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Acid and Base Solutions Acids and bases are found in many

common substances and are important in life processes.

Group Work: Make a list of some common acids and bases. How do we know which is which?

There are several models for what constitutes an acid or a base -- three models to be discussed.

14.1 Acids and Bases: A Brief

Review

Acid: Base: tastes sour tastes bitter

stings skin feels slippery

corrosive to metals

releases CO2 from carbonates

turns litmus red turns litmus blue

turns phenolphthalein colorless turns ph. pink

React together to form a salt with loss of the characteristic acid/base properties

Arrhenius Definition Acids produce hydrogen ions (H+ ) in

aqueous solution.

Bases produce hydroxide ions (OH-) when dissolved in water.

Limited Definition.

– Only one kind of base.

– NH3 ammonia could not be an Arrhenius base.

Arrhenius Definition Not realistic: H+ has a radius of 10-13

cm, which gives a very concentrated charge, so it associates with H2O as H(H2O)4

+, which we usually simplify to H3O

+ or H+ (aq)

OH– is also associated with H2O as OH(H2O)3

– which we usually write as OH–(aq).

Bronsted-Lowry Definitions An acid is an proton (H+) donor (victim)

and a base is a proton acceptor (thief).

Acids and bases always come in pairs: Can’t have a thief without a victim!

Example: HCl

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

Acid Base conjugate conjugate

acid base

This is a competitive equilibrium. Can you describe it with your EQ model?

Acid/Base Reactions Activity Identify the Conjugate acid and base

pairs in each reaction ( you may use A, B, CA, CB abbreviations)

Discuss what the K value should be for the reaction as written given observations and extent data. Then make a statement of relative base strength.

Whiteboards when done:

– Rank order the bases in order of strength

– Rank order the acids in order of strength

Acid/Base Pairs General equation

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

Acid + Base Conjugate acid + Conjugate base

This is an equilibrium.

Competition for H+ between H2O and A-

The stronger base controls direction.

If H2O is a stronger base it takes the H+

Equilibrium moves to right.

Acid/Base Pairs – Group Work Write a balanced equation showing how the

following substances behave as acids in water and identify the conjugate acid-base pairs. HNO3 HCO3

- H3PO4 H2PO4-

HNO3(aq) + H2O(l) ⇌ H3O+ (aq) + NO3

-(aq)

HCO3-(aq) + H2O(l) ⇌ H3O

+ (aq) + CO32-(aq)

H3PO4(aq) + H2O(l) ⇌ H3O+ (aq) + H2PO4

-(aq)

H2PO4-(aq) + H2O(l) ⇌ H3O

+ (aq) + HPO42-(aq)

acid 1 base 2 acid 2 base 1

2

Acid Base Strongest HClO4 ClO4

- Weakest acids H2SO4 HSO4

- bases HI I- HBr Br- HCl Cl- HNO3 NO3

- H3O

+ H2O HSO4

- SO42-

H2SO3 HSO3-

H3PO4 H2PO4-

HNO2 NO2-

HF F- CH3CO2H CH3CO2

- H2CO3 HCO3

- H2S HS- NH4

+ NH3 HCN CN- HCO3

- CO32-

HS- S2- H2O OH- Weakest NH3 NH2

- Strongest acids OH- O2- bases

back

Leveling Effect

All acids above H3O+ in the table are

strong acids, which dissociate completely in aqueous solution.

All bases below OH- in the table are strong bases, which dissociate completely in aqueous solution.

The table can be used to make predictions, based on the principle that the stronger acid reacts with the stronger base to form a weaker acid and a weaker base.

Predicting Acid-Base Reactions

HCl + HSO3- ⇌ H2SO3 + Cl-

stronger stronger weaker weaker acid base acid base

We must also consider H2O as a possible acid or base. Thus, HNO3 will transfer its proton to H2O, not to Cl- because H2O is a stronger base than Cl-.

Group Work Write an equation showing the position

of equilibrium for the following mixtures. Remember that H2O can also be either an acid or a base.

HSO4- and F-

HS- and HCO3-

HSO4- + F- ⇌ SO4

2- + HF

HS- + HCO3- ⇌ H2S + CO3

2-

14.2 Acid dissociation constant Ka The equilibrium constant for the

general equation.

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

Ka = [H3O+][A-]

[HA]

H3O+ is often written H+ ignoring the water in equation (it is implied).

Acid dissociation constant Ka HA(aq) H+(aq) + A-(aq)

Ka = [H+][A-]

[HA]

We can write the expression for any acid.

Strong acids dissociate completely.

Equilibrium far to right.

Conjugate base must be weak.

Back to Pairs

Strong acids

Ka is large

[H+] is equal to [HA]

A- is a weaker base than water

Weak acids

Ka is small

[H+] <<< [HA]

A- is a stronger base than water

Types of Acids

Polyprotic Acids- more than 1 acidic hydrogen (diprotic, triprotic).

Oxyacids - Proton is attached to the oxygen of an ion.

Organic acids contain the Carboxyl group -COOH with the H attached to O

Generally very weak.

Amphoteric Behave as both an acid and a base.

Water autoionizes

2H2O(l) H3O+(aq) + OH-(aq)

KW= [H3O+][OH-]=[H+][OH-]

At 25ºC KW = 1.0 x10-14

In EVERY aqueous solution.

Neutral solution [H+] = [OH-]= 1.0 x10-7

Acidic solution [H+] > [OH-]

Basic solution [H+] < [OH-]

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14.3 pH Scale pH= -log[H+]

Used because [H+] is usually very small

As pH decreases, [H+] increases exponentially

Sig figs only the digits after the decimal place of a pH are significant

[H+] = 1.0 x 10-8 pH= 8.00 2 sig figs

pOH= -log[OH-]

pKa = -log K

Measuring pH

litmus or pH paper

color changes of indicators

voltage generated by

electrodes (pH meter)

pH Indicators

Relationships KW = [H+][OH-]

-log KW = -log([H+][OH-])

-log KW = -log[H+]+ -log[OH-]

pKW = pH + pOH

So…

KW = 1.0 x10-14 14.00 = pH + pOH

[H+], [OH-], pH and pOH – Given any one of these we can find the other

three through equilibrium relationships (Kw)

Basic Acidic Neutral

100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14

[H+]

0 1 3 5 7 9 11 13 14

pH

Basic 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14

[OH-]

0 1 3 5 7 9 11 13 14

pOH

14.5 pH and Ka of Acid Solutions Always write down the major ions in

solution.

Remember these are equilibria.

Remember the chemistry.

Don’t try to memorize; there is no one way to do this. Apply good chemistry!

Strong Acid: HNO3

HNO3 is completely dissociated into the ions, H3O

+ and NO3- (work sample 14.7)

Strong Acids

HBr, HI, HCl, HNO3, H2SO4, HClO4 ALWAYS WRITE THE MAJOR

SPECIES

Completely dissociated

[H+]eq = [HA]i

[OH-] is going to be small because of equilibrium

Kw = 10-14 = [H+][OH-]

Changes in pH with Dilution

Group Work: pH for factors of 10 dilution?

What is the pH of 1.0 M HCl?

pH = 0.00

What is the pH of 0.10 M HCl (a 1:10 dilution)?

pH = 1.00

What is the pH of 0.010 M HCl?

pH = 2.00

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Changes in pH with Dilution What is the pH of 1.0 x 10-3 M HCl?

pH = 3.00

What is the pH of 1.0 x 10- 4 M HCl?

pH = 4.00

What is the pH of 1.0 x 10-5 M HCl?

pH = 5.00

What is the pH of 1.0 x 10-6 M HCl?

pH = 5.996

Changes in pH with Dilution What is the pH of 1.0 x 10-7 M HCl?

pH = 6.791

What is the pH of 1.0 x 10-8 M HCl?

pH = 6.996

Why does the pH stop changing at a value of about 7?

Water has a pH of 7 due to autodissociation, so it is never possible to get a pH higher than 7 by addition of water.

– If [HA] < 10-7 water contributes the H+

14.5 pH of Weak Acids Except for the strong acids, most acids

do not ionize completely. These acids are called weak acids.

HF(aq) + H2O(l) ⇌ H3O+(aq) + F-(aq)

pH and Ka One way to find the value of Ka for a weak

acid is using concentration & pH data. The pH of 0.500 M HNO2 is 1.827. What is

Ka of HNO2? [H3O

+] = 10-1.827 = 0.0149 M HNO2 + H2O H3O

+ + NO2-

I 0.500 -- 0 0 C -0.0149 +0.0149 +0.0149 E 0.485 0.0149 0.0149 Ka = [H3O

+][NO2-]/[HNO2]

Ka = (0.0149)2/0.485 = 4.58 x 10-4

Weak Acid Scenario

Analyze a 0.50 M acetic acid HC2H3O2

solution (acetic acid: Ka = 1.8 x10-5)

Percent dissociation = amount dissociated x 100

initial concentration

For a weak acid percent dissociation increases as acid becomes more dilute.

Calculate the % dissociation of 1.00 M

and 0.100 M Acetic acid (Ka = 1.8 x 10-5 As [HA]0 decreases [H+] decreases but

% dissociation increases.

– Le Chatelier principle with dilution (pg 642)

Weak Acid Summary Ka will be small.

ALWAYS WRITE THE MAJOR SPECIES.

Determine whether most of the H+ will come from the acid or the water.

Compare Ka or Kw

Rest is just like chapter 13.

Some Weak Acids

A mixture of Weak Acids The process is the same.

Determine the major species.

The stronger acid will predominate.

Bigger Ka if concentrations are comparable

Calculate the pH of a mixture 1.20 M

HF (Ka = 7.2 x 10-4) and 3.4 M HOC6H5

(Ka = 1.6 x 10-10)

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The other way What is the Ka of a weak acid that is

8.1 % dissociated as 0.100 M solution?

Compare to Sample Exercise 14.11

14.6 Bases The OH-

is a strong base.

Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved.

The hydroxides of alkaline earths Ca(OH)2 etc. are strong dibasic bases, but they don’t dissolve well in water.

Used as antacids because [OH- ] can’t

build up.

14.6 Scenario

Analyze a 0.150 M ammonia solution

(ammonia: Kb = 1.8 x10-5)

Bases without OH-

Bases are proton acceptors.

NH3 + H2O NH4+ + OH-

It is the lone pair on nitrogen that accepts the proton.

Many weak bases contain N

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

Kb = [BH+][OH- ]

[B]

Strength of Bases Hydroxides are strong.

Others are weak.

Smaller Kb = weaker base.

Consider a solution of 4.0 M pyridine (Kb = 1.7 x 10-9)

– What are the major species present in solution?

– Determine the [OH-] and pH of the solution.

N:

Polyprotic Acid Scenario Describe the dissociation of the weak

diprotic acid, sulfurous acid, H2SO3

Polyprotic Acid Observations

From concentration and pH data, we can tell that all protons do not dissociate the same.

The first H+ comes of much easier than the second.

Ka for the first step is much bigger than Ka for the second.

14.7 Polyprotic acids Always dissociate stepwise.

Denoted Ka1, Ka2, Ka3, etc… (Table 14.4)

H2CO3 + H2O H+ + HCO3-

Ka1= 4.3 x 10-7 HCO3

- + H2O H+ + CO3-2

Ka2= 4.3 x 10-10

Conjugate base in 1st step is acid in 2nd.

Calculate the Concentration… …of all the ions in a solution of 5.00 M

sulfurous acid.

Ka1 = 1.5 x 10-2

Ka2 = 1.0 x 10-7

Find the pH of the solution

See Sample Exercise 14.15 for another example.

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Sulfuric acid, an interesting case Calculate the pH and concentration of all

species in a 1.5 M solution of H2SO4

– Pg 651, Table 14.4 contains Ka’s for polyprotics

Calculate the concentrations in a 1.5 x 10-2 M solution of H2SO4

See Sample Exc. 14.16 & 17

Summary pg. 655

– In first step H2SO4 is a strong acid.

– 2nd step, it is a weak acid: Ka2 = 1.2 x 10-2

14.8 Acid/Base Properties of Salts

Salts are ionic compounds. – Recall Model: Ionic compounds dissociate into

ions in solution. The ions move around independently.

These cations and anions can (but do not have to!) act as acids or bases depending on how they react with water (referred to as a hydrolysis reaction)

Ch 14, #99 discuss on WB’s

Salts that make Neutral Solutions Salts containing the cation from a

strong base and the anion from a strong acid are neutral.

Hydrolysis is not observed with ions derived from strong acids or bases: Cations of group I and II (except Be2+)

Anions: Cl-, Br-, I-, NO3-, ClO4

-

for example NaCl, KNO3

There is no equilibrium established from these ions.

pH of a salt scenario

Determine the pH of a solution of 1.00 M NaCN.

– First prepare a whiteboard of why NaCN might

affect the pH of a solution.

Ka of HCN is 6.2 x 10-10

Relative base strength:

OH- > CN- > H2O

The anion of a weak acid is a weak base.

The CN- ion competes with OH- for the H+

Basic Salts If the anion of a salt is the conjugate base

of a weak acid the solution will be basic.

Consider an aqueous solution of NaF

The aqueous species are Na+, F-, and H2O

F- + H2O HF + OH-

The equilibrium constant for hydrolysis is just a Ka or Kb, depending on the type of hydrolysis.

Basic Salts

Kb =[HF][OH-]

[F- ]

but Ka = [H+][F-]

[HF]

Ka x Kb = [HF][OH-] x [H+][F-] [F- ] [HF]

Basic Salts Ka x Kb = [HF][OH-] x [H+][F-]

[F- ] [HF]

Basic Salts Ka x Kb = [HF][OH-] x [H+][F-]

[F- ] [HF]

Basic Salts Ka x Kb = [HF][OH-] x [H+][F-]

[F- ] [HF]

Ka x Kb =[OH-] [H+]

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Basic Salts Ka x Kb = [HF][OH-] x [H+][F-]

[F- ] [HF]

Ka x Kb =[OH-] [H+]

Ka x Kb = Kw

pH of a salt Scenario 2 Calculate the pH of a solution of 0.40 M

NH4Cl (the Kb of NH3 1.8 x 10-5).

If the cation of a salt is the conjugate acid of a weak base the solution will be acidic.

The same reasoning as with bases leads to: Ka x Kb = KW

– Use acidic dissociation and Ka to determine pH

Other acidic salts are those of highly charged metal ions.

Acidic Salts – Metal ions

Ions are often modified when dissolved in solution. Examine the photos of Fe(III) salts and solutions. Why do they have different colors?

Fe(NO3)3.6H2O contains pink Fe(H2O)6

3+ Solutions may hydrolyze to give yellow

Fe(H2O)5OH2+ or even reddish brown Fe(H2O)3(OH)3

FeCl3.6H2O contains ions such as yellow

Fe(H2O)5Cl2+

Hydrolysis of Metals Hydrolysis is more

important for more highly charged ions and smaller ions.

Hydrogen from water easier to remove.

Hydrolysis of Metals Table below gives values of Ka for metal ions: Ion Radius Ka Na+ 95 pm 3.3 x 10-15

Li+ 60 pm 1.5 x 10-14

Be2+ 31 pm 3.2 x 10-7

Mg2+ 65 pm 3.8 x 10-12

Ba2+ 135 pm 1.5 x 10-14

Cr3+ 69 pm 9.8 x 10-5

Zr4+ 78 pm 6.0 x 10-1

Greater values of Ka for ions with larger charge and smaller size.

Hydrolysis of Metals Scenario Calculate the pH of a 0.20 M CrCl3

solution. (From the table, the Ka value for Cr3+ is 9.8 x 10-5)

Anion of weak acid, cation of weak base

If both of the ions have acid/base properties, compare the Ka of the cation to the Kb of the anion

Ka > Kb acidic

Ka < Kb basic

Ka = Kb neutral

Predict whether an aqueous solution of ammonium cyanide will be acidic, basic, or neutral.

Summary Table 14.6, pg 660

14.9 Structure and Acid base

Properties Draw lewis structures (that obey octet rule) for

the following oxyacids.

HClO4 Ka = Large (~107)

HClO3 Ka = ~1

HClO2 Ka = 1.2 x 10 –2

HClO Ka = 3.5 x 10 –8

Can you explain the relative strengths of these acids using your knowledge of atomic structure and bonding?

Strength of oxyacids

Cl O H

O

O

O

Electron Density

8

Strength of oxyacids

Cl O H

O

O

Electron Density

Strength of oxyacids

Electron Density

Cl O H O

Strength of oxyacids

Electron Density

Cl O H

Strength of oxyacids The more oxygen bonded to the central

atom, the more acidic the hydrogen.

The oxygens are electronegative

Pulls electron density away from hydrogen

HClO4 > HClO3 > HClO2 > HClO

– Remember that the H is attached to an oxygen

atom.

14.9 Structure and Acid base

Properties Draw lewis structures (that obey octet rule) for

the following oxyacids.

HOCl Ka = 4 x 10 –8

HOBr Ka = 2 x 10 –9

HOI Ka = 2 x 10 –11

HOCH3 Ka = ~ 10 –15

Can you explain the relative strengths of these acids using your knowledge of atomic structure and bonding?

14.9 Structure and Acid base

Properties Any molecule with an H in it is a

potential acid.

The stronger the X-H bond the less acidic (compare bond dissociation energies).

The more polar the X-H bond the stronger the acid (use electronegativities).

For oxyacids, the more polar the X-O-H bond, the stronger the acid.

14.10 Acid-Base Properties of

Oxides Non-metal oxides dissolved in water

can make acids.

SO3 (g) + H2O(l) H2SO4(aq)

Ionic oxides dissolve in water to produce bases.

CaO(s) + H2O(l) Ca(OH)2(aq)

14.11 Lewis Acids and Bases Most general definition.

Acids are electron pair acceptors.

Bases are electron pair donors.

B F

F

F

:N

H

H

H

Lewis Acids and Bases Boron triflouride wants more electrons.

B F

F

F

:N

H

H

H

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Lewis Acids and Bases Boron triflouride wants more electrons.

BF3 is Lewis acid NH3 is a Lewis base.

B F

F

F

N

H

H

H

Lewis Acids and Bases

Al+3 ( ) H

H O

Al ( ) 6

H

H O

+ 6

+3