Transcript Chapter 19 Acids, Bases, and Salts

Acids, Bases, and Salts
By: Chris Squires
Effects of Acid Rain on Marble
(marble is calcium carbonate)
George Washington:
BEFORE acid rain
George Washington:
AFTER acid rain
Acids
have a
pH
less
than 7
OPERATIONAL ACIDS
Reacts with metals to form H2 gas.
HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)
Taste Sour
Conducts electricity.
Neutralizes BASES
Acids Turn Litmus Red
Blue litmus paper turns red in
contact with an acid
Sulfuric Acid = H2SO4
 Highest
volume
of any
chemical in
the U.S.
60 billion pounds/year
Acids React with Carbonates
and Bicarbonates
HCl + NaHCO3
Hydrochloric acid + sodium bicarbonate
NaCl + H2O + CO2
salt + water + carbon dioxide
An old-time home remedy for
relieving an upset stomach
OPERATIONAL BASES
BASES ARE SLIPPERY
BASES are Corrosive
BASE
BASES
BASES OPERATE
Bases
have a
pH
greater
than 7
Bases Turn Litmus Blue
Red litmus paper
turns blue in contact
with a base (and blue
paper stays blue).
Properties of Bases (metallic hydroxides)
 React
with acids to form water
and a salt.
 Taste bitter & Feel slippery
 Electrolytes in aqueous solution
 Change the color of indicators
(bases go Blue , with litmus).
Examples of Bases

Sodium hydroxide, NaOH
(lye for drain cleaner; soap)

Potassium hydroxide,
KOH (alkaline batteries)

Magnesium hydroxide,
Mg(OH)2 (Milk of Magnesia)
Acid Theories Base
+
H
H+
Acid-Base Theories
 a)
Arrhenius,
 b) Brønsted-Lowry
 c) Lewis (not covered).
Svante Arrhenius (1859-1927)
Theoretical Definitions

Arrhenius acids and bases
Acid: when dissolved in water increases H+.
Base: Substance when dissolved in water,
increases hydroxide ions OH-
1. Arrhenius Definition - 1887
produce hydrogen ions (H+)
in aqueous solution (HCl → H+ + Cl-)
 Bases produce OH- when dissolved
in water.
 Acids
(NaOH → Na1+ + OH1-)
 Limited
to aqueous solutions.
 Only one kind of base (hydroxides)
 NH3 (ammonia) could not be an
Arrhenius base: no OH- produced.
3 flaws with Arrhenius
1. A proton (H+) does not exist ALONE in
water, rather it combines with waters
lone pairs to form H3O+(aq)
2. NH3 (g) is a base when dissolved in
water; but does not contain the OH-(aq) 3.
3. Amphiprotics like HCO3-(aq) appear to
contain the H+ ion BUT turns red litmus
blue and ARE bases
Johannes Brønsted
(1879-1947)
Denmark
Thomas Lowry
(1874-1936)
England
2. Brønsted-Lowry - 1923
A broader definition than Arrhenius
 Acid is hydrogen-ion donor (H+ or
proton); base is hydrogen-ion acceptor.
 Acids and bases always come in pairs.
 HCl is an acid.
– When it dissolves in water, it gives
it’s proton to water.
HCl(g) + H2O(l) ↔ H3O+(aq) + Cl-(aq)
 Water is a base; makes hydronium ion.

B/L Definition

Brønsted–Lowry: must have both
1. an Acid: Proton donor
and
2. a Base:
Proton acceptor
A Brønsted–Lowry acid…
…must have a removable (acidic) proton.
HCl, H2O, H2SO4
A Brønsted–Lowry base…
…must accept this H+
NH3, H2O, CO3 2- and usually has negative charge
Why Ammonia is a Base
 Ammonia
can be explained as a
base by using Brønsted-Lowry:
NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq)
Ammonia is the hydrogen ion
acceptor (base), and water is the
hydrogen ion donor (acid).
This causes the OH1- concentration
to be greater than in pure water,
and the ammonia solution is basic
Acids and bases come in pairs
A “conjugate base” is the remainder of
the original acid, after it donates it’s
hydrogen ion
 A “conjugate acid” is the particle
formed when the original base gains a
hydrogen ion


Thus, a conjugate acid-base pair is related by
the loss or gain of a single hydrogen ion.
Acids and bases come in pairs

General equation is:
HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq)

Acid + Base ↔ Conjugate acid + Conjugate base
NH3 + H2O ↔ NH41+ + OH1base acid
c.a.
c.b.
 HCl + H2O ↔ H3O1+ + Cl1acid base
c.a.
c.b.
 Amphoteric – a substance that can act as
both an acid and base- as water shows

Conjugate Acids and Bases:


From the Latin word conjugare, meaning “to
join together.”
Reactions between acids and bases always
yield their conjugate bases and acids.
If can also be BOTH Acid and
Base…
...it is amphiprotic.
–
HCO3
–
HSO4
H2O
Three THEORETICAL definitions
Theory:
Acid
When
Arrhenius
increases H+
1880’s
Brønsted
proton donor
1923
Who
Lowry
ditto
same1923
Hydrogen Ions from Water
Water ionizes, or falls apart into ions:
H2O ↔ H+ + OH Called the “self ionization” of water
 Occurs to a very small extent:
[H+ ] = [OH-] = 1 x 10-7 M
 Since they are equal a neutral solution
results


Kw = [H+ ] x [OH-] = 1 x 10-14 M2
How to measure pH with wide-range paper
1. Moisten the pH
indicator paper strip
with a few drops of
solution, by using a
stirring rod.
2.Compare the color
to the chart on the vial
– then read the pH
value.
How Do We Measure pH?
– Litmus paper
• turns blue above
~pH = 7
• turns red below
~pH = 7
– An indicator
• Compound that
changes color in
solution.
Some of the
many pH
Indicators
and their
pH range
Other “p” Scales
The “p” in pH tells us to take the
negative log of the quantity
 Some similar examples are

– pOH –log [OH-]
– pKw –log Kw
The pH concept – from 0 to 14

pH = “hydrogen power”
 definition:
pH = -log[H+]
in neutral pH = -log(1 x 10-7) = 7
 in acidic solution [H+] > 10-7
 pH < -log(10-7)

– pH < 7 (from 0 to 7 is the acid range)
– in base, pH > 7 (7 to 14 is base range)
Calculating pOH
 pOH = -log [OH-]
+
 [H ]
 pH
x
[OH ]
=1x
+ pOH = 14
-14
2
10 M
pH and Significant Figures
 For
pH calculations, the hydrogen
ion concentration is usually
expressed in scientific notation
 [H+] = 0.0010 M = 1.0 x 10-3 M, and
0.0010 has 2 significant figures
 the pH = 3.00, with the two
numbers to the right of the decimal
corresponding to the 2 sig figs
STRONG V WEAK
STRONG VS WEAK
Strength
In any acid-base reaction, the
equilibrium favors the reaction that
moves the proton to the stronger base.
HCl(aq) + H2O(l)  H3O+(aq) + Cl–(aq)
Cl– equilibrium lies so far to the right K is not
measured (K>>1), so is not a base, it is neutral
Strength

Strong acids
completely dissociate
in water.
– Their conjugate bases
are actually so weak as
not to be able to act as
a base at all.

Weak acids partially
ionize in water.
Strength

Acids and Bases are classified acording
to the degree to which they ionize in
water:
– Strong are completely ionized in
aqueous solution; this means they
ionize 100 %
– Weak ionize only slightly in aqueous
solution
 Strength
is very different from
Concentration
Strength
– means it forms many
ions when dissolved (100 %
ionization)
 Mg(OH)2 is a strong base- it falls
completely apart (nearly 100%
when dissolved).
–But, not much dissolves- so it
is not concentrated
 Strong
Strong Acid Dissociation
(makes 100 % ions)
Weak Acid Dissociation
(only partially ionizes)
STRONG V WEAK ACIDS
Organic Acids (those with carbon)
Organic acids all contain the carboxyl group,
(-COOH), sometimes several of them.
CH3COOH – of the 4 hydrogen, only 1 ionizable
(due to being bonded to the highly electronegative Oxygen)
The carboxyl group is a poor proton donor, so
ALL organic acids are weak acids.
Measuring strength
Ionization is reversible for WEAK acids:
HA + H2O ↔ H+ + A(Note that the arrow
 This makes an equilibrium goes both directions.)
 Acid dissociation constant = Ka
+
(Note that water is NOT shown,
 Ka = [H ][A ]
because its concentration is
[HA]
constant, and built into Ka)
 Stronger acid = more products (ions),
thus a larger Ka

What about bases?

Strong bases dissociate completely.

MOH + H2O ↔ M+ + OH-

Base dissociation constant = Kb

Kb =
(M = a metal)
[M+ ][OH-]
[MOH]
 Stronger
base = more dissociated
ions are produced, thus a
 larger Kb
and a small Ka.
Calculating pH from Ka
Calculate the pH of a 0.30 M solution of
acetic acid, C2H3O2H, at 25°C.
Ka for acetic acid at 25°C is 1.8  10-5.
Calculating pH from Ka
Use the ICE table:
[C2H3O2], M
Initial
Change
Equilibrium
0.30
–x
0.30 – x
[H3O+], M [C2H3O2−], M
0
+x
x
0
+x
x
Calculating pH from Ka
Use the ICE table:
Initial
Change
Equilibrium
[C2H3O2], M
[H3O+], M
[C2H3O2−], M
0.30
0
0
–x
+x
+x
0.30 – x ≈ 0.30
x
x
Simplify: how big is x relative to 0.30?
Calculating pH from Ka
Now,
(1.8  10-5) (0.30) = x2
5.4  10-6 = x2
2.3  10-3 = x
pH = = 2.64
Calculating Ka from the pH

The pH at 25°C is 2.38 for a 0.10 M solution
of formic acid, HCOOH,. Calculate Ka
Ka, requires equilibrium concentrations. We can find [H3O+],
which is the same as [HCOO−], from pH, so pH = X
Calculating Ka from the pH
pH = –log [H3O+]
– 2.38 = log [H3O+]
10-2.38 = 10log [H3O+] = [H3O+]
4.2  10-3 = [H3O+] = [HCOO–] = X
Calculating Ka from pH
All values
In Mol/L
Initially
[HCOOH]
[H3O+]
[HCOO−]
0.10
0
0
Change
–4.2  10-3
Equilibrium
0.10 – 4.2  10-3
= 0.0958 = 0.10
+4.2  10-3 +4.2  10-3
4.2  10-3
4.2  10 - 3
Calculating Ka from pH
Ka =
[4.2  10-3] [4.2  10-3]
[0.10]
= 1.8  10-4
Calculating Percent Ionization
[A-]eq = [H3O+]eq = 4.2  10-3 M
[A-]eq + [HCOOH]eq = [HCOOH]initial= 0.10 M
%rx = [4.2  10-3 / 0.10] (x 100)
= 4.2%
Weak Bases
The Kb for this reaction is
pH of Basic Solutions
What is the pH of a 0.15M solution of NH3?
[NH4+] [OH−]
Kb =
= 1.8  10-5
[NH3]
pH of Basic Solutions
[NH3]
[NH4+]
[OH−]
Initial
0.15
0
0
Equil
0.15 - x  0.15
x
x
pH of Basic Solutions
2
(x)
1.8  10-5 =
(0.15-X)
Ignore X
(1.8  10-5) (0.15) = x2
2.7  10-6 = x2
-3 = x2
1.6

10
pOH = –log (1.6  10-3)
pOH = 2.80
pH = 11.20
Strength vs. Concentration
The words concentrated and dilute tell
how much of an acid or base is
dissolved in solution - refers to the # of
moles of acid or base
 The words strong and weak refer to
the extent of ionization
 Is a concentrated, weak acid possible?

Acid and Base Strength
Acetate is a stronger
base than H2O, LOWER
on table,
The stronger base
“wins” the proton.
HC2H3O2(aq) + H2O
H3O+(aq) + C2H3O2–(aq)
TITRATIONS
Titration
 Titration
is the process of adding a
known amount of solution of known
concentration to determine the
concentration of another solution
 Remember? - a balanced equation is
a mole ratio
 The equivalence point is when the moles
of hydrogen ions is equal to the moles
of hydroxide ions (= neutralized!)
Titration
 The
concentration of acid (or base)
in solution can be determined by
performing a neutralization reaction
–An indicator is used to show
when neutralization has occurred
–Often we use phenolphthaleinbecause it is colorless in neutral
and acid; turns pink in base
Steps - Neutralization reaction
#1. A measured volume of acid of
unknown concentration is added to
a flask
#2. Several drops of indicator added
#3. A base of known concentration is
slowly added, until the indicator
changes color; measure the volume
Neutralization
 The
solution of known
concentration is called the
standard solution
– added by using a buret
 Continue
adding until the indicator
changes color
– called the “end point” of the titration
Polyprotic Acids?
 Some
compounds have more than
one ionizable hydrogen to release
 HNO3 nitric acid - monoprotic
 H2SO4 sulfuric acid - diprotic - 2 H+
 H3PO4 - triprotic - 3 H+
 IMPT
for TITRATING BUT NOT FOR
pH calculating…why?
Salt Hydrolysis
A
salt is an ionic compound that:
–comes from the anion of an acid
–comes from the cation of a base
–is formed from a neutralization
reaction
–some neutral; others acidic or basic
 “Salt hydrolysis” - a salt that reacts
with water to produce an acid or base
Salt Hydrolysis
 To
see if the resulting salt is
acidic or basic, check the by
dissociating the salt
NaCl, a neutral salt
(NH4)2SO4, acidic salt
CH3COOK, basic salt
Buffers
 Buffers
are solutions in which the
pH remains relatively constant,
even when small amounts of acid
or base are added
–made from a pair of chemicals:
a weak acid and one of it’s
salts; or a weak base and one
of it’s salts
Buffers
A
buffer system is better able to
resist changes in pH than pure water
 Since it is a pair of chemicals:
–one chemical neutralizes any acid
added, while the other chemical
would neutralize any additional
base
–AND, they produce each other
in the process!!!
Buffers
 Example: Ethanoic (acetic) acid
and sodium ethanoate (also
called sodium acetate)
 The buffer capacity is the
amount of acid or base that can
be added before a significant
change in pH
Buffers
 The
two buffers that are crucial to
maintain the pH of human blood are:
1. carbonic acid (H2CO3) & hydrogen
carbonate (HCO31-)
2. dihydrogen phosphate (H2PO41-) &
monohydrogen phoshate (HPO42-)
Aspirin (which
is a type of
acid)
sometimes
causes
stomach
upset; thus by
adding a
“buffer”, it
does not
cause the
acid irritation.
Bufferin is
one brand of
a buffered
aspirin that
is sold in
stores.