ARRHENIUS THEORY ACIDS and BASES

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					 Chapter 14
Acids & bases
         ARRHENIUS THEORY
          ACIDS and BASES
ACIDS - proton or H+ donor
BASE - hydroxide or OH- donor
NEUTRALIZATION - H+ + OH-  H2O
Problems:
     * H3O+ -- hydronium ion rather than H+
     * OH(H2O)3- -- not OH-
     * Other substances have acidic or basic
       properties also.
H+ surrounded by four H-bonded H2O molecules
OH-
surrounded by
three H-bonded
H2O molecules
        Bronsted-Lowery Theory
           Acids and Bases
Acid – any substance donating a proton, H+
Base – any substance accepting a proton
Conjugate Acid-Base Pairs:
     e.g. HF + NH3  NH4+ + F-
         acid 1   base 2         acid 2    base 1
AMPHOTERIC substances have both acidic and basic
properties.
Mono-, di-, tri-,………. to polyprotic acids.
Acidic versus nonacidic H atoms in compounds.
     For each of the following reactions, identify the acid, the base,
     the conjugate base, and the conjugate acid:
a. H2O + H2O  H3O+ + OH-
                O                                     O-
b.   CH3O- + CH3CCH3  CH3OH + CH3C                      CH2
c.   H2S + NH3  HS- + NH4+
d.   H2SO4 + H2O  H3O+ + HSO4-
e.   H+ + OH-  H2O
f.   H2PO4- + H2O  H3PO4 + OH-
g.   H2PO4- + H2PO4-  H3PO4 + HPO42-
h.   H2PO4- + H2O  HPO42- + H3O+
i.   Fe(H2O)63+ + H2O  Fe(H2O)5(OH) 2+ + H3O+
j.   HCN + CO32- CN- + HCO3-
k.   CO2 + 2H2O  HCO3- + H3O+
Graphical representations of strong and weak acid equilibria
Acid
strength
versus
conjugate
base
strength
  BRONSTED-LOWRY THEORY
          ACID and BASE STRENGTHS:
A proton transfers from a a stronger acid to a stronger
     base, from a weaker acid and weaker base.

        LEVELING EFFECT of SOLVENTS:
The strongest acid in a solvent is the conjugate acid of
the solvent. The strongest base is the conjugate base.
             Acid      H3O+ in water
             Base      OH- in water
       Relative strengths of some Bronsted-Lowry acids and
                       their conjugate bases
                 Acid                  Base
Strongest        HClO4                 ClO4-          Weakest
Acid             H2SO4                 HSO4-          bases
                 HI                    I-
                 HBr                   Br-
                 HCl                   Cl-
                 HNO3                  NO3-
                 H3O+                  H2O
                 HSO4-                 SO42-
                 H2SO3                 HSO3-
                 H3PO4                 H2PO4-
                 HNO2                  NO2-
                 HF                    F-
                 CH3CO2H               CH3CO2-
                 H2CO3                 HCO3-
                 H2S                   HS-
                 NH4+                  NH3
                 HCN                   CN-
                 HCO3-                 CO32-
                 HS-                   S2-
                 H2O                   OH-
Weakest          NH3                   NH2-           Strongest
Acid             OH-                   O2-            bases
     STRONG ACIDS, BASES
          IN WATER
Autoionization
H2O(l) + H2O(l)  H3O+(aq) + OH-(aq)
Kw = [H3O+] [OH-] = 1.0 x 10-14 at 25oC
At equilibrium
     [H3O+] = [OH-]
     so each = 1.0 x 10-7
     Kw changes with temperature
     but [H3O+] = [OH-]
pH, pOH, pM, pX SCALES

 pH = - log [H3O+]
 pOH = - log [OH-]
 pM = - log [M]
   etc.
   pH, pOH CALCULATIONS
      pH = - log [H3O+]
      pOH = - log [OH-]
 Kw = [H3O+] [OH-] = 1.0 x 10-14
 so p[H3O +] + p[OH-] = 14.00

   or    pH = 14.00 – pOH
CALULATE SOME pH and pOH VALUES
          [H+]    pH
          10-14   14    1 M NaOH
          10-13   13
Basic     10-12   12    Ammonia
          10-11   11   (Household
          10-10   10   Cleaner)
          10-9    9
          10-8    8     Blood
Neutral   10-7    7     Pure Water
          10-6    6     Milk
          10-5    5
          10-4    4
          10-3    3    Vinegar
                         Lemon juice
Acidic    10-2    2    Stomach acid
          10-1    1
          1       0     1 M HCl
Calculate the pH of each solution:
a.    [H+] = 1.4 x 10-3 M            e.   [OH-] = 8 x 10-11 M
b.    [H+] = 2.5 x 10-10 M           f.   [OH-] = 5.0 M
c,.   [H+] = 6.1 M                   g.   pOH = 10.5
d.    [OH-] = 3.5 x 10-2 M           h.   pOH = 2.3

Calculate [H+] and [OH-] for each solution:
a.    pH = 7.41 (the normal pH of blood)
b.    pH = 15.3
c.    pH = -1.0                      e. pOH = 5.0
d.    pH = 3.2                       f. pOH = 9.6

How many significant figures are there in the numbers: 10.78, 6.78,
0.78? If these were pH values, to how many significant figures can you
express the [H+]? Explain any discrepancies between your answers to
the two questions.
Values of Kw as a function of temperature are as follows:
   Temp (oC)                            Kw
       0                           1.14 x 10-15
      25                           1.00 x 10-14
      35                           2.09 X 10-14
      40                           2.92 x 10-14
      50                           5.47 x 10-14
a. Is the autoionization of water exothermic or
   endothermic?
b. What is the pH of pure water at 50oC?
c. Restate your answers to water at 50oC. Which of the
   three criteria for neutrality is most general?
d. From a plot of ln(Kw) versus 1/T (using the Kelvin
   scale), estimate Kw at 37oC, normal physiological
   temperature.
e. What is the pH of a neutral solution at 37oC?
-28
-29           Y = -9.2338 – 6870.6x R^2 = 0.999

-30
-31
-32

-33

-34
-35
0.0028   0.0030     0.0032     0.0034     0.0036
                     1/T
             pH MEASUREMENT
-- INDICATORS. colored weak acids and bases
-- pH Meters. Glass membrane with a voltage (potential)
              difference across the glass.


        pH and BODY CHEMISTRY
Normal pH 7.3 to 7.5
Acidosis pH < 7.3
Alkalosis pH > 7.45
Body Chemistry is “buffered” with carbonates and
phosphates which reduce the acid/alkalinity changes.
These reduce the pH effects.
         WEAK ACIDS
    IONIZATION CONSTANTS
   HA(aq) + H2O(l)  H3O+(aq) + A-(aq)
                             -
                 [H 3O ] [A ]
            Ka 
                     [HA]
Ka values at 25oC are known and tabulated for a
          large number of weak acids.
Graphic
representations
of strong and
weak acid
equilibria
 Values of Ka for Some Common Monoprotic Acids

Formula        Name                                Value of Ka*
HSO4-          Hydrogen sulfate ion                1.2 x 10-2
HClO2          Chlorous acid                       1.2 x 10-2




                                                                  Increasing acid strength
HC2H2ClO2      Monochloroacetic acid               1.35 x 10-3
HF             Hydrofluoric acid                   7.2 x 10-4
HNO2           Nitrous acid                        4.0 x 10-4
HC2H3O2        Acetic acid                         1.8 x 10-5
[Al(H2O)6]3+   Hydrated aluminum (III) ion         1.4 x10-5
HOCl           Hypochlorous acid                   3.5 x 10-8
HCN            Hydrocyanic acid                    6.2 x 10-10
NH 4+          Ammonium ion                        5.6 x 10-10
HOC6H5         Phenol                              1.6 x 10-10
*The units of Ka are mol/L, but are customarily omitted.
Write the dissociation reaction and the
corresponding equilibrium expression for each of
the following acids in water.
 a. H3PO4        f. Acetic acid, CH3CO2H (HC2H3O2)
 b. H2PO4-       g. Phenol, C6H5OH
 c. HPO42-       h. Benzoic acid, C6H5CO2H
 d. Ti(H2O)64+ i. Glycine, H2NCH2CO2H
 e. HCN
Write the reaction and the corresponding Kb
equilibrium expression for each of the following
substances acting as bases in water.
  a. PO43-         g. Glycine, NH2CH2CO2H
  b. HPO42-        h. Ethylamine, CH3CH2NH2
  c. H2PO4-        I. Aniline, C6H5NH2
  d. NH3           j. Dimethylamine, (CH3)2NH
  e. CN-
  f. Pyridine, C5H5N
  WEAK ACID CALCULATIONS
  HA(aq) + H2O(l)  H3O+(aq) + A-(aq)
      2H2O(l)  H3O+(aq) + OH-(aq)
                            
                        [H3O ]
         % Ionization          x 100%
                         [HA]
To simplify calculations, if % ionization is < 5%,
  then CHA  [HA] OR [HA] = CHA – [H3O+],
 Set up pH equilibria calculations in tables as in
         previous equilibria problems.
        Solving Weak Acid Equilibrium Problems
List the major species in the solution
Choose the species that can produce H+, and write balanced equations for
the reactions producing H+
Using the values of the equilibrium constants for the reactions you have
written, decide which equilibrium will dominate in producing H+
Write the equilibrium expression for the dominant equilibrium.
List the initial concentrations of the species participating in the dominant
equilibrium.
Define the change needed to achieve equilibrium; that is, define x..
Write the equilibrium concentrations in terms of x.
Substitute the equilibrium concentrations into the equilibrium expression.
Solve for x the “easy” way; that is , by assuming that [HA]0-x  [HA]0.
Use the 5% rule to verify whether the approximation is valid.
Calculate [H+] and pH.
For trichlorophenol (HC6H2Cl3O), Ka = 1 x 10-6,
Calculate the concentrations of all species and the
pH of a 0.05 M solution of trichlorophenol in water.

A solution is prepared by dissolving 0.56 g benzoic
acid (C6H5CO2H), Ka = 6.4 x 10-5) in enough water
to make 1.0 L of solution. Calculate [C6H5CO2H].
[C6H5CO2-], [H+], [OH-], and the pH in this
solution.

Calculate the pH of a solution containing a mixture
of 0.050 M HNO3 and 0.50M HC2H3O2.
         WEAK BASES
    IONIZATION CONSTANTS
   B(aq) + H2O(l)  BH+(aq) + OH-(aq)
                              _
                    [BH ] [OH ]
              Kb 
                         [B]
Kb values at 25oC are tabulated or may be calculated
                 from Kw and Ka
        Kw = KaKb so Kb = Kw/Ka
      Where Ka is the conjugate acid constant
Thallium (Tl) hydroxide is a strong base used in the
synthesis of some organic compounds. Calculate
the pH of a solution containing 2.48 g TlOH per
liter.



For the reaction of hydrazine (N2H4) in water.
   H2NNH2 + H2O  H2NNH3+ + OH-
Kb is 3.0 x 10-6. Calculate the concentrations of all
species and the pH of a 2.0 M solution of hydrazine
in water.
     Two Weak Acid in Solution


A solution of 0.100 M HClO, Ka =3.5 x 10-8
and 0.100 M Formic acid, HCO2H, Ka = 1.8
x 10-4 are mixed. What is the resulting
solution pH?
 Polyprotic Acid Equilibria

What is the pH of a solution
0.100 M sulfurous acid,
H2SO3, Ka1 = 1.5 x 10 -2 and

Ka2 = 1.0 x 10-7?
ACID-BASE PROPERTIES OF THE OXIDES
              (PART I)
ACID BASE PROPERTIES OF THE OXIDES
               (PART II)
ACID-BASE PROPERTIES OF THE OXIDES
             (PART III)
ACID BASE PROPERTIES OF THE OXIDES
             (PART IV)
ACID-BASE PROPERTIES OF THE OXIDES
                (PART V)
ACID BASE PROPERTIES OF THE OXIDES
             PART (VI)
ACID BASE PROPERTIES OF THE OXIDES
             PART (VII)
ACID-BASE PROPERTIES OF THE OXIDES
            PART (VIII)`
    BRONSTED-LOWRY
        THEORY

OXIDES, HYDROXIDES, ANHYDRIDES
Acid, base reactions
Dehydrations (formation of anhydrides)
Hydration of oxides
                ACID STRENGTHS
                 of WEAK ACIDS
Oxoacids:           Ka  up as the central atom
                    oxidation state  up.
                    Ka  up as central atom of same oxidation state
                          moves left to right in the periodic table.
                    Ka  up as the central atom of the same oxidation
                          state moves UP in the same Group or family.
Polyprotic acids:
                    Ka decreases by approximately 105 for each
                    successive H+ ionized.
Binary acids:       (only H and another element)
                    Within a period, Ka  up as electronegativity of
                    the other element.
                    Within a group, Ka  up going down the group
                    to higher mass and larger size.
SnO2 +   ? H2O  ?



?HCrO4-  ?Cr2O72- + ?



?HMnO4-  ? Mn (VI) compound + ?
HYDROLOYSIS OF IONIC SALTS


     The pH of each type salt in
 solution depends on the Ka or Kb of
        the hydrolyzing ion(s).
Salt Derived   Ions Undergoing
From:          Hydrolysis        pH            Examples


Strong base    Neither           Neutral       NaCl, KNO3,
 strong acid                     pH = 7        BaCl2, CaBr2

Strong base    Anion             Basic         LiCN, KNO2, CaF2
 weak acid                       pH > 7        NaCH3CO2

Weak base      Cation            Acidic,       NH4Cl, Al(NO3)3,
strong acid                      pH < 7        (CH3)3NHBr

Weak base,     Both              Acidic        NH4NO2
weak acid                        if Kb< Ka;
                                 Neutral       NH4CH3CO2
                                 if Kb = Ka;
                                 basic         NH4CN
                                 if Kb > Ka
Arrange the following 0.10 M solutions in order
from most acidic to most basic
 KOH, KBr, KCN, NH4Br, NH4CN, HCN



Given that the Ka value for acetic acid is 1.8 x 10-5
and the Ka value for hypochlorous acid is 3 x 10-8,
which is the stronger base, OCl- or C2H3O2-
              Acid-Base Equilibria
What is the pH of a solution of 0.150 M sodium
nitrite, NaNO2?
HNO2, Ka = 4.0 x 10-4.



What is the pH of a solution of 0.150 M
hydrazinnium chloride, H2NNH3+?
H2NNH3+, Kb = 3.0 x 10-6
Calculate the pH of each of the following solutions.
   a. 0.10 M CH3NH3Cl
   b. 0.050 M NaCN
   c. 0.20 M Na2CO3 (consider only the reaction )
       CO32- + H2O  HCO3- + OH-


Sodium azid (NaN3) is sometimes added to water to
kill bacteria. Calculate the concentration of all
species in a 0.010 M solution of NaN3. The Ka
value for hydrazoic acid (HN3) is 1.9 x 10-5.
        LEWIS THEORY
       ACIDS and BASES
     Lewis base - an electron pair donor
    Lewis acid - an electron pair acceptor

  Lewis acid + Lewis base  Adduct
                   (coordination compound)

e.g. Cu2+(aq) + 4CN-(aq)  Cu(CN)42-(aq)
  Look at Lewis Dot structures for lone pairs.

				
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