Acids and Bases

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

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Objectives SWBAT Distinguish between Arrhenius,
Bronsted Lowry and Lewis acids and bases.
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Arrhenius Definition

• Definitions
• Acids produce hydrogen ions in aqueous solution.
• Bases produce hydroxide ions when dissolved in
  water.
• Limited to aqueous solutions.
• Only one kind of base (hydroxide).
• NH3 ammonia could not be an Arrhenius base.

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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

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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.

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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. ?

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Remember

• Hydroxides (and some oxides) are strong bases.
• All other common bases are weak.
• Weak bases establish an equilibrium system like
  acids.

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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.

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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.

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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

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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.

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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

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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.

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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-

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• 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-

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pH

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

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pH

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

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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

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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.

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Basic
Acidic
Neutral
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Strong Acids

• HBr, HI, HCl, HNO3, H2SO4, HClO4
• These acids completely dissociate
• Therefore, H HA
• 10-14 HOH-

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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.

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Example

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

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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)

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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

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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?

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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.

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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

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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
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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

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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.

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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.

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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

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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

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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

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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

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Basic Salts

• Ka x Kb HFOH- x HF- F-
  HF

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Basic Salts

• Ka x Kb HFOH- x HF- F-
  HF
• Ka x Kb OH- H
• Ka x Kb KW

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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

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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.

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Anion of weak acid, cation of weak base

• Ka gt Kb acidic
• Ka lt Kb basic
• Ka Kb Neutral

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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.

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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

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Strength of Oxyacids
Electron Density
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Strength of Oxyacids
Electron Density
O
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Strength of Oxyacids
Electron Density
O
O
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Strength of Oxyacids
Electron Density
O
O
O
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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
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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)

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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
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Lewis Acids and Bases

• Boron triflouride wants more electrons.

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