Chapter 9 Acids, Bases
Chapter 9 Acids, Bases
Acids & Bases are substances that affect the pH of solutions.
Acids & Bases typically are, or behave as, IONIC compounds.
Acids:
Are corrosive
Taste sour
React with indicators
Neutralize bases
Ex. HCl (hydrochloric acid), H2SO4 (sulfuric acid)
Bases:
Are slippery
Taste bitter
React with indicators
Neutralize acids
Ex. NaOH (sodium hydroxide), NH4OH (ammonium hydroxide)
Nomenclature
Acids
Most are “hydrogen” bonded with an anion
Examples: HNO3 HC2H3O2
Bases
Most are metal hydroxides
Examples: KOH Ba(OH)2
Base Formulas
Formulas generally end with “OH”
Ex. NaOH KOH Ca(OH)2
Base Names
All names are generally two parts ending in the word “hydroxide”.
The first part is the name of the metal element bonded to the hydroxide.
Examples
Sodium hydroxide NaOH
Calcium hydroxide Ca(OH)2
Ammonia NH3*
This is the exception to this rule; however, when dissolved in water ammonia becomes ammonium hydroxide (NH4OH)
Acid Formulas
Formulas generally start with “H”
Ex. HCl HNO3 H2SO4
Acid Names
All names are two parts ending in the word “acid”.
The first part depends on how many oxygen atoms are in the compound.
No Oxygen
Hydro ______-ic acid
The blank is the root of the name for the element to which hydrogen is bonded.
Ex. HCl hydrochloric acid
HI hydroiodic acid
Oxyacids Know the anion!
___ate = ____ic acid
Ex. HNO3 H3PO4
____ite = ____ous acid
Ex. HNO2 H3PO3 H2SO3
Summary
Acids with LESS Oxygens end in “ous acid” (remember “ite” ions)
Acids with MORE Oxygens end in “ic acid” (remember “ate” ions)
Definitions of Acids & Bases
Acids
Arrhenius - acids donate H+ (in soln)
Bronsted-Lowery -acids donate H+ (in soln)
Bases
Arrhenius - bases donate OH- (in soln)
Bronsted-Lowery - bases accept H+ (in soln)
*Coordinate covalent bond
Conjugate Acid-Base Pairs
The transfer of protons illustrates the characteristics of conjugate pairs
HNO2 + H2O <==> H3O+ + NO2-
NO2- is the conjugate base of HNO2
H3O+ is the conjugate acid of H2O
Proticity: Acids can be classified according to the number of hydrogen ions (protons) they can transfer per molecule during an acid-base reaction.
Monoprotic: HCl, HNO3
Diprotic: H2CO3
Triprotic: H3PO4
Acid-Base Strength
Strong
“ions” completely dissociate in water
ACIDS:
HCl, HBr, HI,
HClO4, H2SO4, HNO3
BASES:
LiOH, NaOH, KOH,
Ca(OH)2, Sr(OH)2, Ba(OH)2
Weak: All non-strong acids & bases
“ions” partially dissociate in water
Equilibrium systems
A comparison of the number of acidic species present in strong acid and weak acid solutions of the same concentration.
Weak A/B equilibrium: Two reactions (forward & reverse) occur at the same rate
HA <==>H+ + A-
BOH <=> B+ + OH-
Equilibrium expressions are ways to show the mathematical relationships
Keq = [Products]n
[Reactants]m
n & m are the coefficients of each substance
Ionization Constants for Acids & Bases
HA(aq) + H2O(l) <==> H3O+(aq) + A-(aq)
Ka = ------
B(aq) + H2O(l) <==> BH+(aq) + OH-(aq)
Kb = ------
Neutralization reactions - a special type of DR rxn
AX + BY --> AY + BX
HCl + KOH --> KCl + HOH
Acid + Base --> Salt + Water
To balance these rxns. Balance the H in the acid with the OH in the base :)!
For a complete neutralization reaction (that reaches an “equivalence point”), stoichiometric equivalents of the acid and base must be used.
Neutralization equations
HCl(aq) + NaOH(aq) --> NaCl(aq) + H2O(l)
H2SO4 + Ba(OH)2 -->
H3PO4 + KOH -->
HNO3 + Al(OH)3 -->
pH
Self-Ionization of Water: Water molecules can break apart when they collide
H2O(l) <==> H+(aq) + OH-(aq)
Kw = ------
Kw = 1.0 x 10-14 M2
Adding an acid or a base changes the relative amounts of [H+] and [OH-] but not the value of Kw.
Ionic Concentration
If [H+] = [OH-] the solution is neutral
If [H+] > [OH-] the solution is acidic
If [H+] < [OH-] the solution is basic
[H+] x [OH-] = 1.0 x 10-14M2
pH: a logarithmic scale of a solution’s molar hydrogen (hydronium) ion concentration
This is a way to express the relative acidity/basicity of a solution.
pH = -log[H+]
High [H+] causes low pH
Low [H+] causes high pH
Therefore, strong acids have lower pH!
pOH = -log[OH-]
pH scale
0 - 14 is the usual range
pH < 7 = acid
pH > 7 = base
pH = 7 = neutral
pH + pOH = 14
[H+] à pHCalculations
If the [H+] = 3.35 x 10-5 M, what is the pH of the solution?
If the [OH-] = 2.8 x 10-4M, what is the pH of the solution?
pH à [H+] calculations
What is the [H+] for a solution with a pH = 3.92?
pH = -log[H+]
3.92 = -log[H+]
-3.92 =log[H+]
10-3.92= [H+]
[H+] = 1.20 x 10-4M
Practice: determine the [H+] for the solutions with the following values.
pH = 7.55 pH = 10.4
pH = 2.12 pOH = 4.5
pKa: to express acid strength!
pKa = -logKa
Acetic acid (HC2H3O2)
Ka = 1.8 x 10-5
Salt Hydrolysis
Some aqueous salt solutions have the ability to split (hydrolyze) water and form compounds which result in larger [H+] or [OH-] in the solution.
Example: Aluminum chloride
AlCl3(aq) --> Al+3(aq) + 3Cl-(aq)
Cation of WB Anion of WA
Aluminum ion will react with OH- in solution:
Remember: H2O <==> H+ + OH-
Al+3(aq) + H2O(l) <==> Al(OH)3(aq) + 3H+(aq)
Chloride ion will NOT react with H+ in solution!
Rules for Determining pH Strength wins!
Strong Acid + Strong Base --> Neutral soln
HCl + NaOH --> NaCl + H2O
Strong Acid + Weak Base --> Acidic soln
HCl + Al(OH)3 --> AlCl3 + H2O
Weak Acid + Strong Base --> Basic soln
H2S + NaOH --> Na2S + H2O
Weak Acid + Weak Base --> depends on the salt
HNO2 + NH4OH --> NH4NO2 + H2O
Buffers
Buffers are solutions in which the pH remains relatively constant when small amounts of acid or base are added
Two active chemical species:
A substance to react with & remove added base
A substance to react with & remove added acid.
Buffers are solutions of a weak acid and one of its conjugate base OR a weak base and one of its conjugate base.
Carbonic acid and Sodium bicarbonate
H2CO3 <==> H+ + HCO3-
NaHCO3 --> Na+ + HCO3-
Buffering Action in Human Blood
H2CO3 <==> H+ + HCO3-
High concentration High concentration
Ratio: 1 : 10
Add a base [OH-] and the equilibrium position shifts ; pH doesn’t change much
Add an acid [H+] and the equilibrium position shifts ; pH doesn’t change much
Reason: high [ ] of acid and anion can accommodate large shifts of EQ position.
Lots of acid is produced in the body daily.
Henderson-Hasselbalch Equation
Buffers are most effective when the acid-to-conjugate base ratio is 1:1.
Ka = ------
If [HA] = [A-], then Ka = [H3O+], thus pKa = pH
H-H eqn:
pH = pKa + log------
The implication of this is: If [A-] > [HA], then pH > pKa and vice versa
Titration
At the completion of the reaction (equivalence point) the
# moles acid = # moles base
So,
MaVa = MbVb
Chemical Titration
This process can be done for any reaction in which a stoichiometric equivalence is reached and can be identified by an indicator
At the equivalence point an indicator will change color permanently.