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Acids and Bases

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Chapter 16 Acids and Bases – PowerPoint PPT presentation

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Title: Acids and Bases


1
Chapter 16 Acids and Bases
2
  • Drill
  • Determine which strong acid and strong base the
    following salts were derived from
  • LiCl
  • Ba3(PO4)2
  • CaSO4
  • Sr(NO3)2

3
Objectives iWBAT Distinguish between Arrhenius,
Bronsted Lowry and Lewis acids and bases. Write
conjugate acid/base pairs.
4
Arrhenius Definition
  • Definitions
  • Acids produce hydrogen ions in aqueous solution.
  • Bases produce hydroxide ions when dissolved in
    water.
  • Limited to aqueous solutions.
  • Uses only one kind of base (hydroxide).
  • NH3 ammonia could not be an Arrhenius base.

5
Bronsted-Lowry Definitions
  • DefinitionAn acid is a proton (H) donor and a
    base is a proton acceptor.
  • Acids and bases always come in pairs.
  • HCl is an acid.
  • When it dissolves in water, it gives its proton
    to water.
  • HCl(g) H2O(l) H3O Cl-
  • Water is a base that makes a hydronium ion

6
Example Reaction
  • HCl(g) H2O(l) ? H3O Cl-

Base
Hydronium Ion
Acid
7
Remember
  • Strong acids completely dissociate in water.
  • HCl H2O ? H3O 1 Cl-1
  • This reaction goes to completion and there is no
    HCl left in the solution.
  • Use a single direction arrow.

8
Strong Acids
9
Remember
  • Weak acids only partially dissociate.
  • CH3COOH NH3 ? CH3COO-1 NH41
  • This is an equilibrium reaction.
  • There are significant amounts of reactants and
    products in the solution.
  • Use a double headed arrow. ?

10
Hydroxides (and some oxides) are strong bases.
11
Weak Bases
  • All other common bases are weak.
  • Weak bases establish an equilibrium system like
    acids.
  • Some examples of weak bases are on the next
    slide.

12
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13
Acid Base Pairs
  • General equation
  • HA(aq) H2O(l) H3O(aq) A-(aq)
  • Acid Base Conjugate acid Conjugate
    base
  • This is an equilibrium situation.
  • There is competition for H between H2O and A-
  • The stronger base controls direction of the rxn.
  • If H2O is a stronger base it takes the H
  • Equilibrium would then move to right.

14
Practice
  • Try the Bronsted Lowry conjugate acid/base
    worksheet
  • Keep in mind
  • The acid turns into a base and
  • The base turns into an acid.

15
Day 2
16
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17
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18
This is an abbreviated chart with Common
Acids/Bases
19
Use the following reaction and the conjugate
acid/base chart to determine which direction the
equilibrium will lie.
  • CH3COOH NH3 ? CH3COO-1 NH41
  • CH3COOH is a stronger acid than NH41
  • NH3 is a stronger base than CH3COO-1
  • The equilibrium will favor the side in which the
    weaker acid and base a present.
  • Equilibrium will lie to the right.

20
Acid Dissociation Constant Ka
  • HA(aq) H2O(l) H3O(aq) A -1(aq)
  • Ka H3O1A-1 HA
  • H3O1 is often written H1 ignoring the water in
    equation (it is implied).
  • Since this is the equilibrium constant associated
    with weak acid dissociation, this particular Kc
    is most commonly called the acid dissociation
    constant Ka

21
Acid Dissociation Constant Ka
  • HA(aq) H(aq) A-(aq)
  • Ka HA- HA
  • We can write the expression for any acid.
  • Strong acids dissociate completely.
  • Equilibrium lies far to right.
  • Conjugate base must be weak.

22
Back to Pairs
  • Strong acids
  • Ka is large
  • H is equal to HA
  • A-1 is a weaker base than water
  • Weak acids
  • Ka is small
  • H ltltlt HA
  • A-1 is a stronger base than water

23
Types of Acids
  • Monoprotic Acids have only one hydrogen.
  • Polyprotic Acids more than 1 acidic hydrogen
    (diprotic, triprotic).
  • OxyacidsProton is attached to the oxygen of an
    ion.
  • Organic acidscontain the Carboxyl group -COOH
    with the H attached to O
  • Generally very weak.

24
  • Oxyacid examples
  • Sulfuric acid (H2SO4), phosphoric acid (H3PO4),
    and nitric acid (HNO3) are all oxyacids.
  • Organic acid examples
  • Lactic acid Uric acidAcetic acid Oxalic
    acidFormic acid Citric acid

25
Amphoteric
  • Amphoteric means that the substance can behave as
    both an acid and a base.
  • Water molecules interact with each other and
    ionize. At the same time, the ions in solution
    reform molecules of water as shown in the
    following reaction. (This means that water
    auto-ionizes)
  • 2H2O(l) H3O1(aq) OH-1 (aq)
  • KW H3OOH- HOH-

26
  • In pure water the concentrations of H3O1 and
    OH-1 will always be equal.
  • H OH- 1.0 x 10-7
  • At 25ºC KW 1.0 x10-14
  • Therefore
  • Neutral solution H OH- 1.0 x10-7
  • Acidic solution H gt OH-
  • Basic solution H lt OH-

27
pH
  • In 1909, Danish biochemist S. P. L Sorensen
    introduced the pH system.
  • pH representing power of hydrogen

28
pH
  • pH -logH
  • Used because H is usually very small
  • As pH decreases, H increases exponentially
  • Other equations
  • pOH -logOH-
  • pKa -log K

29
Sig Figs for pH
  • Sig figs the number of sig figs in the lead
    number is the number of decimal places for the pH
    value. (only the digits after the decimal place
    of a pH are significant)
  • H 1.0 x 10-8 pH 8.00 2 sig figs

30
Relationships
  • Derivation
  • KW HOH-
  • -log KW -log(HOH-)
  • -log KW -logH -logOH-
  • pKW pH pOH
  • 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.

31
Basic
Acidic
Neutral
32
Strong Acids
  • HBr, HI, HCl, HNO3, H2SO4, HClO4
  • These acids completely dissociate
  • Therefore, H HA
  • 10-14 HOH-

33
Weak Acids
  • Ka will be small.
  • ALWAYS WRITE THE MAJOR SPECIES.
  • It will be an equilibrium problem from the start.
  • Determine whether most of the H will come from
    the acid or the water.
  • Compare Ka or Kw
  • Rest is just like equilibrium chapter.

34
Example
  • Calculate the pH of 2.0 M acetic acid HC2H3O2
    with a Ka 1.8 x10-5
  • Calculate pOH, OH-, H

35
A Mixture of Weak Acids
  • The process is the same.
  • Determine the major species.
  • The stronger 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)

36
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
    .00100 M Acetic acid (Ka 1.8 x 10-5
  • As HA0 decreases H decreases but
    dissociation increases.
  • Le Chatelier

37
The Other Way
  • What is the Ka of a weak acid that is 8.1
    dissociated as 0.100 M solution?

38
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 dont
    dissolve well in water.
  • Used as antacids because OH- cant build up.

39
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 BHOH- B

40
Strength of Bases
  • Hydroxides are strong.
  • Others are weak.
  • Smaller Kb weaker base.
  • Calculate the pH of a solution of 4.0 M pyridine
    (Kb 1.7 x 10-9)

N
41
Polyprotic Acids
  • Always dissociate stepwise.
  • The first H comes of much easier than the
    second.
  • Ka for the first step is much bigger than Ka for
    the second.
  • Denoted Ka1, Ka2, Ka3

42
Polyprotic Acids
  • What does K stand for?
  • Is it easier to remove the first or second
    ionizable proton?
  • Is is easier to remove the first.
  • The K values become successively smaller as
    successive protons are removed.
  • You will need to do 2 or more ice boxes.

43
Polyprotic Acid
  • H2CO3 H HCO3-1 Ka1 4.3 x 10-7
  • HCO3-1 H CO3-2 Ka2 4.3 x 10-10
  • Base in first step is acid in second.
  • In calculations we can normally ignore the second
    dissociation.

44
Calculate the Concentration
  • Of all the ions in a solution of 1.00 M Arsenic
    acid H3AsO4
  • Ka1 5.0 x 10-3
  • Ka2 8.0 x 10-8
  • Ka3 6.0 x 10-10

45
Sulfuric Acid is Special
  • In first step it is a strong acid.
  • Ka2 1.2 x 10-2
  • Calculate the concentrations in a 2.0 M solution
    of H2SO4
  • Calculate the concentrations in a 2.0 x 10-3 M
    solution of H2SO4

46
Salts
  • A salt is an ionic compound formed by the
    reaction between an acid and a base.
  • Salts are strong electrolytes that completely
    dissociate into ions in water.
  • Salts of the cation of strong bases and the anion
    of strong acids are neutral.
  • for example NaCl, KNO3

47
Salts that Produce Neutral Solutions
48
Basic Salts
  • If the anion of a salt is the conjugate base of a
    weak acid - basic solution.
  • In an aqueous solution of NaF
  • The major species are Na, F-, and H2O
  • F- H2O HF OH-
  • Kb HFOH- F-
  • but Ka HF- HF

49
Basic Salts
  • Ka x Kb HFOH- x HF- F-
    HF

50
Basic Salts
  • Ka x Kb HFOH- x HF- F-
    HF
  • Ka x Kb OH- H
  • Ka x Kb KW

51
Ka tells us Kb
  • The anion of a weak acid is a weak base.
  • Calculate the pH of a solution of 1.00 M NaCN. Ka
    of HCN is 6.2 x 10-10
  • The CN- ion competes with OH- for the H

52
Acidic Salts
  • A salt with the cation of a weak base and the
    anion of a strong acid will be basic.
  • The same development as bases leads to
  • Ka x Kb KW
  • Calculate the pH of a solution of 0.40 M NH4Cl
    (the Kb of NH3 1.8 x 10-5).
  • Other acidic salts are those of highly charged
    metal ions.

53
Anion of weak acid, cation of weak base
  • Ka gt Kb acidic
  • Ka lt Kb basic
  • Ka Kb Neutral

54
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).
  • The more polar H-O-X bond -stronger acid.

55
Strength of Oxyacids
  • The more oxygen hooked to the central atom, the
    more acidic the hydrogen.
  • HClO4 gt HClO3 gt HClO2 gt HClO
  • Remember that the H is attached to an oxygen
    atom.
  • The oxygens are electronegative
  • Pull electrons away from hydrogen

56
Strength of Oxyacids
Electron Density
57
Strength of Oxyacids
Electron Density
O
58
Strength of Oxyacids
Electron Density
O
O
59
Strength of Oxyacids
Electron Density
O
O
O
60
Hydrated Metals
  • Highly charged metal ions pull the electrons of
    surrounding water molecules toward them.
  • Make it easier for H to come off.

H
Al3
O
H
61
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)

62
Lewis Acids and Bases
  • Most general definition.
  • Acids are electron pair acceptors.
  • Bases are electron pair donors.

F
H
B
F
N
H
F
H
63
Lewis Acids and Bases
  • Boron triflouride wants more electrons.

F
H
B
F
N
H
F
H
64
Lewis Acids and Bases
  • Boron triflouride wants more electrons.
  • BF3 is Lewis base NH3 is a Lewis Acid.

F
H
F
H
B
N
F
H
65
Lewis Acids and Bases
(
H
Al3
6
O
H
3
(
H
Al
O
H
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