Transcript Acids & Bases - Rogue Community College

What’s
wrong in
this picture?
Acids & Bases
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)
Baking soda (NaHCO3)
Litmus is a
vegetable dye
obtained from
certain lichens
found principally in
the Netherlands.
Typical with Acids
Typical with Bases
The difference between the aqueous solution
processes of ionization and dissociation.
Nomenclature
Acids
 Most are
“hydrogen”
bonded with an
anion
Examples:



HNO3 (nitric acid)
HC2H3O2 (acetic acid)
HCl (hydrochloric acid)
Bases
 Most are metal
hydroxides
Examples:




NaOH (sodium hydroxide)
KOH (potassium hydroxide)
Ba(OH)2 (barium hydroxide)
NH4OH* (ammonium
hydroxide)
Definitions
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+ + NO2NO2- is the conjugate base of HNO2
H3O+ is the conjugate acid of H2O
Protocity
Monoprotic

HCl, HNO3
Diprotic

H2CO3
Triprotic

H3PO4
Acids can be
classified according
to the number of
hydrogen ions
(protons) they can
transfer per
molecule during an
acid-base reaction.
Acid-Base Strength
(You can dilute an acid or a base but you can’t change its 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
 “ions” partially
dissociate in water
All non-strong acids &
bases
These exist as
equilibrium systems in
solution, thus, their
“weakness” exists
within a range defined
by their Keq values.
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+ + ABOH <=> 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)
B(aq) + H2O(l) <==> BH+(aq) + OH-(aq)
Ka = ---------------------
Kb = ---------------------
Neutralization reactions
- a special type of DR rxn
AX + BY --> AY + BX
HCl + KOH --> HOH + KCl
Acid + Base --> Water + Salt
To balance these rxns. Balance the H in the acid
with the OH in the base :)!
For a complete reaction, stoichiometric
equivalents of the acid and base must be used.
Neutralization equations
HCl(aq) + NaOH(aq) --> H2O(l) + NaCl(aq)
H2SO4 + Ba(OH)2 -->
H3PO4 + KOH -->
HNO3 + Al(OH)3 -->
The acid-base reaction between
sulfuric acid and barium hydroxide
produces the insoluble salt barium
sulfate.
Calculations
A sample of 0.0084 mol HCl is dissolved in
water to make 1500 mL solution. Calculate the
molarity of the HCl solution and the [H3O+].
Self-Ionization of Water (pH is a
derivative of this concept)
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
The relationship between H3O+
and OH- in aqueous solution is an
inverse proportion.
Calculations
If the [OH-] = 3.5 x 10-3 M, what is [H+]?
pH: a logarithmic scale of a solution’s hydrogen
(hydronium) ion concentration (molarity)
This is a way to express the relative
acidity/basicity of a solution.
pH = -log[H+]
 Therefore, each difference in pH of 1.0 is
equivalent to a concentration change by a
factor of 10



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
Calculations
If the [H+] = 3.35 x 10-5 M, what is the
pH of the solution?

On your calculator:
- log (3.35 x 10-5) =
Calculations
If the [OH-] = 2.8 x 10-4M, what is the
pH of the solution?
pH --> [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
[H+] = 2.82 x 10-8 M
pH = 10.4
[H+] = 3.98 x 10-11 M
pH = 2.12
[H+] = 7.86 x 10-3 M
pOH = 4.5
[H+] = 3.16 x 10-10 M
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 sol’n

HCl + NaOH --> NaCl + H2O
Strong Acid + Weak Base --> Acidic sol’n

HCl + Al(OH)3 --> AlCl3 + H2O
Weak Acid + Strong Base --> Basic sol’n

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+ + HCO3NaHCO3 --> Na+ + HCO3-
Ka = 1.7 x 10-3
Buffering Action in Human Blood
H2CO3 <==> H+ + HCO3High 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.
Buffer Systems
Titration
At the completion of a neutralization reaction (equivalence
point) the
# moles acid = # moles base
So,
MaVa = MbVb
but, keep the reaction stoichiometry in mind.
Diagram showing setup
for titration procedures.
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.
Calculations
How many mL of 0.10M NaOH solution
are needed to neutralize 15 mL of
0.20M H3PO4 solution?