1. The Arrhenius Definition
a. Acid=a substance which produces H+ in water solution.
b. Base=a substance which produces OH- in water solution.
2. The Bronsted-Lowry Definition
a. Acid=a substance which can DONATE a PROTON (H+).
b. Base=a substance which can ACCEPT a PROTON(H+).
c. Substances which can donate one proton are called monoprotic acids, substances which can donate 2 or more protons are called polyprotic acids. H2SO4, H3PO4, H2C2O4
d. Substances which can accept more than one proton are called polyprotic bases. S2-
S2- + H3O+ ----> HS- + H2O
HS- + H3O+ ----> H2S + H2O
e. Substances which can act as either Bronsted acids or bases are called amphiprotic.
Water is a good example.
f. HCl as an Acid & Ammonia as a base:
HCl + H2O ----> H3O+ + Cl-
NH3 + H2O ----> NH4+ + -OH
If we look at these equations there are really two acids and two bases in each reaction.
Conjugate acid-base pairs = a pair of compounds or ions that differ by the presence of an H+ unit are called a conjugate acid-base pair:
(1) Conjugate Acid= Base + H+
NH3 = base and NH4+= conjugate acid
(2) Conjugate Base= Acid - H+
HCl = acid and Cl- = conjugate base
g. Strong Acids=highly ionized-give up H+ easily,(highly ionized). See Figure 16.4- Page 598.
h. Strong Acid=Weak Conjugate Base
i. Usually the stronger acid and stronger base are written on the left hand side of the equation with the weaker acid and base on the right. All proton transfer reactions run from the stronger acid-base combination toward the weaker acid- base combination.
Figure 16.4 is useful in writing acid-base reactions.
B. Auto-Ionization of water-The hydronium ion-Water is the most common solvent used in chemistry and its ionization is an important factor in acid-base chemistry.
Kc = [H3O+] [OH-] / [H2O]2
1. For pure water [H3O+] = [OH-] = 1.0 x 10-7 mol/L
and in l L of water there are 55.6 moles of H2O; therefore, the concentration of water remains almost constant. A new constant called Kw is used: Kw= Kc[H2O]2= [H3O+] [OH-].
2. at 25oC Kw = (1.0 x 10-7 mol/L)(1.0 x 10-7 mol/L)
= 1 x 10-14
[H3O+] > [OH-] ----> solution is acidic
[H3O+] < [OH-] ----> solution is basic
[H3O+]= [OH-] ----> solution is neutral
3. If either the [OH-] or [H3O+] is known, the other value can be determined.
Kw / [H3O+] = [OH-] OR Kw / [OH-] =[H3O+]
4. The pH Concept- in 1909 Sorenson proposed the pH scale to avoid having to use values like
4.6 x 10-8.
a. pH = -log [H3O+] and [H3O+]= 10-pH and pOH= -log[OH-]
Therefore, since [H3O+] [OH-] = 1.0 x 10-14 =Kw
pH + pOH = 14 = pKw
b. Examples of calculation of pH-
these calculations used base 10 log not ln! Can use log tables or a calculator.
c. Measurement of pH
1. Acid-base indicators-usually some organic material which has one color in a certain range of pH and a different color in a different range of pH. Generally the pH range is not very narrow. Phenolphthalein indicator is colorless at pH < 8.2 but is pink pH=8.2-9.8.
b. The pH meter-A voltmeter which uses electrodes whose voltage depends on the
[H3O+] concentration.Discussed in detail later.
C. Relative Strengths of Strong Acids
1. In Water
HCl, HBr, HI seem to have about the same strength in H2O but in ethanol, the strength varies: HCl < HBr < HI
WHY? Leveling effect of water = H2O is much stronger base than Cl-, Br- or I- so all acids appear equally strong.
2. In other solvents
a. HF< HCl < HBr < HI and this trend is observed generally . H-Nonmetal acid strength increases
down in a group.
H-Nonmetal acid strength increases left to right
across a period: CH4 < NH3 < H2O < HF
b. For oxyacids, H + Nonmetal + O
Acid strength increases across a period as the
electronegativity of the central atom increases.
H3PO4 < H2SO4 < HClO4
Acid strength increases in a group with increasing oxidation state of central atom (or with increasing number of O atoms bonded to central atom).
HClO < HClO2 < HClO3 < HClO4
D. Strong Acids and Strong Bases versus Weak Acids and Weak Bases.
1. Strong Acids and Bases are completely ionized in water. The hydronium ion is the strongest acid that can exist in water. The hydroxide ion is the strongest base that can exist in water.
2. Weak Acids and Bases are not completely ionized in water.The extent of their ionization can be expressed as an equilibrium expression called Ka or Kb.
a. Writing Ka or Kb
For Weak Acids:
HA(aq) + H2O(l)
H3O+(aq) + A-(aq)Kc= [H3O+] [A-]
[HA] [H2O]
but [H2O] is constant, so
Kc[H2O] = Ka = [H3O+] [A-]
[HA]
Ka is called the acid ionization constant or the ionization constant.
Ka gives a quantitative method for comparing the relative strength of weak acids. The larger the value of Ka the stronger will be the acid. Appendix D on pages 1022-1023 lists Ka values for many acids.
For Weak Bases:
B(aq) + H2O(l)
BH+(aq) + -OH(aq)Kb= [BH+] [-OH]
[B]
Kb is the base ionization constant. The larger the value of Kb the stronger the weak base will be. Appendix D, pg 1023 lists Kb values for some bases.
b. Ka and Kb for conjugate acid/base pairs
Ka (for the acid) a 1/ Kb (for conjugate base)
Ka = Kw / Kb
so Ka X Kb = Kw
3. Calculations using Ka and Kb
4. Predicting reaction products using Ka and Kb
Whether a Bronsted acid/base reaction will occur in the forward direction depends on the relative strengths of the conjugate acid /base pairs (or on their Ka values)
A Bronsted acid donates a proton to the conjugate base of a weaker acid:
HNO2 + C2H3O2- -----> ?
Since HNO2 is a stronger acid than HC2H3O2, the
reaction proceeds:
HNO2 + C2H3O2- -----> HC2H3O2 + NO2-
acid base c. a. c. b.
E. Polyprotic Acids-Acids with more than one acidic hydrogen per formula unit. There is a Ka for each loss of H+. Generally the pH of a solution containing a polyprotic acid depends primarily on the extent of the first ionization step.
F. ACID - BASE Neutralizations
Reactions between acids and bases fall into four possible combinations:
1. Strong Acid + Strong Base ----> Salt + Water
a. Net ionic equation:
b. K = 1/Kw = 1 x 1014
c. The solution produced will be neutral (pH=7) at the equivalence point.
d. HCl + NaOH -----> Na+(aq) + Cl-(aq) + H2O
2. Strong Base + Weak Acid -----> Salt + Water
a. Net Ionic Equation:
b. Knet = Ka/Kw or 1/Kb (for A-)
c. The solution will be basic at the equivalence point, due to the reaction of the stronger conjugate base A- with water. The Kb will determine the pH of the solution.
d. Example
3. Strong Acid + Weak Base -----> Salt + Water
a. Net Ionic Equation:
Where B is the weak base and HB+ its conjugate acid.
b. Knet= Kb/Kw or 1/Ka since the reverse reaction is the reaction of the conjugate acid with water.
c. The solution at the equivalence point will be ACIDIC due to the reaction of HB+ with water. The pH will be determined by the Ka value of the conjugate acid.
d. Example
4. Weak Acid and Weak Base-more complicated and depends on the Ka of the acid and Kb of the base of the reactants.
a. Qualitative predictions can be made if K’s are known.
(1) IF Ka > Kb then the solution will be acidic since more H3O+ exists at equilibrium.
(2) When a weak acid and weak base are mixed in equal molar amounts, a salt is produced. The pH of the resulting solution depends on the relative K values of the conjugate acid and base.
b. Example
G. Solutions of Salts
Hydrolysis=reaction of a salt with water to change the [H3O+]/ [-OH].
The salt solution may be the result of
a. an acid - base reaction
b. adding a salt to water
1. How will dissolving NaCl in water effect the pH?
NaCl is the salt of a strong acid + a strong base.
NaCl(s) -----> Na+(aq) + Cl-(aq), pH‰7
Neither Na+ or Cl- react with water to give hydronium
or hydroxide ions. Na+ is a weak acid and Cl- is a weak base.
2. How will adding NH4Cl to water effect the pH?
NH4Cl is the salt of a strong acid + a weak base.
NH4Cl -----> NH4+(aq) + Cl-(aq)
Then:
Ka for NH4+ = Kw/Kb (for NH3)
[H3O+] > [OH-] so pH < 7
3. How will adding NaC2H3O2 to water effect the pH?
NaC2H3O2 is the salt of a weak acid + a strong base
NaC2H3O2 ----->Na+(aq) + C2H3O2-(aq)
Kb = Kw / Ka (for HC2H3O2)
[OH-] > [H3O+] so pH > 7
4. General Statements:
a. Salts of a strong acid and weak base give acidic solutions.
b. Salts of strong bases and weak acids give basic solutions.
c. Salts of strong acids and strong bases give neutral solutions.
5. Calculations
H. The Common Ion Effect-The limitation on acid (or base) dissociation caused by addition (or presence) of some species at equilibrium. This is related to LeChatelier’s Principle.
1. General Case for a Weak Acid:
HA(aq) + H2O(l)
H3O+(aq) + A-(aq) If A- is added prior to the equivalence point, the equilibrium will be shifted to the left, and will effect the pH of the solution. The solution must become less acidic.
2. General Case for a Weak Base:
B(aq) + H2O(l)
HB+(aq) + -OH(aq)If HB+ is added prior to the equivalence point the equilibrium will shift to the left and the pH will be effected. The solution must become less basic.
3. Sample Calculations.
I. Buffer Solution-protects against large changes in the pH of the solution even when strong acids or bases are added to the solution.
1. pH control is extremely important in: living systems, industry and chemical and biological research.
2. Buffer systems usually consist of two solutes
a. Weak acid and its salt:
CH3CO2H / Na+CH3CO2-
b. Weak base and its salt:
NH3 / NH4Cl
3. How Does a Buffer Work?
4. pH of a Buffer Solution
a. pKa = -log Ka
b. Consider the case of a weak acid/ salt buffer. HA (aq) + H2O(l)
H3O+(aq) + A-(aq)Ka= [H3O(aq)] [A-(aq)] = [H3O+(aq)] [Salt]
[HA(aq)] [HA(aq)]
Solving for [H3O+(aq)]= Ka [HA(aq)]
[Salt]
Take -log of both sides of the above equation.
pH = -log [HA(aq)] + ( -log Ka)
[Salt]
or pH= -log [HA(aq)] + pKa
[Salt]
or pH= pKa + log [Salt]
[HA(aq)]
This form of the equation is called the Henderson- Hasselbalch equation. It is very useful!
c. Using the same reasoning for a buffer of a weak base/salt, the equations below can be derived.
[-OH] = Kb [B]
[Salt]
pOH = pKb + log [Salt]
[B]
5. Preparation of a Buffer
a. If [Salt] = [Acid] then pH = PKa, pick an acid whose pKa is ± 1 of the desired pH.
b. Example of preparation of a buffer.
J. Acid-Base Titration Curves
1. End Point vs Equivalence Point
a. End point-the point in a titration when the acid-base indicator changes color.
b. Acid-base indicators are designed to change color near the equivalence point( the point in a titration at which one reactant has been exactly consumed by addition of the other reactant).
c. At the equivalence point:
moles acid = moles base
L acid x M acid= L base x M base
Remember: Molarity = mol/L or mol=L x M
2. Titration of a Strong Acid with a Strong Base(or vice versa)
a. The equivalence point will be at pH=7.
b. Graph
3. Strong Base titrated with Weak Acid
a. There will be a significant period in the titration where the pH remains almost constant = buffer region.
b. The pH of the solution at the equivalence point will be >7 (basic).
HA(aq) + -OH (aq)
H2O(l) + A-(aq) A-(aq) + H2O(l)
HA(aq) + -OH(aq) c. At the point in the titration when 1/2 of the acid has been neutralized, 1/2 HA remains and 1/2 A- has been formed . At this point, you have [A-]/[HA]=1, therefore the equation can be written:
pH = pK + log [A-] = pK + log 1 = pK + 0
[HA]
pH = pK (for weak acid or base)
This allows the Ka or Kb for a weak acid or base to be determined from an acid-base titration.
d. Graph
K. Lewis Acids and Bases-1930’s; G.N. Lewis
1. Definitions:
a. An acid is any substance that can ACCEPT a PAIR OF ELECTRONS from another atom to form a new bond. All metal cations, neutral compounds of Group 3A elements, etc.
b. A base is any substance that can DONATE A PAIR OF ELECTRONS to another atom to form a new bond. All negative ions and any molecule containing atoms with unshared electron pairs in the valence shell.
c. The new bond formed is a coordinate covalent bond.
2. Coordination number of a metal ion in a complex is the number of donor atoms bonded to the central ion, usually 4 or 6.
3. Examples of Lewis Acids and Bases and Coordination number.
coordination number = 4
4. Formation Constants-the extent to which a Lewis acid-base reaction occurs can be expressed in terms of an equilibrium constant. The greater the value for this constant, the greater the extent of the reaction.
