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Chapter 18 Acids and Bases. Items from Chapter 17. Reversible Reactions - p. 416 In a reversible reaction, the reactions occur simultaneously in both directions double arrows used to indicate this 2SO 2(g) + O 2(g)  2SO 3(g)

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items from chapter 17
Items from Chapter 17...
  • Reversible Reactions - p. 416
    • In a reversible reaction, the reactions occur simultaneously in both directions
    • double arrows used to indicate this

2SO2(g) + O2(g) 2SO3(g)

    • In principle, almost all reactions are reversible to some extent
items from chapter 171
Items from Chapter 17...
  • Le Chatelier’s Principle - p.421
    • If a stress is applied to a system in dynamic equilibrium, the system changes to relieve the stress.
    • Stresses that upset the equilibrium in a chemical system include: changes in concentration, changes in temperature, and changes in pressure
items from chapter 172
Items from Chapter 17...
  • Equilibrium Constants (Keq) - p. 418
    • Chemists generally express the position of equilibrium in terms of numerical values
    • These values relate to the amounts of reactants and products at equilibrium
items from chapter 173
Items from Chapter 17...
  • Equilibrium Constants - p. 418
    • consider this reaction:

aA + bB  cC + dD

    • The equilibrium constant (Keq) is the ratio of product concentration to the reactant concentration at equilibrium, with each concentration raised to a power (= the coefficient)
items from chapter 174
Items from Chapter 17...
  • Equilibrium Constants - p. 418
    • consider this reaction:

aA + bB  cC + dD

    • Thus, the “equilibrium constantexpression” has the general form:

[C]c x [D]d

[A]a x [B]b

( [ ] = molarity )

Keq =

items from chapter 175
Items from Chapter 17...
  • Equilibrium Constants - p. 418
    • the equilibrium constants provide valuable information, such as whether products or reactants are favored:

Keq > 1, products favored at equilibrium

Keq < 1, reactants favored at equilibrium

  • Sample Problem 17-8, p. 419
section 18 1 describing acids and bases
Section 18.1Describing Acids and Bases
  • OBJECTIVES:
    • List the properties of acids and bases.
section 18 1 describing acids and bases1
Section 18.1Describing Acids and Bases
  • OBJECTIVES:
    • Name an acid or base, when given the formula.
properties of acids
Properties of acids
  • Taste sour (don’t try this at home).
  • Conduct electricity.
    • Some are strong, others are weak electrolytes.
  • React with metals to form hydrogen gas.
  • Change indicators (blue litmus to red).
  • React with hydroxides to form water and a salt.
properties of bases
Properties of bases
  • React with acids to form water and a salt.
  • Taste bitter.
  • Feel slippery (don’t try this either).
  • Can be strong or weak electrolytes.
  • Change indicators (red litmus turns blue).
names and formulas of acids
Names and Formulas of Acids
  • An acid is a chemical that produces hydrogen ions (H1+) when dissolved in water
  • Thus, general formula = HX, where X is a monatomic or polyatomic anion
  • HCl(g) named hydrogen chloride
  • HCl(aq) is named as an acid
  • Name focuses on the anion present
slide13

Naming Acid

Hydrogen _______ide becomes hydro____ic acid

Hydrogen_______ate becomes _________ic acid

Hydrogen_______ite becomes _______ous acid

chlor

chlor

names and formulas of acids1
Names and Formulas of Acids
  • When anion ends with -ide, the acid starts with hydro-, and the stem of the anion has the suffix -ic followed by the word acid

2. When anion ends with -ite, the anion has the suffix -ous, then acid

3. When anion ends with -ate, the anion suffix is -ic and then acid

names and formulas of bases
Names and Formulas of Bases
  • A base produces hydroxide ions (OH1-) when dissolved in water.
  • Named the same way as any other ionic compound
    • name the cation, followed by anion
  • To write the formula: write symbols; write charges; then cross (if needed)
section 18 2 hydrogen ions and acidity
Section 18.2Hydrogen Ions and Acidity
  • OBJECTIVES:
    • Given the hydrogen-ion or hydroxide-ion concentration, classify a solution as neutral, acidic, or basic.
section 18 2 hydrogen ions and acidity1
Section 18.2Hydrogen Ions and Acidity
  • OBJECTIVES:
    • Convert hydrogen-ion concentrations into values of pH, and hydroxide-ion concentrations into values of pOH.
hydrogen ions from water
Hydrogen Ions from Water
  • Water ionizes, or falls apart into ions:

H2O ® H1+ + OH1-

  • Called the “self ionization” of water
  • Occurs to a very small extent:

[H1+ ] = [OH1-] = 1 x 10-7 M

  • Since they are equal, a neutral solution results from water
  • Kw = [H1+ ] x [OH1-] = 1 x 10-14 (mol/L)2
  • Kw is called the “ion product constant”
ion product constant
Ion Product Constant
  • H2O H+ + OH-
  • Kw is constant in every aqueous solution: [H+] x [OH-] = 1 x 10-14
  • If [H+] >1x 10-7 then [OH-] < 1x10-7
  • If [H+] < 1x10-7 then [OH-] > 1x10-7
  • If we know one,the other can be determined
  • If [H+] > 10-7, it is acidic and [OH-] < 10-7
  • If [H+] < 10-7, it is basic and [OH-] > 10-7
  • Basic solutions are also called “alkaline”
logarithms and the ph concept
Logarithms and the pH concept
  • Logarithms are powers of ten.
    • Know how to use the log buttons on your calculator
  • definition: pH = -log[H+]
  • in neutral pH = -log(1 x 10-7) = 7
  • in acidic solution [H+] > 10-7
  • pH < 7 (from 0 to 7 is the acid range)
  • in base, pH > 7 (7 to 14 is base range)
ph and poh
pH and pOH
  • pOH = -log [OH-]
  • [H+] x [OH-] = 1 x 10-14 M2
  • pH + pOH = 14
  • Thus, a solution with a pOH less than 7 is basic; with a pOH greater than 7 is an acid
slide22

pH

0

1

3

5

7

9

11

13

14

Acidic

Neutral

Basic

14

13

11

9

7

5

3

1

0

pOH

slide23

100

10-1

10-3

10-5

10-7

10-9

10-11

10-13

10-14

[H+]

Acidic

Neutral

Basic

10-14

10-13

10-11

10-9

10-7

10-5

10-3

10-1

100

[OH-]

examples
Examples:
  • Sample 18-1, p.434
  • Sample 18-2, p.435
  • Sample 18-3, p.436
  • Sample 18-4, p.438
  • Sample 18-5, p.439
formula s to know
Formula’s to Know

Ion Product Constant

Kw = [H3O+] [OH-] = 1 x 10-14

Acid Equations

pH = - log[H30+]

[H3O+] = antilog (-pH)

Base Equations

pOH = -log [OH-]

[OH-] = antilog (-pOH)

pH Scale Equation

This is from your formula sheet

pH + pOH = 14

Note H+ is the same as H3O+

measuring ph
Measuring pH
  • Why measure pH?
    • Everything from swimming pools, soil conditions for plants, medical diagnosis, soaps and shampoos, etc.
  • Sometimes we can use indicators, other times we might need a pH meter
acid base indicators
Acid-Base Indicators
  • An indicator is an acid or base that undergoes dissociation in a known pH range, and has different colors in solution (more later in chapter)
  • HIn H1+ + In1-
  • Examples: litmus, phenolphthalein, bromthymol blue: Fig 18.14, p.445
acid base indicators1
Acid-Base Indicators
  • Although useful, there are limitations to indicators:
    • usually given for a certain temperature

(25 oC), thus may change at different temperatures

    • what if the solution already has color?
    • ability of human eye to distinguish colors
acid base indicators2
Acid-Base Indicators
  • A pH meter(simulation)may give more definitive results
    • some are large, others portable
    • works by measuring the voltage between two electrodes
    • needs to be calibrated
    • Fig. 18.17, p.446
section 18 3 acid base theories
Section 18.3Acid-Base Theories
  • OBJECTIVES:
    • Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and Lewis
section 18 3 acid base theories1
Section 18.3Acid-Base Theories
  • OBJECTIVES:
    • Identify conjugate acid-base pairs in acid-base reactions.
svante arrhenius
Svante Arrhenius
  • Swedish chemist (1859-1927) - Nobel prize winner in chemistry (1903)
  • One of the first chemists to explain the chemical theory of the behavior of acids and bases
1 arrhenius definition
1. Arrhenius Definition
  • Acids produce hydrogen ions (H1+) in aqueous solution.
  • Bases produce hydroxide ions (OH1-) when dissolved in water.
  • Limited to aqueous solutions.
  • Only one kind of base (hydroxides)
  • NH3 (ammonia) would not be an Arrhenius base.

H2O

HCl  H+ + Cl-

H2O

NaOH  Na+ + OH-

polyprotic acids
Polyprotic Acids
  • Some compounds have more than 1 ionizable hydrogen.
  • HNO3 nitric acid – monoprotic
  • H2SO4 sulfuric acid - diprotic - 2 H+
  • H3PO4 phosphoric acid - triprotic - 3 H+
  • Having more than one ionizable hydrogen does not mean stronger!
polyprotic acids1
Polyprotic Acids
  • However, not all compounds that have hydrogen are acids
  • Also, not all the hydrogen in an acid may be released as ions

CH3COOH CH3COO- + H+

    • only those that have very polar bonds are ionizable - this is when the hydrogen is joined to a very electronegative element

NH4+ or H2O

arrhenius examples
Arrhenius examples...
  • Consider HCl
  • What about CH4(methane)?
  • CH3COOH (ethanoic acid, or acetic acid) - it has 4 hydrogens like methane does…?
  • Table 18.3, p. 442 for bases
what is a proton
What is a Proton
  • A proton is a positively charged subatomic particle (no neutrons, no electrons)
  • A hydrogen atom is 1 proton(+) and 1 electron (-)
  • A hydrogen ion or proton (H+) is a hydrogen atom that has lost its electron leaving only the proton
  • H  H+(proton) + 1 electron
2 br nsted lowry definitions
2. Brønsted-Lowry Definitions
  • Broader definition than Arrhenius
  • Acid is hydrogen-ion donor (H+ or proton); base is hydrogen-ion acceptor.
  • Acids and bases always come in pairs
  • HCl is an acid.
    • When it dissolves in water, it gives it’s proton to water.
  • HCl(g) + H2O(l) H3O+ + Cl-
  • Water is a base; makes hydronium ion.
acids and bases come in pairs
Acids and bases come in pairs...
  • A conjugate base is the remainder of the original acid, after it donates it’s hydrogen ion
  • A conjugate acid is the particle formed when the original base gains a hydrogen ion
  • Indicators are weak acids or bases that have a different color from their original acid and base
acids and bases come in pairs1
Acids and bases come in pairs...
  • General equation is:
  • HA(aq) + H2O(l) H3O+(aq) + A-(aq)
  • Acid Base Conjugate acid Conjugate base
  • NH3 + H2O NH41+ + OH1-

base acid c.a. c.b.

slide41

HCl + H2O H3O1+ + Cl1-

acid base c.a. c.b.

Amphoteric - acts as acid or base

(H20, HSO4-, H2PO4-) must be able to give a H+ or accept an H+

3 lewis acids and bases
3. Lewis Acids and Bases
  • Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction
  • Lewis Acid - electron deficient
  • Lewis Base - electron rich
  • Most general of all 3 definitions; acids don’t even need hydrogen!
  • Sample Problem 18-6, p.447
lewis acids bases
Lewis Acids/Bases
  • Several categories of substances can be considered Lewis acids: Electron Poor
  • 1) positive ions2) having less than a full octet in the valence shell3) polar double bonds (one end)4) expandable valence shells
  • Several categories of substances can be considered Lewis bases: Electron Rich
  • 1) negative ions2) one of more unshared pairs in the valence shell3) polar double bonds (the other end)4) the presence of a double bond
section 18 4 strengths of acids and bases
Section 18.4Strengths of Acids and Bases
  • OBJECTIVES:
    • Define strong acids and weak acids.
section 18 4 strengths of acids and bases1
Section 18.4Strengths of Acids and Bases
  • OBJECTIVES:
    • Calculate an acid dissociation constant (Ka) from concentration and pH measurements.
section 18 4 strengths of acids and bases2
Section 18.4Strengths of Acids and Bases
  • OBJECTIVES:
    • Arrange acids by strength according to their acid dissociation constants (Ka).
section 18 4 strengths of acids and bases3
Section 18.4Strengths of Acids and Bases
  • OBJECTIVES:
    • Arrange bases by strength according to their base dissociation constants (Kb).
strength
Strength
  • Strong acids and bases are strong electrolytes
    • They fall apart (ionize) completely.
    • Weak acids don’t completely ionize.
  • Strength different from concentration
  • Strong-forms many ions when dissolved
  • Mg(OH)2 is a strong base- it falls completely apart when dissolved.
    • But, not much dissolves- not concentrated
measuring strength
Measuring strength
  • Ionization is reversible.
  • HA H+ + A-
  • This makes an equilibrium
  • Acid dissociation constant = Ka
slide50

HA H+ + A-

Ka = [H+ ][A- ]

[HA]

(water is constant so we can ignore it)

Stronger acid = more products (ions), thus a larger Ka

(Table 18.4, p.448) (Table 18.5 p. 450)

what about bases
What about bases?
  • Strong bases dissociate completely.
  • B + H2O BH+ + OH-
  • Base dissociation constant = Kb
slide52

B + H2O BH+ + OH-

Kb = [BH+][OH-] [B]

(we ignore the water)

Stronger base = more dissociated, thus a larger Kb.

strength vs concentration
Strength vs. Concentration
  • The words concentrated and dilute tell how much of an acid or base is dissolved in solution - refers to the number of moles of acid or base in a given volume
  • The words strong and weak refer to the extent of ionization of an acid or base
  • Is concentrated weak acid possible?
strong acids
Strong Acids
  • HClO4
  • HI
  • HBr
  • HCl
  • HNO3
  • H2SO4
  • HA  H+ + A-
  • 100% Ionization occurs
strong bases
Strong Bases
  • NaOH
  • Mg(OH)2
  • KOH
  • Ca(OH)2
  • BOH B+ + OH-
  • 100% Ionization occurs
concentrated acid or base
Concentrated Acid or Base
  • Based on molarity
  • mol/L
  • Dilute solution = 0.5 mol/L
  • Concentrated solution = 18 mol/L
practice
Practice
  • Write the expression for HNO2
  • Write the Kb for NH3
  • Sample 18.7, p. 452
  • Carefully study Key Terms and equations, p. 453
  • Be sure to take all the self-tests that are available!
18 9 calculating dissociation constants
18.9 Calculating Dissociation Constants

A 0.100 M solution of acetic acid is only partially

ionized.

Using a measure of pH, the [H+] is calculated

as 1.34 x 10 -3 M.

What is the acid dissociation constant of acetic acid?

solution
solution

CH3COOH <-> CH3COO - + H +

Given – Initial concentration of acetic acid is 0.100 M

- Equilibrium concentration of H + is 1.34 x 10 -3 M

Solution – I need to use the…..

ICE TABLE

slide60

CH3COOH <-> CH3COO - + H +

0.100

0

0

-1.34 x 10 -3

+1.34 x 10 -3

+1.34 x 10 -3

0.0987

1.34 x 10 -3

1.34 x 10 -3

slide61

Products

Ka =

Reactants

[CH3COO -]1

[H+]1

Ka =

[CH3COOH]1

[1.34 x 10 -3]

[1.34 x 10 -3]

Ka =

[0.0987]

1CH3COOH <-> 1CH3COO - + 1H +

Ka = 1.82 x 10 – 5 which indicates that acetic acid is a weak acid( less than 1)