Chapter 16 : Acid-Base Equilibira

Lauren Querido

Table of Contents

16.1 Review16.2 Brønsted-Lowry

Acids and Bases16.3 Autoionization of

Water16.4 pH Scale16.5 Strong Acids and

Bases16.6 Weak Acids16.7 Weak Bases

16.8 Relationship Between Ka and Kb

16.9 Acid-Base Properties of Salt Solutions

16.10 Acid-Base Behavior and Chemical Structure

16.11 Lewis Acids and

Bases

16.1 Review

• Acids– Sour in taste

– Litmus paper turns red

• Bases– Bitter, slippery

– Litmus paper turns blue

• When acids and bases mix, their properties disappear!

Arrhenius Acids and Bases

• Svante Arrhenius (1880)– In aqueous solutions:

• Acids will increase the concentration of H+ ions when dissolved in water.

• Bases will increase the concentration of OH-

ions when dissolved in water.

16.2 Brønsted-Lowry Acids and Bases

• 1923 Brønsted and Lowry made a more general definition

– Brønsted-Lowry Acid is a substance that can transfer a proton. It must have a hydrogen atom that can be lost as H+.

– Brønsted-Lowry Base is asubstance that can accept a proton. Must have a nonbonding pair of electrons to gain a H+ ion.

Conjugate Acid-Base Pairs

• Conjugate base- Removal of proton from the acid• Conjugate acid- Addition of proton to the base

Relative Strengths of Acids and Bases

• The stronger the acid, the weaker its conjugate base.• The stronger the base, the weaker its conjugate acid.

1. Strong acids completely transfer protons to water.

2. Weak acids partly dissociate in aqueous solutions and exist as a mixture of acid molecules and component ions.

3. Negligible acidity contain Hydrogen but do not demonstrate acidic behavior. Ex: CH4

• Position of equilibrium favors transfer of proton from stronger acid to stronger base.

16.3 Autoionization of Water

• Ion product of water–

– 1.0 x 10 –14 = [H+] [OH-]– This is used to calculate concentrations of H+

and OH- .• If [H+] = [OH-], than neutral equation• If [H+] > [OH-], than acidic equation• If [H+] < [OH-], than basic equation

=

16.4 The pH Scale

• pH = -log [H+]

• pH of 7 is neutral

• Acidic solution 0 < pH < 7

• Basic solution 14 > pH > 7

• Other p scales are – pOH = -log [OH-]

– pOH + pH = -log Kw = 14.0

Examples on the pH Scale

Measuring pH

• A pH meter consists of a pair of electrodes connected to a meter which pH is generated when placed in the solution.

• An acid-base indicator turns a color if placed in acid or base. Ex: litmus paper

16.5 Strong Acids and Bases

• Strong Acids– 7 most common strong acids are

• HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4

– In acidic reactions, equilibrium lies entirely to the right side.

– Completely dissociates – Example:

• HNO3 => H+ + NO3-

Strong Bases• Most common strong bases are ionic hydroxides

of alkali metals (1A) and heavier alkaline earth metals (2A). Examples: LiOH, RbOH, CsOH, NaOH, KOH, and Ca(OH)2, Sr(OH)2, and Ba(OH)2.

• Other strong bases react with water to form OH- such as Na2O, CaO.

• Also, anions O2-, H-, and N3- are stronger bases than OH- and therefore remove a proton from H2O.– Example: N3- + H2O => NH3 + 3OH-

16.6 Weak Acids

• A weak acid only partially ionizes in aqueous solutions.

• General weak acid equation– HX H+ + X- where H is Hydrogen

– Many weak acids contain some Hydrogen atoms bonded to carbon atoms and oxygen atoms (organic compounds).

– Ka is the acid dissociation constant.

– The larger the value of Ka , the stronger the acid.

Calculating Ka from pH

• Use and ICE box!• Sample exercise

– A student prepared a .10 M solution of formic acid and measures its pH which was 2.38.

• A) calculate Ka for formic acid

• B) what percentage of the acid is ionized in the .10M solution?

Answer

a) HCHO2 H+ + CHO2-

Ka = [H+][CHO2-]

[HCHO2]

pH= -log[H+]

=10 –2.38 = 4.2 X 10-3M

Ka = [4.2 X 10–3 ][4.2 X 10–3] [.10]

1.8 X 10-4 = [4.2 X 10–3 ][4.2 X 10–3] [.10]

b) Percent Ionization =Concentration of H+Initial concentration

of component = 4.2% HCHO2 H+ CHO-

I .10 M 0 M 0 M

C -4.2 X 10–3 +4.2 X 10–3 +4.2 X 10–3

E .10 - 4.2 X 10–3 +4.2 X 10–3 +4.2 X 10–3

Using Ka to Calculate pH

The best way to explain this is by an example.

Calculate the pH of a .30 M solution of acetic acid at 25o C. (Ka = 1.8 X 10-5)

So… HC2H3O2 H+ + C2H3O2-

Ka = [H+][C2H3O2-] = 1.8 X 10-5

[HC2H3O2]

What now?

HC2H3O2

H+ C2H3O2-

I .30 M 0 M 0 M

C -x +x +x

E .30-x x x

Ka = (x)(x) = 1.8 X 10-5 (.30 – x)

Either do the quadratic equation or in this case you can take out x in the denominator.

[H+] = x = 2.3 X 10-3

pH = -log 2.3 X 10-3 = 2.64

Polyprotic Acids

• Polyprotic acids have more than one ionizable Hydrogen atom.

• Example: • H2SO3 H+ + HSO4

-

• HSO4- H+ + SO32-

• The second Ka (Ka2) is much smaller than Ka1 because it is easier to remove the first proton.

16.7 Weak Bases• Weak base + water => conjugate acid + hydroxide

ion• Kb is the base-dissociation constant (equilibrium in

which base reacts when H2O to form conjugate acid and OH- ion).

• Types of weak bases:1. Neutral substances that have atoms with a non-bonding

pair of electrons that can serve as a proton acceptor.– Most of these contain amines, N-H which is sometimes replaced

with a bond between C or N Ex: NH2CH3

2. Anions of weak acids– Ex: ClO- + H2O HClO + H+

– ClO- is the weak base

16.8 Relationship Between Ka and Kb

• Reaction 1 + reaction 2 = reaction 3

Which leads to K1 x K2 = K3

Which leads to Ka x Kb = Kw • Kw is the ion-product constant for water

– Kw = 1 x 10-14

• As the strength of the acid increases, the strength of the base decreases and visa-versa.

• pKa + pKb = pKw = 14.00

16.9 Acid-Base Properties of Salt Solutions

• Hydrolysis is the process at which ions react with water and produce H+ or OH-

X- + H2O HX + OH-

• Anions of strong acids do not influence pH– Ex: NO3

-

• Anions that still have ionizable protons are amphoteric – Ex: HSO3

- from H2SO4

• Most cations (except 1A elements and Ca+2, Sr+2. Ba+2) act as weak acids in solution.

Predicting the pH of a Solution1. Salts derived from a strong acid and a strong base makes a

neutral pH (pH of 7).

• NaOH + HCl => NaCl + H2O

2. Salts derived from a strong base and a weak acid will yield a pH of above 7 because the anion hydrolyzes to produce OH- ions and the cation does not hydrolyze.

• NaOH + HClO => NaClO + H2

3. Salts derived from a weak base and a strong acid will result in a pH that is below 7 because the cation hydrolyzes to produce H+ ions and the anion does not hydrolyze.

• Al(OH)3 + 3HNO3 => Al(NO3)3 + 3H2O

4. Salts derived from a weak base and a weak acid will yield a pH that is dependant on the constant value of the constant dissociations (Ka and Kb).• if the base is more basic than the acid is acidic,

then the solution will have a pH that is greater than 7.

• If the acid is more acidic, than the pH will be less than 7.• NH4

+ + CN- NH4CN

• NH4+ Ka = 5.6 X 10-10

• CN- Kb= 2.0 X 10-5

• Therefore, the pH of NH4CN is greater than 7

16.10 Acid-Base Behavior and Chemical Structure

• Factors that effect acid strength– If H-X bond is polarized (X is more electronegative) the H

acts as a proton acceptor.– Non-polar bonds (CH4) produce neutral solutions. – Weaker bonds dissociate more easily than very strong bonds. – HF is a weak acid because of this.– The greater the stability of the conjugate base, the weaker the

acid.– Ultimately, there are three factors effecting acid strength:

• Polarity of H-X bond• Strength of H-X bond• Stability of conjugate base, X-

Binary Acids

• Binary acids are composed of Hydrogen and a non-metal. – Ex: HCl, HF, H2S, etc.

• The more polar the bond,the stronger it is• The weaker the bond, the stronger the acid. • Strength of the bond decreases (acidity increases)

as the element increases in size or moves down a group.

• Acid strength increases (acidity decreases) moving from left to right

Group

4A 5A 6A 7APeriod 2 CH4

No acid or base properties

NH3

Weak base

H2O

-------

HFWeak acid

Period 3 SiH4No acid or base

properties

PH3

Weak base

H2SWeak acid

HClStrong acid

Increasing acid strength

Increasing acid strength

Increasing base strength

Increasing base strength

Oxyacids

• Oxyacids are acids with an OH group is bound to a central atom.– Example: H2SO4

–

OH- Bonding• To determine if an OH group acts as an acid or base, consider

this:

• If Y is a metal than sources of OH- behave as bases.• If Y is a non-metal than the compound will not readily lose the

OH- ion.– The electronegativity will increase and so will the acidity.

• The increasing number of Oxygen atoms stabilizes the conjugate base and thus increases the strength of the acid.

Oxyacid Rules of Thumb

1. Oxyacids that have the same number of OH groups and the same number of Oxygen atoms, acid strength increases with increasing elecronegativity of the central atom

• Example: HClO > HBrO > HIO (> = more acidic)

2. For oxyacids with the same central atom, acid strength increases with increasing number of Oxygen atoms that are attached.

• Example: HClO < HClO2 < HClO3 < HClO4 ( < less acidic)

Carboxylic Acid• Carboxylic acids are organic compounds.

• -COOH is the functional group

• -R is either a Hydrogen

or Carbon based group

– If an extra Oxygen is added than it stabilizes the conjugate base and increases the acidity.

– If conjugate base has resonance structures, it spreads the negative charge evenly over the compound.

– Acid strength of carboxylic acid increases as the number of electronegative atoms increase.

6.11 Lewis Acids and Bases• G.N. Lewis proposed this:• Lewis Acids have an incomplete

octet of electrons. Function as

electron pair acceptors• Lewis Bases act as electron pair

donators

Hydrolysis of Metal Ions

• Hydration is a process when when metals attract unshared electron pairs of water molecules.– The metal acts as Lewis acid– The water acts as Lewis base

– Ex: Fe(H2O)6+3 Fe(H2O)5(OH)2+ + H+

– So, general equation• M(H2O)n

c M(H2O)n-1(OH)c-1 + H+