Chapter 16 : Acid-Base Equilibira

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Chapter 16 Acid-Base Equilibira

• Lauren Querido

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Table of Contents

• 16.1 Review
• 16.2 Brønsted-Lowry Acids and Bases
• 16.3 Autoionization of Water
• 16.4 pH Scale
• 16.5 Strong Acids and Bases
• 16.6 Weak Acids
• 16.7 Weak Bases

• 16.8 Relationship Between Ka and Kb
• 16.9 Acid-Base Properties of Salt Solutions
• 16.10 Acid-Base Behavior and Chemical
  Structure
• 16.11 Lewis Acids and
• Bases

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16.1 Review

• Acids
• Sour in taste
• Litmus paper turns red
• Bases
• Bitter, slippery
• Litmus paper turns blue
• When acids and bases mix, their properties
  disappear!

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Arrhenius Acids and Bases

• Svante Arrhenius (1880)
• In aqueous solutions
• Acids will increase the concentration of H ions
  when dissolved in water.
• Bases will increase the concentration of OH-ions
  when dissolved in water.

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16.2 Brønsted-Lowry Acids and Bases

• 1923 Brønsted and Lowry made a more general
  definition
• Brønsted-Lowry Acid is a substance that can
  transfer a proton. It must have a hydrogen atom
  that can be lost as H.
• Brønsted-Lowry Base is asubstance that can accept
  a proton. Must have a nonbonding pair of
  electrons to gain a H ion.

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Conjugate Acid-Base Pairs

• Conjugate base- Removal of proton from the acid
• Conjugate acid- Addition of proton to the base

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Relative Strengths of Acids and Bases

• The stronger the acid, the weaker its conjugate
  base.
• The stronger the base, the weaker its conjugate
  acid.
• 1. Strong acids completely transfer protons to
  water.
• 2. Weak acids partly dissociate in aqueous
  solutions and exist as a mixture of acid
  molecules and component ions.
• 3. Negligible acidity contain Hydrogen but do
  not demonstrate acidic behavior. Ex CH4
• Position of equilibrium favors transfer of proton
  from stronger acid to stronger base.

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16.3 Autoionization of Water

• Ion product of water
• 1.0 x 10 14 H OH-
• This is used to calculate concentrations of H
  and OH- .
• If H OH-, than neutral equation
• If H gt OH-, than acidic equation
• If H lt OH-, than basic equation

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16.4 The pH Scale

• pH -log H
• pH of 7 is neutral
• Acidic solution 0 lt pH lt 7
• Basic solution 14 gt pH gt 7
• Other p scales are
• pOH -log OH-
• pOH pH -log Kw 14.0

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Examples on the pH Scale
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Measuring pH

• A pH meter consists of a pair of electrodes
  connected to a meter which pH is generated when
  placed in the solution.
• An acid-base indicator turns a color if placed
  in acid or base. Ex litmus paper

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16.5 Strong Acids and Bases

• Strong Acids
• 7 most common strong acids are
• HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4
• In acidic reactions, equilibrium lies entirely to
  the right side.
• Completely dissociates
• Example
• HNO3 gt H NO3-

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Strong Bases

• Most common strong bases are ionic hydroxides of
  alkali metals (1A) and heavier alkaline earth
  metals (2A). Examples LiOH, RbOH, CsOH, NaOH,
  KOH, and Ca(OH)2, Sr(OH)2, and Ba(OH)2.
• Other strong bases react with water to form OH-
  such as Na2O, CaO.
• Also, anions O2-, H-, and N3- are stronger bases
  than OH- and therefore remove a proton from H2O.
• Example N3- H2O gt NH3 3OH-

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16.6 Weak Acids

• A weak acid only partially ionizes in aqueous
  solutions.
• General weak acid equation
• HX ? H X- where H is Hydrogen
• Many weak acids contain some Hydrogen atoms
  bonded to carbon atoms and oxygen atoms (organic
  compounds).
• Ka is the acid dissociation constant.
• The larger the value of Ka , the stronger the
  acid.

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Calculating Ka from pH

• Use and ICE box!
• Sample exercise
• A student prepared a .10 M solution of formic
  acid and measures its pH which was 2.38.
• A) calculate Ka for formic acid
• B) what percentage of the acid is ionized in the
  .10M solution?

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Answer

• a) HCHO2 ? H CHO2-
• Ka HCHO2- HCHO2
• pH -logH
• 10 2.38 4.2 X 10-3M

• Ka 4.2 X 103 4.2 X 103 .10
• 1.8 X 10-4 4.2 X 103 4.2 X 103
  .10
• b) Percent Ionization
• Concentration of H
• Initial concentration of component
• 4.2

HCHO2 ? H CHO-
I .10 M 0 M 0 M
C -4.2 X 103 4.2 X 103 4.2 X 103
E .10 - 4.2 X 103 4.2 X 103 4.2 X 103
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Using Ka to Calculate pH

• The best way to explain this is by an example.
• Calculate the pH of a .30 M solution of acetic
  acid at 25o C. (Ka 1.8 X 10-5)
• So HC2H3O2 ? H C2H3O2-
• Ka HC2H3O2- 1.8 X 10-5
• HC2H3O2
• What now?

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HC2H3O2 ? H C2H3O2-
I .30 M 0 M 0 M
C -x x x
E .30-x x x

Ka (x)(x) 1.8 X 10-5 (.30
x) Either do the quadratic equation or in this
case you can take out x in the denominator.
H x 2.3 X 10-3 pH -log 2.3 X 10-3
2.64
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Polyprotic Acids

• Polyprotic acids have more than one ionizable
  Hydrogen atom.
• Example
• H2SO3 ? H HSO4-
• HSO4- ? H SO32-
• The second Ka (Ka2) is much smaller than Ka1
  because it is easier to remove the first proton.

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16.7 Weak Bases

• Weak base water gt conjugate acid hydroxide
  ion
• Kb is the base-dissociation constant (equilibrium
  in which base reacts when H2O to form conjugate
  acid and OH- ion).
• Types of weak bases
• Neutral substances that have atoms with a
  non-bonding pair of electrons that can serve as a
  proton acceptor.
• Most of these contain amines, N-H which is
  sometimes replaced with a bond between C or N Ex
  NH2CH3
• Anions of weak acids
• Ex ClO- H2O ? HClO H
• ClO- is the weak base

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16.8 Relationship Between Ka and Kb

• Reaction 1 reaction 2 reaction 3
• Which leads to K1 x K2 K3
• Which leads to Ka x Kb Kw
• Kw is the ion-product constant for water
• Kw 1 x 10-14
• As the strength of the acid increases, the
  strength of the base decreases and visa-versa.
• pKa pKb pKw 14.00

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16.9 Acid-Base Properties of Salt Solutions

• Hydrolysis is the process at which ions react
  with water and produce H or OH-
• X- H2O ?HX OH-
• Anions of strong acids do not influence pH
• Ex NO3-
• Anions that still have ionizable protons are
  amphoteric
• Ex HSO3- from H2SO4
• Most cations (except 1A elements and Ca2, Sr2.
  Ba2) act as weak acids in solution.

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Predicting the pH of a Solution

• 1. Salts derived from a strong acid and a strong
  base makes a neutral pH (pH of 7).
• NaOH HCl gt NaCl H2O
• 2. Salts derived from a strong base and a weak
  acid will yield a pH of above 7 because the anion
  hydrolyzes to produce OH- ions and the cation
  does not hydrolyze.
• NaOH HClO gt NaClO H2
• 3. Salts derived from a weak base and a strong
  acid will result in a pH that is below 7 because
  the cation hydrolyzes to produce H ions and the
  anion does not hydrolyze.
• Al(OH)3 3HNO3 gt Al(NO3)3 3H2O

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• 4. Salts derived from a weak base and a weak acid
  will yield a pH that is dependant on the constant
  value of the constant dissociations (Ka and Kb).
• if the base is more basic than the acid is
  acidic, then the solution will have a pH that is
  greater than 7.
• If the acid is more acidic, than the pH will be
  less than 7.
• NH4 CN- ? NH4CN
• NH4 Ka 5.6 X 10-10
• CN- Kb 2.0 X 10-5
• Therefore, the pH of NH4CN is greater than 7

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16.10 Acid-Base Behavior and Chemical Structure

• Factors that effect acid strength
• If H-X bond is polarized (X is more
  electronegative) the H acts as a proton acceptor.
• Non-polar bonds (CH4) produce neutral solutions.
• Weaker bonds dissociate more easily than very
  strong bonds.
• HF is a weak acid because of this.
• The greater the stability of the conjugate base,
  the weaker the acid.
• Ultimately, there are three factors effecting
  acid strength
• Polarity of H-X bond
• Strength of H-X bond
• Stability of conjugate base, X-

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Binary Acids

• Binary acids are composed of Hydrogen and a
  non-metal.
• Ex HCl, HF, H2S, etc.
• The more polar the bond,the stronger it is
• The weaker the bond, the stronger the acid.
• Strength of the bond decreases (acidity
  increases) as the element increases in size or
  moves down a group.
• Acid strength increases (acidity decreases)
  moving from left to right

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Group Group Group Group
4A 5A 6A 7A
Period 2 CH4 No acid or base properties NH3 Weak base H2O ------- HF Weak acid
Period 3 SiH4 No acid or base properties PH3 Weak base H2S Weak acid HCl Strong acid

Increasing base strength
Increasing acid strength
Increasing acid strength
Increasing base strength
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Oxyacids

• Oxyacids are acids with an OH group is bound to a
  central atom.
• Example H2SO4

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OH- Bonding

• To determine if an OH group acts as an acid or
  base, consider this
• If Y is a metal than sources of OH- behave as
  bases.
• If Y is a non-metal than the compound will not
  readily lose the OH- ion.
• The electronegativity will increase and so will
  the acidity.
• The increasing number of Oxygen atoms stabilizes
  the conjugate base and thus increases the
  strength of the acid.

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Oxyacid Rules of Thumb

• Oxyacids that have the same number of OH groups
  and the same number of Oxygen atoms, acid
  strength increases with increasing
  elecronegativity of the central atom
• Example HClO gt HBrO gt HIO (gt more acidic)
• 2. For oxyacids with the same central atom,
  acid strength increases with increasing number of
  Oxygen atoms that are attached.
• Example HClO lt HClO2 lt HClO3 lt HClO4 ( lt less
  acidic)

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Carboxylic Acid

• Carboxylic acids are organic compounds.
• -COOH is the
  functional group
• -R is either a
  Hydrogen
• or Carbon based group
• If an extra Oxygen is added than it stabilizes
  the conjugate base and increases the acidity.
• If conjugate base has resonance structures, it
  spreads the negative charge evenly over the
  compound.
• Acid strength of carboxylic acid increases as the
  number of electronegative atoms increase.

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6.11 Lewis Acids and Bases

• G.N. Lewis proposed this
• Lewis Acids have an incomplete
• octet of electrons. Function as
• electron pair acceptors
• Lewis Bases act as electron pair
• donators

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Hydrolysis of Metal Ions

• Hydration is a process when when metals attract
  unshared electron pairs of water molecules.
• The metal acts as Lewis acid
• The water acts as Lewis base
• Ex Fe(H2O)63 ? Fe(H2O)5(OH)2 H
• So, general equation
• M(H2O)nc ? M(H2O)n-1(OH)c-1 H

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