Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

96
Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1

Transcript of Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Page 1: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Chapter 19

Acids & Bases

Section 19.1

Acid – Base Theories

1

Page 2: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Acids

· vinegar citrus fruits· carbonated drinks car battery· lemon juice tea

Bases

· calcium hydroxide in mortar antacids· household cleaning agents

2

Page 3: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Properties of Acids

1. Give foods a tart or sour taste

e.g. lemon & vinegar for example

2.Aqueous solutions of acids are electrolytes (conduct electricity)

3.Acids cause certain chemical indicators to change color.

4.Acid + Base Salt + water

3

Page 4: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Properties of Bases

1.Bases have a bitter taste

e.g. soap

2.Bases have a slippery feel

3.Aqueous solutions of bases are electrolytes (conduct electricity)

4.Bases cause certain chemical indicators to change color.

5.Acid + Base Salt + water

4

Page 5: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Arrhenius Acids & Bases

• Acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution.

• Bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution

In 1887, Swedish chemist Svante Arrhenius proposed a revolutionary way of defining and thinking about acids and bases

5

Page 6: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Monoprotic acids: acids that contain one ionizable hydrogen

HNO3 – nitric acid

Diprotic acids: acids that contain two ionizable hydrogens

H2SO4 – sulfuric acid

Triprotic acids: acids that contain three ionizable hydrogens

H3PO4 – phosphoric acid

6

Page 7: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

• Not all compounds that contain hydrogen are acidse.g. CH4 – methane has weak polar C – H bonds and no ionizable hydrogens. Not an acid.

• Not all hydrogens in an acid may be released as hydrogen ions.

• Only hydrogens in very polar bonds are ionizable. In the case where hydrogen is joined to a very electronegative element.

e.g. HCl hydrogen chloride very polar covalent molecule

7

Page 8: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

• When HCl dissolves in water, it releases hydrogen ions because the hydrogen ions are stabilized by solvation.

H2O

H – Cl (g) H+ (aq) + Cl- (aq)

Hydrogen Hydrogen Chloride chloride ion ion

Ionizes to form an aqueous solution of hydronium ions and chloride ions

HCl + H2O H3O+ + Cl-

8

Page 9: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

• Ethanoic acid CH3COOH is a monoprotic acid due to its structure

H O

H C C O H

H

The three H attached to the carbon are in weak polar bonds. They do not ionize.

Only the H bonded to the highly electronegative O can be ionized

9

Page 10: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Sodium hydroxide dissociates into sodium ions and hydroxide ions in aqueous solution.

H2O

NaOH (s) Na+ (aq) + OH- (aq)

Sodium Sodium HydroxideHydroxide Ion ion

Potassium hydroxide dissociates into potassium ions and hydroxide ions in aqueous solution.

H2O

KOH (s) K+ (aq) + OH- (aq)

Potassium Potassium HydroxideHydroxide Ion ion

10

Page 11: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Arrhenius Bases Group one, the alkali metals, react with water to produce solutions that are basic.

Group one metals are very soluble in water and can produce concentrated solutions.

Group two metals are not very soluble in water. Their solutions are always very dilute.

11

Page 12: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Bronsted – Lowry Acids and Bases

The Bronsted – Lowry theory defines

Acid: a hydrogen-ion donor

Base: a hydrogen-ion acceptor

All acids and bases included in the Arrhenius theory are also acids and bases according to the Bronsted-Lowry theory.

12

Page 13: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)

base acid conjugate conjugate acid base

• Ammonia is the hydrogen-ion acceptor therefore it is a base.

• Water is the hydrogen-ion donor and therefore it is an acid.

• Hydrogen ions are transferred from water to ammonia, which causes the hydroxide-ion concentration to be greater than it is in pure water.

• When ammonia dissolves and reacts with water, NH4+

is the conjugate acid of the base NH3.

• OH- is the conjugate base of acid H2O

13

Page 14: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Conjugate Acids and Bases

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

acid base conjugate conjugate acid base

•HCl is the hydrogen-ion donor – thus it is an acid.

•Water is the hydrogen-ion acceptor – thus it is a base

14

Page 15: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Conjugate Acid – Base Pair

Conjugate acid: the particle formed when a base gains a hydrogen ion

Conjugate base: the particle that remains when an acid has donated a hydrogen ion.

Conjugate acids and bases are always paired with a base or an acid, respectively.

Conjugate acid-base pairs consists of two substances related by the loss or gain of a single hydrogen ion.

15

Page 16: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Common Conjugate Acid – Base Pairs

Acid Conjugate BaseHCl Cl-

H2SO4 HSO4-

H3O+ H2O

HSO4- SO4

2-

CH3COOH CH3COO-

H2CO3 HCO3-

HCO3- CO3

2-

NH4+ NH3

H2O OH-

16

Page 17: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Bronsted – Lowry Acids and Bases

A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion (H3O+)

Amphoteric – a substance that can act as both an acid and a base

e.g. water

H2SO4 + H2O H3O+ + HSO4-

NH3 + H2O NH4+ + OH-

17

Page 18: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Lewis Acids and Bases

Gilbert Lewis proposed a third Acid Base theory

Acid – accepts a pair of electrons during a reaction

Base – donates a pair of electrons during a reaction

Concept is more general than either the Arrhenius theory or the Bronsted-Lowry theory.

18

Page 19: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Lewis Acids and Bases

Lewis Acid – a substance that can accept a pair of electrons to form a covalent bond.

Lewis Base – a substance that can donate a pair of electrons to form a covalent bond.

..H+ : O – H :O:

..

H HLewis LewisAcid Base

19

Page 20: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Acid Base Definitions

Type Acid Base

Arrhenius H+ producer OH- producer

Bronsted - Lowry H+ donor H+ acceptor

LewisElectron-pair

acceptorElectron-pair

donor

20

Page 21: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

End of Section 19.121

Page 22: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Section 19.2

Hydrogen Ions & Acidity

22

Page 23: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Hydrogen Ions From Water

A water molecule that loses a hydrogen ion becomes a negatively charged hydroxide ion

A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion

H2O (l) OH- (aq) + H+ (aq) Hydroxide ion Hydrogen ion

Self ionization of water – the reaction in which water molecules produce ions

23

Page 24: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Self Ionization of Water

Hydrogen ions in aqueous solution have several names. • Some chemists call them protons • Some chemists call them hydrogen ions or

hydronium ions.

For our purposes, either H+ or H3O+ will represent hydrogen ions in aqueous solution.

H2O + H2O H3O+ + OH-

24

Page 25: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

• In pure water at 25˚C, the equilibrium concentration of hydrogen ions and hydroxide ions are each only 1 x 10-7.

• In other words the concentration of OH- and H+ are equal in pure water

Any aqueous solution in which H+ and OH- are equal is a neutral solution.

When [H+] increases [OH-] decreases

When [H+] decreases [OH-] increases

25

Page 26: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

For aqueous solutions, the product of the hydrogen ion concentration and the hydroxide ion concentration equals 1.0 x 10-14

[H+] x [OH-] = 1 x 10-14

1 x 10-7 x 1 x 10-7 = 1 x 10-14

This equation is true for all dilute aqueous solutions at 25˚C.

Ion-Product Constant for Water (Kw) the product of the concentrations of the hydrogen ions and hydroxide ions in water

Kw = [H+] x [OH-] = 1.0 x 10-1426

Page 27: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

But not all solutions are neutral

When some substances dissolve in water, they release hydrogen ions.

When hydrogen chloride dissolves in water, it forms hydrochloric acid.

H2O

HCl (g) H+ (aq) + Cl- (aq)

27

Page 28: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Ion Product Constant for Water

In the previous HCl solution, the hydrogen-ion concentration is greater than the hydroxide-ion concentration.

Acidic Solution:A solution in which [H+] is greater than [OH-].

The [H+] of an acidic solution is greater than 1 x 10-7 M

28

Page 29: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

When sodium hydroxide dissolves in water, it forms hydroxide ions in solution.

H20

NaOH(s) Na+(aq) + OH-(aq)

In the above solution, the hydrogen-ion concentration is less than the hydroxide-ion concentration.

Basic Solution:A solution in which [H+] is less than [OH-]

The [H+] of a basic solution is less than 1 x 10-7

Basic solutions are also known as alkaline solutions. 29

Page 30: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

The pH Concept

The pH scale was proposed by Danish Scientist Soren Sorensen in 1909.

The pH scale is used to express [H+]

1 2 3 4 5 6 7 8 9 10 11 12 13 14

Strongly Neutral Strongly Acidic Basic

30

Page 31: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Calculating pH

The pH of a solution is the negative logarithm of the hydrogen-ion concentration.

pH = - log [H+]

31

Page 32: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

In neutral solution, the [H+] = 1 x 10-7M

pH = - log [H+]

pH = - log (1 x 10-7)

pH = 7

32

Page 33: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Classifying Solutions

A solution in which [H+] is greater than 1 x 10-7 has a pH less than 7 is called acidic.

A solution in which [H+] is less than 1 x 10-7 has a pH greater than 7 is called basic.

The pH of pure water or a neutral aqueous solution is 7

Acidic solution: pH < 7.0 [H+] > 1 x 10-7 M

Neutral solution: pH = 7.0 [H+] = 1 x 10-7 M

Basic solution: pH > 7.0 [H+] < 1 x 10-7 M

33

Page 34: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.
Page 35: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

pH can be read from the value of [H+] if it is written in scientific notation and has a coefficient of 1.

Then the pH of the solution equals the exponent, with the sign changed from minus to plus

e.g. [H+] = 1 x 10-2 has a pH of 2

e.g. [H+] = 1 x 10-13 has a pH of 13

35

Page 36: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

If the pH is an integer, it is also possible to directly write the value of [H+].

pH = 9 ; then [H+] = 1 x 10-9 M

pH = 4 ; then [H+] = 1 x 10-4M

36

Page 37: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Acidic solution: pOH > 7 [OH-] < 1 x 10-7 M

Neutral solution: pOH = 7 [OH-] = 1 x 10-7 M

Basic solution: pOH < 7 [OH-] > 1 x 10-7 M

The pOH of a solution equals the negative logarithm of the hydroxide-ion concentration

pOH = - log [OH-]

A neutral solution has a pOH of 7

37

Page 38: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

pH and pOH Relationship

pOH + pH = 14

pH = 14 – pOH

pOH = 14 – pH 38

Page 39: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

39

Page 40: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Problem Example Colas are slightly acidic. If the [H+] in a solution is 1.0 X 10 - 5 M , is the solution acidic, basic or neutral. What is the [OH-] of this solution?

[H+] = 1.0 X 10 - 5 M which is greater than 1.0 X 10 -7 M

so the solution is acidic

Kw = [OH-] x [H+] = 1.0 X 10 - 14

[OH-] = 1.0 X 10 - 14 ÷ [H+]

[OH-] = 1.0 X 10 - 14 ÷ 1.0 X 10 - 5

[OH-] = 1.0 X 10 - 9 M40

Page 41: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Problem Example

What is the pH of a solution with a hydrogen-ion concentration of 4.2 x 10 - 10 M?

pH = - log [H+]

pH = - log (4.2 x 10 - 10)

pH = 9.38

41

Page 42: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Problem Example pH of an unknown solution is 6.35. What is its hydrogen-ion concentration?

pH = -log [H+]

6.35 = -log [H+]

- 6.35 = log [H+]

Using calculator find the anti log of - 6.35

[H+] = 4.5 x 10 - 7 M42

Page 43: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Problem Example

What is the pH of a solution if the [OH-] = 4.0 X10 – 11 M?

Kw = [H+] x [OH-] = 1 x 10 -14

[H+] = 1 x 10 -14 ÷ [OH-]

[H+] = 1 x 10 -14 ÷ 4.0 x 10 -11

[H+] = 2.5 x 10 - 4 M

43

Page 44: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Problem Example (con’t)

What is the pH of a solution if [OH-] = 4.0 X 10 - 11 M?

pH = - log [H+]

pH = - log (2.5 x 10 - 4)

pH = 3.60

44

Page 45: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Acid – Base IndicatorsIndicator - is an acid or a base that undergoes dissociation in a known pH range

An indicator is a valuable tool for measuring pH because its acid form and base form have different color in solution.

For each indicator, the change from dominating acid form to dominating base form occurs in a narrow range of approximately two pH units.

Within this range, the color of the solution is a mixture of the colors of the acid and the base forms.

Knowing the pH range over which this color change occurs, can give you a rough estimate of the pH of the solution.

45

Page 46: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

46

Page 47: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

End of Section 19.2

47

Page 48: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Section 19.3

Hydrogen Ions & Acidity

48

Page 49: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Strong Acids

Acids are classified as strong or weak depending on the degree to which they ionize in water.

• In general, strong acids are completely ionized in

aqueous solution.

HNO3 - nitric acid HCl - hydrochloric acidH2SO4 - sulfuric acid HClO4 - perchloric acid HBr - hydrobromic acidHI - hydroiodic acid

HCl(g) + H2O(l) H3O +(aq) + Cl ¯(aq)

Page 50: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Weak Acids

Weak acids ionize only slightly in aqueous solution. • Some Weak Acids

Acetic Acid CH3COOH

Boric Acid H3BO3 (all three are weak)

Phosphoric Acid H3PO4 (all three are weak)

Sulfuric Acid HSO4- (first ionization is strong)

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO ¯

(aq)

ethanoic acid water hydronium ethanoate

ion ion

Page 51: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Acid Strength

A strong acid completely dissociates in water ([H3O+] is high).

A weak acid remains largely undissociated. ([H3O+] is low).

Page 52: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Equilibrium Constant (Keq)

Write the equilibrium-constant expression from the balanced chemical equation.

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO ¯(aq)

Keq = [H3O+] x [ CH3COO ¯]

[H3COOH] x [H2O] [H2O] constant in dilute solutions

Page 53: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Acid Dissociation Constant (Ka)

Ka = Ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form.

CH3COOH (aq) + H2O (l) H3O +(aq) + CH3COO ¯(aq)

Acid Dissociation Constant Ka = [H3O+] x [ CH3COO¯] [CH3COOH]

Page 54: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Acid Dissociation Constant (Ka)

Acid dissociation constant reflects the fraction of an acid in the ionized form.

(Ka sometimes called ionization constant) If the value of the Ka is small, then the degree of

dissociation or ionization of the acid in the solution is small.

Weak acids – small Ka valuesStronger the acid – larger the Ka

Page 55: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Acid Dissociation Constant (Ka)

Nitrous acid (HNO2) has a Ka of 4.4 x 10¯4

Acetic acid (CH3COOH) has a Ka of 1.8 x 10¯5

Nitrous acid is more ionized in solution and a stronger acid

Page 56: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Acids

Strong AcidsHave high [H3O+] Large dissociation constant (Ka)

Weak AcidsHave low [H3O+] Small dissociation constant (Ka)

Page 57: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

AcidsDiprotic and triprotic acids lose their hydrogens one

at a time.

Each ionization reaction has a separate dissociation constant.

H3PO4 – 3 separate dissociation constants.

Page 58: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

58

A 0.012 M solution of formic acid HCOOH is partially ionized. [H+] is 1.02 x 10-4 M. what is the acid dissociation constant (Ka) of formic acid?

Concentration [HCOOH] [H+] [HCOO¯ ]

Initial 0.012 0 0

Change – 1.02 x 10 - 4 1.02 x 10 - 4 1.02 x 10 – 4

Equilibrium 0.011989 1.02 x 10 - 4 1.02 x 10 – 4

Substitute the equilibrium values into the expression for Ka :

Ka = [H+] [HCOO¯ ]

[HCOOH]=

(1.02 x 10 – 4) (1.02 x 10 – 4)

0.011989= 8.68 x10 - 7

Classwork page: 610 # 22 & 23

Page 59: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Strong Bases

Bases are classified as strong or weak depending on the degree to which they ionize in water.

• In general, strong bases are completely ionized in

aqueous solution.

Ca(OH)2 - calcium hydroxide NaOH - sodium hydroxideKOH - potassium hydroxide

NaOH(s) + H2O(l) Na +(aq) + OH ¯(aq)

Page 60: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Weak Bases

Weak bases ionize only slightly in aqueous solution. • Some Weak bases

Ammonia NH3

Methylamine CH3NH2

NH3 (aq) + H2O(l) NH4+(aq) + OH ¯ (aq)

ammonia water ammonium ion hydroxide ion

Page 61: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Base Dissociation Constant (Kb)

• Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solution.

• Some strong bases are not very soluble in water (calcium hydroxide and magnesium hydroxide)

• Small amounts that do not dissolve dissociate completely

Weak bases react with water to form the hydroxide

ion and the conjugate acid of the base.

NH3(aq) + H2O(l) NH4+ (aq) + OH¯ (aq)

Ammonia Water Ammonium Ion Hydroxide ion

Page 62: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Base Dissociation Constant (Kb)

Only about 1% of ammonia is present as NH4+

Base Dissociation Constant (Kb)Kb = [NH4

+] x [OH¯ ] [NH3]

NH3(aq) + H2O(l) NH4+(aq) + OH ¯(aq)

Ammonia Water Ammonium Ion Hydroxide ion

Page 63: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Concentration and Strength

The words concentrated and dilute indicate how much of an acid or base is dissolved in solution.Number of moles of the acid or base in a given volume.

The words strong and weak refer to the extent of ionization or dissociation of an acid or base.

How many of the particles ionize or dissociate into ions

A sample of HCl added to a large volume of water becomes more dilute, but it is still a strong acid.

Vinegar is a dilute solution of a weak acid.

Page 64: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

End of section 19.3

Page 65: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Section 19.4

Neutralization Reactions

65

Page 66: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Acid – Base Reactions

If you mix a solution of a strong acid containing hydronium ions with a solution of a strong base that has an equal number of hydroxide ions, a neutral solution results.

Final solution has properties that are characteristic of neither an acidic nor a basic solution. HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

H2SO4 (aq) + 2KOH (aq) K2SO4 (aq) + H2O (l)

Page 67: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Neutralization Reactions

Reactions of weak acids and weak bases do not usually produce a neutral solution.

In general, reactions with which an acid and a base react in an aqueous solution to produce a salt and water are called neutralization reactions.

Page 68: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Making Salts

Prepare potassium chloride by mixing equal molar quantities of hydrochloric acid and potassium hydroxide.

HCl + KOH KCl + H2O

Heating the solution to evaporate the water will leave the salt potassium chloride.

In general, the reaction of an acid with a base produced water and salt

Page 69: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

TitrationThe number of moles of hydrogen ions provided by the acid are equivalent to the number of hydroxide ions provided by the base.

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l) 1 mole 1 mole 1 mole 1 mole

H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + 2H2O (l) 1 mole 2 mole 1 mole 2 mole

When an acid & base are mixed, the Equivalence point is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions.

Page 70: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Sample Problem

How many moles of sulfuric acid are required to neutralize 0.50 mol of sodium hydroxide?

H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + 2H2O (l)

Page 71: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Practice ProblemHow many moles of potassium hydroxide are needed to completely neutralize 1.56 mol of phosphoric acid?

H3PO4 (aq) + 3KOH (aq) K3PO4 (aq) + 3H2O (l)

Page 72: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

TitrationYou can determine the concentration of acid or base in a solution by performing a neutralization reaction.

You must use an appropriate acid-base indicator to show when neutralization has occurred. In the lab, typically phenolphthalein for acid base neutralization reactions.

Solutions that contain phenolphthalein turn from colorless to deep pink as the pH of the solution changes from acidic to basic.

Page 73: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Titration

Measured volume of an acid solution of unknown concentration is added to a flask

Page 74: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Titration

Several drops of the indicator are added to the solution while the flask is swirled

Page 75: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Titration

Measured volumes of the base of known concentration are mixed into the acid until the indicator just barely changes color.

Page 76: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

TitrationTitration – the process of adding a known amount of solution of known concentration to determine the concentration of another solution.

Standard solution – the solution of known concentration

End point – the point at which the indicator changes color, the point of neutralization

You can also use titration to find the concentration of a base using a standard acid.

Equivalence point – the point in a titration where the number of moles of hydrogen ions = number of moles of hydroxide ions.

Page 77: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

A 25ml solution of H2SO4 is completely neutralized by 18ml of 1.0M NaOH. What is the concentration of the H2SO4 solution?

H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + 2H2O (l)

Page 78: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Practice Problem

How many milliliters of 0.45M HCl will neutralize 25.0ml of 1.00M KOH?

Page 79: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Practice ProblemWhat is the molarity of H3PO4 if 15.0 ml is completely neutralized by 38.5 ml of 0.150 M Ca(OH)2?

Page 80: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

End of section 19.4

Page 81: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Section 19.5

Salts in Solution

81

Page 82: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Salt Hydrolysis

A salt consists of an anion from an acid and a cation from a base.

The salt forms as a result of a neutralization reaction

Although solutions of many salts are neutral, some are acidic and others are basic.

Page 83: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Salt Hydrolysis

Salt Hydrolysis – the cations or anions ofa dissociated salt remove hydrogen ions from or donate hydrogen ion to water.

Hydrolyzing salts are usually derived from a strong acid and weak base or from a weak acid and a strong base.

In general, salts that produce acidic solutions contain positive ions that release protons to water.

Salts that produce basic solutions contain negative ions that attract protons from water.

Page 84: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Salt Hydrolysis

CH3COONa (aq) CH3COO ¯ (aq) + Na+ (aq)Sodium ethanoate ethanoate ion sodium ion

CH3COONa is the salt from a weak acid CH3COOH and a strong base NaOH

In solution, the salt is completely ionized.

Page 85: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Salt HydrolysisSalt Hydrolysis – the cations or anions of a dissociated

salt remove hydrogen ions from or donate hydrogen ion to water.

CH3COO¯ (aq) + H2O (l) CH3COOH (aq) + OH ¯(aq) BL base BL acid makes hydrogen-ion hydrogen-ion solution acceptor donor basic

This process is called hydrolysis because it splits a hydrogen ion off a water molecule.

Resulting solution contains a hydroxide-ion concentration greater than the hydrogen-ion concentration. Thus the solution is basic

Page 86: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Salt Hydrolysis

NH4Cl (aq) NH4+ (aq) + Cl ¯ (aq)

Ammonium Ammonium ion Chloride ion chloride

NH4Cl is the salt from a strong acid (hydrochloric acid, HCl) and a weak base (ammonia, NH3)

In solution the salt is completely ionized.

Page 87: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Salt Hydrolysis

NH4+ (aq) + H2O (l) NH3 (aq) + H3O+

(aq)

BL acid BL baseHydrogen-ion Hydrogen-ion donor acceptor

This process is also called hydrolysis because it splits a hydrogen ion off a water molecule.

Resulting solution contains a hydrogen-ion concentration greater than the hydroxide-ion concentration. Thus the solution is acidic

Page 88: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Salt Hydrolysis

Equivalence Point

Strong Acid Strong Base pH= 7 neutral

Weak AcidStrong Base

pH > 7 basic

Strong AcidWeak Base

pH < 7 acidic

Equivalence point – the point in a titration where the number of moles of hydrogen ions = number of moles of hydroxide ions

Page 89: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Buffers

Buffer – a solution in which the pH remains relatively constant when small amounts of acid or base are added.

A buffer is a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts.

A buffer solution is better able to resist drastic changes in pH than is pure water.

Page 90: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Buffers

A solution of ethanoic acid (CH3COOH) and sodium ethanoate (CH3COONa) is an example of a typical buffer.

CH3COOH and CH3COO - (source is the completely ionized CH3COONa) act as reservoirs of neutralizing power.

Page 91: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Buffers

CH3COO - (aq) + H+ (aq) CH3COOH (aq)

ethanoate ion Hydrogen ion ethanoic acid

When an acid is added to the solution, the ethanoate ions act as a hydrogen-ion sponge.

CH3COOH (aq) + OH - (aq) CH3COO - (aq) + H2O (l) Ethanoic acid hydroxide ion ethanoate ion water

When a base is added to the solution, the ethanoic acid and the hydroxide ions react to produce water and the ethanoate ion.

Page 92: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Buffers

The ethanoate ion is not strong enough base to accept hydrogen ions from water extensively.

The buffer solution cannot control the pH when too much acid is added, because no more ethanoate ions are present to accept hydrogen ions.

Buffer also become ineffective when too much base is added. No more ethanoic acid molecules are present to donate hydrogen ions.

Page 93: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Buffers

When too much acid or base is added, the buffer capacity is exceeded.

Buffer capacity – the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs.

Page 94: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Buffers

When a base is added to a buffered solution, the acidic form removes hydroxide ions from the solution.

When an acid is added to a buffered solution, the basic form removes hydrogen ions from the solution.

Page 95: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

Buffers & Your BloodYour body function properly only when the pH of your blood lies between 7.35 and 7.45

Your blood contains buffers (hydrogen carbonate ions and carbonic acid)

HCO3- (aq) + H+ (aq) H2CO3 (aq)

Hydrogen Hydrogen ion Carbonic acidcarbonate ion

As long as there are hydrogen carbonate ions available, the excess hydrogen ions are removed, and the pH of the blood changes very little.

Page 96: Chapter 19 Acids & Bases Section 19.1 Acid – Base Theories 1.

End of Section Chapter 19