Transcript Chapter 8: Acids & Bases

Chapter 8: Acids & Bases

Acids are all around us.
– Foods like lemons, limes, and oranges
contain acidic compounds.
– When wine is oxidized, it becomes acetic
acid.
– Our muscles produce lactic acid.
– Our stomach uses hydrochloric acid to
break down food.
Acids
Arrhenius definition – an acid
is a substance that produces
H+ ion when added to water.
HCl(aq)  H+(aq) + Cl-(aq)
 Acids taste sour, are
electrolytes, and neutralize
bases.

Bases
Arrhenius definition – a base
is a substance that produces
OH- when added to water.
NaOH(aq)  Na+(aq) + OH-(aq)
 Bases taste bitter, have a
slippery, soapy feel, and
neutralize acids.

Bronsted-Lowery Theory

Expanded definition of acids & bases.
– Acid = a proton (H+) donor
– Base = a proton acceptor
Important note about the use of H+ in
equations.
 For weak acids and bases, the
reactions are reversible, equilibria.

Bronsted-Lowery Theory

HC2H3O2 + H2O  H3O+ + C2H3O2-

NH3 + H2O  NH4+ + OH-
Conjugates

Write the conjugate base of an acid.
– Remove an H+
1. HBr – H+ = Br−
2. H2S – H+ = HS−

Write the conjugate acid of a base.
– Add an H+
1. NO2− + H+ = HNO2
2. NH3 + H+
= NH4+
Learning Check
1. The conjugate base of HCO3− is
a. CO32−
b. HCO3−
c. H2CO3
2. The conjugate acid of HCO3− is
a. CO32−
b. HCO3−
c. H2CO3
3. The conjugate base of H2O is
a. OH−
b. H2O
c. H3O +
4. The conjugate acid of H2O is
a. OH−
b. H2O
c. H3O+
Strong Acids
Strong acids completely ionize
(dissociate) in water.
 Strong acids are also strong
electrolytes.
 There are six: HClO4, H2SO4, HI, HBr,
HCl, and HNO3.

Weak Acids
Weak acids only partially dissociate to
produce ions in solution.
 Weak acids are weak electrolytes.
 Too many to list, but some common
ones are: H3PO4, HC2H3O2, HF, and
HC6H7O6.

Comparison
Strong Bases
Strong bases completely ionize in water
and are also strong electrolytes.
 Only group 1A and 2A hydroxides are
strong bases.
 All other metal hydroxides are insoluble
in water.

Weak Bases
Weak bases only partially react with
water (accepting a proton) and are,
thus, weak electrolytes.
 Ammonia, NH3, is the most common
weak base.
 Lone pair on N group will accept a
proton.
 Other organic weak bases.

Ionization of Water
Water can act as both an acid and a
base.
 Any substance that does this is called
amphoteric.
 Pure water – two water molecules will
occasionally react with each other
where one is an acid and one is a base.

Ionization of Water
H2O + H2O  H3O+ + OH [ X ] = symbol for molarity.
 In pure water, [H3O+] = [OH-].
 Kw = [H3O+] x [OH-]; where Kw = 1 E-14.
 Thus, [H3O+] = [OH-] = ____________

Acidic, Basic, and Neutral Solutions
Calculating [H3O+] and [OH-]

If we know [H3O+], then [OH-] =

If we know [OH-], then [H3O+] =

Ex) [H3O+] = 2.5 E-5

Ex) [OH-] = 4.8 E-3
pH Scale
One convenient method for measuring
the acidity or basicity of a solution is to
use the pH scale.
 pH is a logarithmic (log) scale and equal
to: pH = -log[H3O+].
 pOH = -log[OH-].
 pH + pOH = 14.

pH Scale

A word about significant figures.

2.4 x 10-3 M
pH = 2.62
Red numbers are the significant digits.
 Blue numbers are exact numbers.

Guide to Using Your Calculator
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1.
2.
3.
4.
Basic Calculators
Enter concentration
Press “log” key
Change the sign
Record answer to
proper s.f.’s

1.
2.
3.
4.
TI-83 or TI-89
Press negative sign
Press “log” key
(select in catalog)
Enter concentration,
close parenthesis
Enter key and round
to proper s.f.’s
pH to Concentration
To convert a pH back to a
concentration, you will use the “antilog”
key = 10x.
 [H3O+] = 10-pH
 [OH-] = 10-pOH
 Can also use the universal power (^)
key.

Comparison of Values
Measuring pH

Can be done with a meter, pH paper, or
an indicator.
Fill in the Chart
[H3O+]
[OH-]
pH
pOH
1.8 E-5
3.72
3.4 E-2
5.48
A/B
Reactions of Acids

An acid will react with most metals.
– Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g)
Reactions of Acids

Acids react with any carbonate (CO3-2) and
bicarbonate (HCO3-) to generate CO2.
Environmental Note: Acid Rain
Normally, rain is slightly acidic – pH of
5.5 to 6.2 – due to dissolved CO2.
 Burning fossil fuels, which contain small
amounts of Sulfur, produces SO3.
 This is converted to H2SO4.


H2SO4 + CaCO3  CaSO4 + H2O + CO2
Neutralization
Acids neutralize bases and vice versa.
 Acid + Base  Water + Salt
 HCl + NaOH  H2O + NaCl
 Balancing more complex acid-base
neutralization.

– Each H+ needs one OH-.
– Each H+ and OH- makes one H2O.
Solution Stoichiometry
Acid-base titration – can use a known
acid or base solution to determine an
unknown counterpart.
 Endpoint – when all of the unknown
acid or base has reacted.
 Indicator – a substance that changes
color when it changes pH.

Solution Stoichiometry
Requires precise
glassware to deliver the
known solution = buret.
 Stopcock allows for
delivery drop-by-drop.
 Volumes can be read to
nearest 0.05mL.

Buffers
When a small amount of acid or base is
added to pure water, the pH swings
drastically.
 Some solutions, though, will resist these
wild swings in pH and are called buffer
solutions.
 Made from a ___________ and a salt
containing the ______________.

Buffers
1.0L of pure water + 0.0100moles
(0.365g) of HCl.
 [H3O+] = 0.010 mol / 1.0L = 0.010M
 pH =
 1.0L of pure water + 0.0100moles
(0.400g) of NaOH.
 [OH-] = 0.010 mol / 1.0L = 0.010M
 pH =

Buffers
Buffer of HF and NaF (note: Na+ is a
spectator ion).
 Ideal buffer would contain a 50 / 50
mixture of each.
 1.0L of 0.10 moles HF (2.0g) and 0.10
moles of NaF (4.2g) will have a pH of
3.17.

Buffers
HF + H2O  H3O+ + F50%
50% (from NaF)
 Addition of strong acid

– reacts with F- ion to generate more HF

Addition of strong base
– reacts with HF to produce water plus more
F-
Buffers
Buffer plus 0.010 moles of HCl.
 pH = 3.08 (from 3.17).
 Buffer plus 0.010 moles of NaOH.
 pH = 3.26 (from 3.17).

Buffers in Blood
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
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1.
2.
Normal blood pH is 7.35 to 7.45.
Outside this range, cells cannot
function properly.
Two buffer systems are present to
maintain this pH.
H2CO3 / HCO3H2PO4- / HPO4-2