sulfate, Kb = 8.3 x 10-13 ; therefore this solution will be acidic.

Kb (for CN-) = = 1.6 x 10-5Kw

Ka(for HCN)

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Type of Salt Examples Comment pH of Solution

Cation is from Neither acts asstrong base; KCl, KNO3 an acid or a Neutralanion is from NaCl, NaNO3 basestrong acid

Cation is from Anion acts asstrong base; NaC2H3O2 a base; cation Basicanion is from KCN, NaF has no effectweak acid on pH

Cation is conjugate Cation acts asacid of weak base; NH4Cl, an acid; anion Acidicanion is from NH4NO3 has no effectstrong acid on pH

Summary of Acidity of Salt solutions

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Cation is conjugate Cation acts as Acidic if :acid of weak base NH4C2H3O2 an acid; anion Ka > Kb

anion is conjugate NH4CN acts as a base Basic if :base of weak acid Kb > Ka

Neutral if : Ka = Kb

Cation is highly Hydrated cationcharged metal ion; Al(NO3)3, acts as an acid; Acidicanion is from FeCl3 anion has nostrong acid effect on pH

Cation is highly charged metal ion; anion is conjugate base of weak acid

Al2(SO4)3

Fe(NO2)3

Hydrated cation acts as an acid; anion acts as a

base

Acidic: Ka>Kb

Basic: Kb>Ka

Neutral: Ka = Kb

Cation is conjugate Cation acts as Acidic if :acid of weak base NH4C2H3O2 an acid; anion Ka > Kb

anion is conjugate NH4CN acts as a base Basic if :base of weak acid Kb > Ka

Neutral if : Ka = Kb

Cation is highly Hydrated cationcharged metal ion; Al(NO3)3, acts as an acid; Acidicanion is from FeCl3 anion has nostrong acid effect on pH

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Selected Problems

5-7, 12, 13, 15-19, 21-24, 26, 29, 31-35, 37, 41-47, 51, 52, 55-58, 61-63, 67-70, 73, 75.