Salts, Acids, and Bases
Chapter 11
1
Solutions
• Solutions are homogeneous mixtures
• Solute is the dissolved substance
– Seems to “disappear” or “Takes on the state” of the solvent
• Solvent is the substance the solute dissolves in
– Does not appear to change state
– When both solute and solvent have the same state, the solvent
is the component present in the highest percentage
• Solutions in which the solvent is water are called
aqueous solutions
– Water is often called the universal solvent
2
Solubility
• When one substance (solute) dissolves in another
(solvent) it is said to be soluble
– Salt is soluble in Water,
– Bromine is soluble in methylene chloride
• When one substance does not dissolve in another they
are said to be insoluble
– Oil is insoluble in Water
• There is usually a limit to the solubility of one substance
in another
– Gases are always soluble in each other
– Some liquids are always mutually soluble
3
Solutions & Solubility
• Molecules that are similar in structure tend to form
solutions
– Like dissolves like
• The solubility of the solute in the solvent depends on the
temperature
– Higher Temp = Larger solubility of solid in liquid
– Lower Temp =Larger solubility of gas in liquid
• The solubility of gases depends on the pressure
– Higher pressure = Larger solubility
4
The precipitation reaction
that occurs when yellow
potassium chormate,
K2CrO4 (aq), is mixed
with a colorless barium
nitrate solution,
Ba(NO3)2 (aq)
5
The Solution Process Ionic Compounds
• When ionic compounds dissolve in water
they dissociate into ions
– ions become surrounded by water molecules hydrated
• When solute particles are surrounded by
solvent molecules we say they are solvated
6
When solid sodium
chloride dissolves,
the ions are
dispersed randomly
throughout the
solution
7
Polar water molecules interact with the positive and
negative ions of a salt
8
Describing Solutions - Qualitatively
• A concentrated solution has a high proportion of solute
to solution
• A dilute solution has a low proportion of solute to
solution
• A saturated solution has the maximum amount of solute
that will dissolve in the solvent
– Depends on temp
• An unsaturated solution has less than the saturation limit
• A supersaturated solution has more than the saturation
limit
– Unstable
9
Solution Concentration
• Parts Per Million
• PPM = grams of solute per 1,000,000 g of
solution
• PPM = mg of solute per liter of solution
– 1000 mg = 1 g
• Mass of Solution = Mass of Solute + Mass of Solvent
10
Chemical Packages - Moles
• We use a package for atoms and molecules
called a mole
• A mole is the number of particles equal to the
number of Carbon atoms in 12 g of C-12
• One mole = 6.022 x 1023 units
• The number of particles in 1 mole is called
Avogadro’s Number
• 1 mole of C atoms weighs 12.01 g and has 6.02 x 1023 atoms
11
One-mole
samples of iron
(nails), iodine
crystals, liquid
mercury, and
powdered sulfur
12
Example #1
Compute the number of moles
and number of atoms in 10.0 g of Al
 Use the Periodic Table to determine the
mass of 1 mole of Al
1 mole Al = 26.98 g
 Use this as a conversion factor for
grams-to-moles
1 mol Al
10.0 g Al x
 0.371 mol Al
26.98 g
13
Example #1
Compute the number of moles
and number of atoms in 10.0 g of Al
 Use Avogadro’s Number to determine the
number of atoms in 1 mole
1 mole Al = 6.02 x 1023 atoms
 Use this as a conversion factor for
moles-to-atoms
23
6.02 x 10 atoms
0.371 mol Al x
 2.23 x 1023 Al atoms
1 mol Al
14
Example #2
Compute the number of moles
and mass of 2.23 x 1023 atoms of Al
 Use Avogadro’s Number to determine the
number of atoms in 1 mole
1 mole Al = 6.02 x 1023 atoms
 Use this as a conversion factor for
atoms-to-moles
1 mol Al
2.23x 10 Al atoms x
 0.370mol Al
23
6.02 x 10 atoms
23
15
Example #2
Compute the number of moles
and mass of 2.23 x 1023 atoms of Al
 Use the Periodic Table to determine the
mass of 1 mole of Al
1 mole Al = 26.98 g
 Use this as a conversion factor for
moles-to-grams
26.98 g
0.370 mol Al x
 9.99 g Al
1 mol Al
16
Molar Mass
• The molar mass is the mass in grams of one
mole of a compound
• The relative weights of molecules can be
calculated from atomic masses
water = H2O = 2(1.008 amu) + 16.00 amu
= 18.02 amu
• 1 mole of H2O will weigh 18.02 g, therefore
the molar mass of H2O is 18.02 g
• 1 mole of H2O will contain 16.00 g of oxygen
and 2.02 g of hydrogen
17
Solution Concentration
Molarity
• moles of solute per 1 liter of solution
• used because it describes how many
molecules of solute in each liter of solution
• If a sugar solution concentration is 2.0 M ,
1 liter of solution contains 2.0 moles of
sugar, 2 liters = 4.0 moles sugar, 0.5 liters
= 1.0 mole sugar, etc.
moles of solute
molarity =
liters of solution
18
Properties of Acids
• Sour taste
• Change color of vegetable dyes
• React with “active” metals
– Like Al, Zn, Fe, but not Cu, Ag or Au
Zn + 2 HCl ZnCl2 + H2
– Corrosive
• React with carbonates, producing CO2
– Marble, baking soda, chalk
CaCO3 + 2 HCl CaCl2 + CO2 + H2O
• React with bases to form ionic salts
– And often water
19
Properties of Bases
•
•
•
•
Also Known As Alkalis
Taste bitter
Feel slippery
Change color of vegetable dyes
– Different color than acid
– Litmus = blue
• React with acids to form ionic salts
– And often water
– Neutralization
20
Arrhenius Theory
• Acids ionize in water to H+1 ions and anions
• Bases ionize in water to OH-1 ions and
cations
• Neutralization reaction involves H+1
combining with OH-1 to make water
• H+ ions are protons
• OH- ions are called hydroxide ions
• Definition only good in water solution
21
Strength of Acids & Bases
• The stronger the acid, the more willing it is to donate H
• Strong acids donate practically all their H’s
HCl + H2O  H3O+1 + Cl-1
• Strong bases will react completely with water to form
hydroxides
CO3-2 + H2O HCO3-1 + OH-1
• Weak acids donate a small fraction of their H’s
– The process is reversible, the conjugate acid and conjugate
base can react to form the original acid and base
HC2H3O2 + H2O  H3O+1 + C2H3O2-1
• Only small fraction of weak base molecules pull H off
water
HCO3-1 + H2O H2CO3 + OH-1
22
Graphical
representation of the
behavior of acids in
aqueous solution
23
Multiprotic Acids
• Monoprotic acids have 1 acid H,
diprotic 2, etc.
– In oxyacids only the H on the O is
acidic
• In strong multiprotic acids, like
H2SO4, only the first H is strong;
transferring the second H is usually
weak
H2SO4 + H2O  H3O+1 + HSO4-1
HSO4-1 + H2O  H3O+1 + SO4-2
24
Water as an Acid and a Base
• Amphoteric substances can act as either an
acid or a base
– Water as an acid, NH3 + H2O  NH4+1 + OH-1
– Water as a base, HCl + H2O  H3O+1 + Cl-1
• Water can even react with itself
H2O + H2O  H3O +1 + OH-1
25
Autoionization of Water
• Water is an extremely weak electrolyte
– therefore there must be a few ions present
H2O + H2O  H3O+1 + OH-1
• all water solutions contain both H3O+1 and OH-1
– the concentration of H3O+1 and OH-1 are equal
– [H3O+1] = [OH-1] = 10-7M @ 25°C
• Kw = [H3O+1] x [OH-1] = 1 x 10-14 @ 25°C
– Kw is called the ion product constant for water
– as [H3O+1] increases, [OH-] decreases
26
Acidic and Basic Solutions
• acidic solutions have a larger [H+1] than [OH-1]
• basic solutions have a larger [OH-1] than [H+1]
• neutral solutions have [H+1]=[OH-1]= 1 x 10-7 M
[H+1]
-14
1
x
10
=
[OH-1]
[OH-1]
-14
1
x
10
=
[H+1]
27
Example #2
Determine the [H+1] and [OH-1] in a
10.0 M H+1 solution
 Determine the given information and the
information you need to find
Given [H+1] = 10.0 M
Find [OH-1]
 Solve the Equation for the Unknown
Amount
1
Kw  [H ] x [OH ]
Kw
-1
[OH ]  1
[H ]
-1
28
Example #2
Determine the [H+1] and [OH-1] in a
10.0 M H+1 solution
 Convert all the information to Scientific
Notation and Plug the given information
into the equation.
Given [H+1] = 10.0 M
= 1.00 x 101 M
Kw = 1.0 x 10-14
Kw
[OH ]  1
[H ]
-1
-14
1.0
x
10
-15
[OH -1 ] 

1.0
x
10
M
1
1.00 x 10
29
pH & pOH
• The acidity/basicity of a solution is often expressed as
pH or pOH
• pH = -log[H3O+1]
pOH = -log[OH-1]
– pHwater = -log[10-7] = 7 = pOHwater
• [H+1] = 10-pH
[OH-1] = 10-pOH
• pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral
• The lower the pH, the more acidic the solution; The higher
the pH, the more basic the solution
• 1 pH unit corresponds to a factor of 10 difference in acidity
• pOH = 14 - pH
30
The pH scale and pH
values of some
common
substances
31
A pH meter
32
Indicator paper being
used to measure the pH of
a solution
33
Example #3
Calculate the pH of a solution with a
[OH-1] = 1.0 x 10-6 M
Find the concentration of [H+1]
Kw
[H ] 
1
[OH ]
1
-14
1.0 x 10
-8
[H ] 

1.0
x
10
M
-6
1.0 x 10
1
34
Example #3
Calculate the pH of a solution with a
[OH-1] = 1.0 x 10-6 M
Enter the [H+1] concentration into your
calculator and press the log key
log(1.0 x 10-8) = -8.0
Change the sign to get the pH
pH = -(-8.0) = 8.0
35
Example #4
Calculate the pH and pOH of a
solution with a [OH-1] = 1.0 x 10-3 M
Enter the [H+1] or [OH-1]concentration into
your calculator and press the log key
log(1.0 x 10-3) = -3.0
Change the sign to get the pH or pOH
pOH = -(-3) = 3.0
Subtract the calculated pH or pOH from
14.00 to get the other value
pH = 14.00 – 3.0 = 11.0
36
Example #5
Calculate the [OH-1] of a solution with a pH of 7.41
If you want to calculate [OH-1] use pOH, if you
want [H+1] use pH. It may be necessary to convert
one to the other using 14 = [H+1] + [OH-1]
pOH = 14.00 – 7.41 = 6.59
Enter the pH or pOH concentration into your
calculator
Change the sign of the pH or pOH
-pOH = -(6.59)
Press the button(s) on you calculator to take the
inverse log or 10x
[OH-1] = 10-6.59 = 2.6 x 10-7
37
Calculating the pH of a Strong,
Monoprotic Acid
• A strong acid will dissociate 100%
HA  H+1 + A-1
• Therefore the molarity of H+1 ions will be
the same as the molarity of the acid
• Once the H+1 molarity is determined, the pH
can be determined
pH = -log[H+1]
38
Example #6
Calculate the pH of a 0.10 M HNO3 solution
Determine the [H+1] from the acid concentration
HNO3  H+1 + NO3-1
0.10 M HNO3 = 0.10 M H+1
Enter the [H+1] concentration into your calculator
and press the log key
log(0.10) = -1.00
Change the sign to get the pH
pH = -(-1.00) = 1.00
39
Neutralization Reactions
• Acid-Base reactions are also called
Neutralization Reactions
• Often we use neutralization reactions to
determine the concentration of an unknown
acid or base
• The procedure is called a titration. With
this procedure we can add just enough acid
solution to neutralize a known volume of a
base solution
– Or visa versa
40
Buffered Solutions
• Buffered Solutions resist change in pH when an acid
or base is added to it.
• Used when need to maintain a certain pH in the
system
– Blood
• A buffer solution contains a weak acid and its
conjugate base
• Buffers work by reacting with added H+1 or OH-1 ions
so they do not accumulate and change the pH
• Buffers will only work as long as there is sufficient
weak acid and conjugate base molecules present
41