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Acid-Base Equilibria. REVIEW. Electrolyte : Substances that dissolves in water to produce solutions that conduct electricity Nonelectrolytes : Substances whose aqueous solutions do not conduct electricity Strong and weak relates to the degree of dissociation or ionization

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slide2

REVIEW

  • Electrolyte: Substances that dissolves in water
  • to produce solutions that conduct electricity
  • Nonelectrolytes: Substances whose aqueous
  • solutions do not conduct electricity
  • Strongandweakrelates to the degree of
  • dissociation or ionization
  • In aqueous solutions protons occasionally exist
  • as hydrated molecules, H(H2O)n+ (n=1)
slide3

Definitions:

Arrhenius:

An acid is a substance that increases the H+ (or H3O+) concentration in an aqueous solution.

HCl + H 2O  H3O+ + Cl-

HCl  H+ + Cl-

A base is a substance that increases the OH- concentration

in an aqueous solution.

NaOH(s)  Na+ + OH-

What about Na2CO3 ????

slide4

Bronsted-Lowry:

HCl(aq) + NaOH(aq) → HOH + NaCl

Acid = a proton donor in a Reaction

Base = a proton acceptor in a Reaction

slide5

Lewis:

An acid is an electron pair acceptor

H+

acid

..

H:O:H

..

A base is an electron pair donor

water

..

:O:H-

..

slide6

Acid/Base reactions:

Produce water and a salt (and sometimes carbon dioxide).

Hint: concentrate on the water first. Remember, water has the formula HOH.

Complete and balance the following:

HCl + KOH 

HOH + KCl

HCl + Ca(OH)2

2

2HOH + CaCl2

Require equal numbers

slide7

1. Ba(OH)2 + H3PO4

2. HC2H3O2 + NaOH

3. H2SO4 + KOH

4. H2CO3 + NaOH

5. Na2CO3 + HCl 

slide8

6. NH4OH + H2SO4

7. NH3 + HCl 

Give a definition of an acid:

An acid is a proton donor (H+)

Give a definition of a base:

A base is a proton acceptor

slide9

Conjugate acids and Conjugate bases

HCl + KOH 

HOH + KCl

acid

base

conj. base

conj. acid

Na2CO3 + 2HCl  H2CO3 + 2NaCl

acid

conj. base

base

conj. acid

Na2CO3 + 2HCl  H2O + CO2(g) + 2NaCl

acid

conj. acid

base

conj. base

slide12

What is a strong Acid?

An Acid that is 100% ionized in water.

Strong Acids:

100% ionized (completely dissociated) in water.

HCl + H2O  H3O+ + Cl-

often written as:

HCl  H+ + Cl-

slide13

Strong Acids:

100% ionized (completely dissociated) in water.

HCl + H2O  H3O+ + Cl-

Strong Acids:

Perchloric HClO4

Chloric, HClO3

Hydrobromic, HBr

Hydrochloric, HCl

Hydroiodic, HI

Nitric, HNO3

Sulfuric, H2SO4

slide14

What is a strong Base?

A base that is completely dissociated in water (highly soluble).

NaOH(s)  Na+ + OH-

Strong Bases:

Group 1A metal hydroxides

(LiOH, NaOH, KOH,

RbOH, CsOH)

Heavy Group 2A metal hydroxides

[Ca(OH)2, Sr(OH)2, and

Ba(OH)2]

slide16

Strong Acids:

100% ionized (completely dissociated) in water.

HCl + H2O  H3O+ + Cl-

Note the “one way arrow”.

Weak Acids:

Only a small % (dissociated) in water.

HC2H3O2 + H2O H3O+ + C2H3O2-

Note the “2-way” arrow.

Why are they different?

slide17

Strong Acids:

HCl HCl

HCl

HCl

HCl

(H2O)

ADD WATER to MOLECULAR ACID

slide18

Strong Acids:

Cl-

H3O+

(H2O)

Cl-

H3O+

H3O+

Cl-

Cl-

H3O+

H3O+

Cl-

Note: No HCl molecules remain in solution, all have been ionized in water.

slide19

Weak Acid Ionization:

HC2H3O2

HC2H3O2

(H2O)

HC2H3O2

HC2H3O2

HC2H3O2

Add water to MOLECULES of WEAK Acid

slide20

Weak Acid Ionization:

HC2H3O2

HC2H3O2

H30+ C2H3O2-

HC2H3O2

(H2O)

HC2H3O2

H30+ C2H3O2-

HC2H3O2

HC2H3O2

Note: At any given time only a small portion of the acid molecules are ionized and since reactions are running in BOTH directions the mixture composition stays the same.

This gives rise to an Equilbrium expression, Ka

slide21

ACID-BASE CONCEPTS

  • Acid-base reactions may be one of the most
  • important class of reactions
  • The most basic of the acid-base concepts is
  • the Arrhenius theory
  • Acids are substances that dissociate in water
  • to produce hydronium ions, H3O+ and bases are
  • substances that dissociate in water to produce
  • hydroxide ions, OH-
slide22

ACID-BASE STRENGTH

  • A strongacid is almost completely dissociated
  • or ionized

HCl(aq)  H+(aq) + Cl-(aq)

  • A weak acid is only partially dissociated or
  • ionized
  • So each reaction is an equilibrium with a K
slide23

STRONG ACIDS AND BASES

  • Common strong acids are either monoprotic or
  • diprotic
  • HA(aq) + H2O(l)  H3O+(aq) + A-(aq) 100%
  • Means [HA] = [H3O+]
  • Similar situation for strong bases
  • Typical strong bases: Group 1A metal and Ca,
  • Sr, and Ba hydroxides
  • Strong acids: HX, HNO3, HClO3, HClO4, H2SO4
slide24

WEAK ACIDS

  • Weak acids and bases are weak electrolytes
  • For a weak acid, HA
  • Ka: thedissociationorionizationconstant
  • For acetic acid
slide25

Ka = [H3O+][CH3COO-]/[CH3COOH]

  • Acids with larger ionization constants ionize or
  • dissociate to a greater extent than acids with
  • smaller ionization constants
  • The larger the value of Ka, the higher [H3O+]
  • and the stronger is the acid
  • HIO3; Ka = 1.6 x 10-1
  • CH3COOH; Ka = 1.8 x 10-5
  • HCN; Ka = 6.2 x 10-10
slide26

AUTOIONIZATION OF WATER

  • Reaction called autoionization of water
  • K = [H3O+][OH-]/[H2O]2
  • K[H2O]2 = [H3O+][OH-]
  • Kw = [H3O+][OH-]
  • Kw is called ion product of water
  • At 25 °C, Kw = 1.0 x 10-14
  • Valid also for dilute aqueous solutions
slide27

THE pH SCALE

  • The hydronium ion concentration is a measure
  • of a solution’s acidity
  • Usually small numbers
  • The pH scale is used express acidity and
  • basicity
  • pH = -log[H3O+]so[H3O+] = 10-pH
  • Note that as pH increases [H3O+] decreases
  • Value of Kw varies with temperature
slide28

“Acidity” depends on the concentration of

  • H3O+ ions
  • Acidic: [H3O+] > [OH-]
  • Basic: [H3O+] < [OH-]
  • Neutral: [H3O+] = [OH-]
  • Notice that neutral does NOT necessarily
  • mean pH 7
  • pH is usually quoted with the same number of
  • significant digits as the concentration
slide29

pH can be measured using a pH meter or by an

  • acid-base indicator
  • An indicator is usually an organic acid that have
  • different colors in solutions of different pH
  • Indicators exist that cover the entire pH scale
  • e.g. bromothymol blue: 6.0 – 7.6
  • phenolphthalein: 8.0 – 10
slide30

The p-scale can also be applied to ionization

  • constants
  • pKa = - log Ka
  • Thelargerthe value of Ka thesmallerthe value of pKa and thestrongerthe acid
  • From the acid-dissociation constant we can
  • calculate equilibrium concentrations as well as
  • pH
slide31

BRØNSTED-LOWRY THEORY

  • An acid is a proton donor
  • A base is a proton acceptor
  • An acid-base reaction is the transfer of a
  • proton from an acid to a base
  • Reactions can be described in terms of what
  • are called conjugate acid-base pairs
  • - species which differ by a proton
  • - a charge
slide32

- B is a proton acceptor; it is a base

- BH+ is a proton donor; it is an acid

So A- is the conjugate base of HA and

BH+ is the conjugate acid of B

  • HF/F- and H3O+/H2O are conjugate pairs
slide33

Relationship of Ka and concentration

  • In dilute solutions of weak acids, the assumption
  • - is that all the H3O+ is coming from the acid
  • - the concentration change of a species is small
  • compared to the initial concentration of that
  • species
  • - at equilibrium, [HA]  [HA]init
  • Always be cautious when using the assumptions
  • given above
slide34

%-dissociation

  • Amount of acid that dissociates can be
  • expressed as a percent
  • In general %-dissociation increases with the
  • value of Ka
  • For any weak acid, HA, %-dissociation
  • increases with dilution
slide35

POLYPROTIC ACIDS

  • Polyprotic acids:Acids that provide more than
  • one hydronium ion in solution
  • e.g. H2SO4, H3PO4
  • Polyprotic acids ionize in a stepwise manner
  • Consider sulfuric acid, H2SO4

(1)

(2)

  • In general, Ka1 > Ka2
slide36

WEAK BASES

  • Weak bases undergo equilibria in water
  • For a general weak base, B
  • Kb is called the base-dissociation constant
slide37

For conjugate acid-base pairs the product of

  • their equilibrium constants is the ionization
  • constant for water
  • Ka x Kb = Kw
slide38

FACTORS AFFECTING ACID STRENGTH

  • The strength of an acid depends on the polarity
  • of the H-E bond
  • The polarity of the bond is related to the bond
  • strength of H-E
  • The weaker the H-E bond the stronger the
  • acid
  • Take the hydrohalic acids:
  • HX; X = F, Cl, Br, I
slide39

HF << HCl < HBr < HI

  • For binary acid in the same group H-A bond
  • strength determines acid strength
  • Same applies to other groups:
  • H2O < H2S < H2Se
  • For binary acids in the same row polarity of the
  • H-E bond determines acid strength
  • :- acid strength increase with the
  • electronegativity of E
slide40

The oxoacids are also important

  • e.g. HNO3, H2SO4, HClO4
  • The acidity is dictated by factors that affect
  • the O-H bond strength
  • The electronegativity of the central element
  • The oxidation number of the central element
slide41

Some small, highly charged metal ions are quite

  • acidic
  • :- they are hydrated and transfers a proton
  • to a water molecule
slide42

SALTS: ACID-BASE PROPERTIES

  • Salts: an ionic compound that is formed when
  • an acid neutralizes a base
  • NaOH(aq) + HCl(aq)  NaCl(aq) + H2O
  • Aqueous solutions of salts can be neutral, acidic
  • or basic
  • Strong acid + strong base  neutral solutions
  • Strong acid + weak base  acidic solutions
  • Strong base + weak acid  basic solutions
slide43

When a salt dissolves in water, its constituent

  • ions may react with water – reaction called
  • hydrolysis
  • NEUTRAL SOLUTIONS
  • Salts of strong acids and strong bases
  • e.g. NaCl
  • Because the ions do not hydrolyze
slide44

Cl- is the conjugate base of HCl – it is a weak

  • base
  • The same argument is made for Na+
  • Essentially [H3O+]/[OH-] ratio does not
  • change
  • {Group 1A metals, Ca2+, Sr2+, Ba2+} and {I-,
  • Br-, Cl-, NO3-, ClO4-} does not hydrolyze
slide45

Acidic Solutions

  • NH3(aq) + HCl(aq)  NH4Cl(aq)
  • Salts from weak bases and strong acids
  • In aqueous solution NH4+ undergo hydrolysis
  • - the chloride ion does not
  • The generation of H3O+ from the reaction
  • makes these solutions acidic
slide46

Basic Solutions

  • Salts from strong bases and weak acids give
  • basic solutions
  • This basic anions of weak acids hydrolyze to
  • form hydroxide ions
  • NaCH3COO  CH3COO- + H3O+
slide47

SALTS FROM WEAK ACID AND BASES

  • e.g. NH4CH3COO, ammonium acetate
  • Aqueous solution of the salts may be basic,
  • acidic or neutral
  • The pH depends on the relative Ka and Kb of
  • the parent acid and base
  • Consider the case when Ka = Kb
  • NH4CH3COO(aq)  NH4+(aq) + CH3COO-(aq)
slide48

Both ions can undergo hydrolysis

  • Hydrolysis constant for the acetate ion (Kbh) is
  • equal to the hydrolysis constant for the
  • ammonium ion (Kah)
  • The same concentration of H3O+ as OH- is
  • produced  solution is NEUTRAL
slide49

If the parent Ka > Kb: solutions acidic

  • e.g. NH4F
  • Ka(HF): 7.2 x 10-4 > Kb (NH3): 1.8 x 10-5
  • Kb(F-): 1.4 x 10-11 < Ka(NH4+): 5.6 x 10-10
  • So NH4+ hydrolyzes to a greater extent than F-
  • more H3O+ is produced than OH-
slide50

If the parent Ka < Kb: solutions basic

  • e.g. NH4CN
  • Ka(HCN): 4.9 x 10-10 < Kb (NH3): 1.8 x 10-5
  • Kb(CN-): 2.0 x 10-10 > Ka(NH4+): 5.6 x 10-10
  • So NH4+ hydrolyzes to a greater extent than F-
  • more OH- is produced than H3O+
slide51

LEWIS THEORY

  • The Lewis acid-base theory is the most
  • complete that were are going to see
  • In this theory:
  • - an acid is any species that can accept an
  • electron pair
  • - a baseis any species that can donate an
  • electron pair
  • The Lewis acid typically have vacant orbitals
slide52

Common Lewis acids: metal ions and electron

  • deficient molecules such as AlCl3, BF3, etc.
  • A typical Lewis base would be ammonia, NH3
  • as well as the common anions
  • Acid-base reactions:
  • BF3 + NH3→ [F3B←NH3]
  • acid base acid-base adduct
  • AlCl3 + Cl-→[AlCl4]-