Arhenius Definition

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Arhenius Definition

• Acids produce hydrogen ions in aqueous solution.
• Bases produce hydroxide ions when dissolved in
  water.
• Theory limited to aqueous solutions.
• Only one kind of base exists (ex. NaOH).
• NH3 ammonia could not be an Arhenius base.

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Brönsted-Lowry Definitions

• And acid is an 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 makes hydronium ion H3O

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Acid/Base Conjugate Pairs

• General equation
• HA(aq) H2O(l) H3O(aq) A-(aq)
• Acid Base Conjugate acid Conjugate
  base
• This is an equilibrium.
• Competition for H between H2O and A-
• The stronger base controls direction.
• If H2O is a stronger base it takes the H
• Equilibrium moves to right.

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Conjugate Acid/Base Pairs (new)

• Conjugate acid-base pair aacid and its
  conjugate base (HA/A- HSO4-/SO4-2)) and base
  and its conjugate acid ( A-/HA H2O/H3O1)
• Every Brönsted acid has its conjugate base
• Every Brönsted base has its conjugate acid

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The Acid/Base Properties of Water

• Water behaves as both an acid and a base.
• Water auto ionizes
• 2H2O(l) H3O(aq) OH-(aq)
• KW H3OOH-HOH-
• At 25ºC KW 1.0 x10-14
• In EVERY aqueous solution.
• Neutral solution H OH- 1.0 x10-7
• Acidic solution H gt OH-
• Basic solution H lt OH-

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Equilibrium Constant for Water

• Kw called ion-product constant product of ion
  concentrations of H and OH- at a a particular
  temperature
• At 25 ºC Kw 1.0 x 10-14
• H3O1 OH- 1.0 x 10-7
• Neutral solution H3O1 OH-
• Acidic solution H3O1 gt OH-
• Basic solution H3O1 lt OH-

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Properties of Water-Example

• Example
• The concentration of OH-1 in a certain household
  ammonia cleaning solution is 0.0025 M.
  Calculate the concentration of H1 ions.
• Answer 4.0 x 10-12 M

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Amphoteric/Amphiprotic Substances

• Substances behave as both acid and a base
• Examples
• HSO4-1(aq) H2O ? SO4-2 H3O1
• HSO4-1(aq) H2O ? H2SO4 OH-1
• HCO3-1 H2O ?
• H2PO4-1 H2O ?
• HS-1 H2O ?

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pH Measure of Acidity

• H3O1 10-pH pH -logH3O
• Used because H is usually very small
• As pH decreases, H increases exponentially
• Sig figs only the digits after the decimal place
  of a pH are significant
• H 1.0 x 10-8 pH 8.00 2 sig figs
• pOH -logOH-
• pKa -log K

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Relationships

• 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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Calculating pH of Solutions

• Always write down the major ions in solution.
• Remember these are equilibria.
• Remember the chemistry.
• Dont try to memorize there is no one way to do
  this.

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pH Problem Sets (1)

• Example 1 determine the pH of 0.00015 M HCl
  solution.
• Answer 3.82
• Example 2 What is the H3O in a solution with
  pH 2.19
• Ans 6.5 x 10-3

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pH- Problem Set (2)

• In NaOH solution OH- is 2.9 x 10-4 M.
  Calculate the pH of the solution
• Answer 10.46

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The Strength of Acids and Bases

• Determined by the degree of ionization of the
  substance (HCl versus CH3COOH)
• The stronger the acid, the weaker its conjugate
  base
• Reaction favored towards the weaker species (weak
  acid or weak base)
• H3O the strongest species in aqueous solution.
• Weaker acids react with water to produce H3O1,
  but the equilibrium is to the left (reactants)

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Strong Acids

• Acids stronger than H3O1
• 100 dissociation in aqueous solution
• Reaction goes to completion
• HCl H2O ? Cl- H3O

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Weak Acids

• Acids weaker than H3O1 react with water in a
  reversible reaction
• Product H3O1 and conjugate base
• HF(aq) H2O(l) ? H3O1(aq) F-(aq)
• Equilibrium primarily goes towards reactants
  (left)

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Bases

• The OH- is the strongest base that can exist in
  H2O solution
• Strong bases dissociate completely (100
  dissociation) NaOH
• NH2-1(aq) H2O(l) ? NH3(aq) OH-(aq)
• Weak bases react with water in a reversible
  reaction
• NH3(aq) H2O(l) ? NH41 OH-

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Problem Set (1)

• Calculate the pH of
• A 1.0 x10-3 M HCl solution (strong acid)
• A 0.020 M Ba(OH)2 solution (strong base)
• Answers a) 3.00
• b) 12.60

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Problem Set (2)

• Predict the direction of the following reaction
  in aqueous solution. Explain
• Note Use Table 15.2, page 611
• HNO2(aq) CN-(aq) ? HCN(aq) NO2-1(aq)
• Answer from left to right

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Types of Acids

• Polyprotic Acids- more than 1 acidic hydrogen
  (diprotic, triprotic).
• Oxyacids - Proton is attached to the oxygen of an
  ion.
• Organic acids contain the Carboxyl group -COOH
  with the H attached to O
• Generally very weak.

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Strong Acids

• HBr, HI, HCl, HNO3, H2SO4, HClO4
• ALWAYS WRITE THE MAJOR SPECIES ( H and Cl-)
• Completely dissociated
• H HA
• OH- is going to be small because of equilibrium
• 10-14 HOH-
• If HAlt 10-7 water contributes H

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Binary Acids and Strength

• Depends on
• Bond dissociation energy (tendency for bond
  breakage)
• Radius of the anion (the larger the anion, the
  smaller the tendency to attract the H)

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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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Binary Acids and Strength
Acid HF HCl HBr HI
Bond Dissoc. kJ/mol 569 431 368 297
Ion Radius, pm 136 181 195 216
Ka 6.6 x 10 -4 106 108 109
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Binary Acids and Electronegativity
?Electronegativity 0.4 .09 1.4 1.9
Acid strength CH4lt NH3lt H2Olt HF
Ka 1.00 x 10 -14 6.6 x 10-4
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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 last chapter.

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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 acid

• H2CO3 H HCO3- Ka1 4.3 x 10-7
• HCO3- 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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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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Acid Base Reactions Review

• Strong acid strong base neutralization
  reaction
• Weak acid-strong base equilibrium
• HF(aq) H2O(aq) ? H3O1(aq) F-(aq)
• Strong acid- weak base nitric acid ammonia
  (equimolar quantities)
• H(aq) NH3 (aq) ? NH41(aq)
• NH41(aq) H2O(l) ? NH3(aq) H3O1(aq)
• Solution I ACIDIC

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Acid Base Reactions Review

• Weak acid- weak base
• NH3 (aq) CH3COOH(aq) ?CHCOO-(aq)
  NH41(aq)
• CH3COO-(aq) H2O(l) ? CH3COOH OH-(aq)
• NH41(aq) H2O(l) ? NH3(aq) H3O1(aq)
• The pH of the solution depends on the relative
  strengths of the acid and base calculate next
  chapter

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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.
• Form hydroxides

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)
• Amphoteric oxides
• Al(OH)3(s) 3H(aq) ?Al3 (aq) 3H2O(l)
• Al(OH)3(s) OH-(aq) ? Al(OH)4-1(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
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• The End

• The End

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Example
Example

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