Solutions
Chemistry 1-2
Mr. Chumbley
Modern Chemistry: Chapter 12
p. 378 – 402
Types of Mixtures
Chapter 12, Section 1
p. 379 – 384
Mixtures
 Some
mixtures are easy to identify, while
others are not
 Mixtures
are considered heterogeneous if
they are not uniform in composition
 Mixtures
are considered homogenous if
they are uniform in composition
Solutions
A
solution is a homogenous mixture of two
or more substances uniformly dispersed
throughout a single phase
 When
a chemical is capable of being
dissolved it is considered soluble
Components of a Solution
 The
solvent is the dissolving medium in a
solution
 The
solute is the substance dissolved in
solution
Types of Solutions
Solute
State
Solvent Example
State
gas
gas
gas
liquid
Oxygen in
nitrogen
Carbon dioxide
in water
liquid
liquid Alcohol in water
liquid
solid
solid
liquid Sugar in water
solid
solid
Mercury in silver
and tin
Copper in nickel
 Solutions
can
exist in a variety
of forms and
combinations
Special Mixtures
A
suspension is a mixture in which the
particles of the solvent are so large that
they settle out unless the mixture is
constantly stirred or agitated
A
colloid is a mixture in which the particles
are intermediate in size between those in
solutions and suspensions
Solutions and Electricity

When ionic compounds dissolve, the positive and
negative ions separate

Since the ions are free to move, electricity can pass
through the solution

An electrolyte is a substance that dissolves in water
to produce a solution that conducts electricity

A non-electrolyte is a substance that dissolves in
water to produce a solution that does not conduct
electricity
The Solution Process
Section 2
P. 385 – 394
Factors that Affect Dissolving

Several factors affect the rate at which
substances dissolve

Increasing the surface area of the solute can
increase the rate of dissolution

Agitation can increase the rate of dissolution

Heating can increase the rate of dissolution
Solubility
 Solubility
is a measure of how well one
substance dissolves in another
 Solution
Equilibrium is the physical state in
which the opposing processes of
dissolution and crystallization of a solute
occur at equal rates
Saturation

There are limits to how much solute can be
dissolved by a solvent

A saturated solution contains the maximum
amount of dissolved solute

An unsaturated solution contains less solute
than a saturated solution

A supersaturated solution contains more
dissolved solute than a saturate solution
Solubility Values
 The
solubility of a substance is the amount
of that substance required to form a
saturated solution with a specific amount
of solvent at a specified temperature
 Generally,
solubility is given as the mass of
solute dissolved by 100 g of water at
varying temperatures
Sample Problem
Potassium nitrate (KNO3) has a solubility
value of 31.6 at 20˚C. What mass of KNO3 is
needed to make a saturated solution using
50 mL of water?
Liquid Solutes and Solvents
 Just
as some solids cannot be dissolved,
some liquids will not dissolve in other
liquids either
 Immiscible
other
 Miscible
liquids are not soluble in each
liquids dissolve freely in one
another in any proportion
Concentration of Solution
Section 3
p. 396 – 402
Concentration
 The
concentration of a solution is a
measure of the amount of solute in a
given amount of solvent or solution
 Solutions
are often referred to as either
concentrated or dilute, but these are not
definite terms
Molarity
 One
way to measure the concentration
of solutions is molarity
 Molarity
(M) is the number of moles of
solute in one liter of solution
amount of solute (mol)
Molarity =
amount of solition (L)
Sample Problem 12A
You have 3.50 L of solution that contains
90.0 of sodium chloride, NaCl. What is the
molarity of the solution?
Sample Problem 12B
You have 0.8 L of a 0.5M HCl solution. How
many moles of HCl does this solution
contain?
Sample Problem 12C
To produce 40.0 g of silver chromate
(Ag2CrO4) in a reaction, you need at least
23.4 g of potassium chromate (K2CrO4) as a
reactant. If all you have is a 6.0 M K2CrO4
solution, what volume is required to have
23.4 g of K2CrO4.
Ions in Aqueous
Solutions and
Colligative
Properties
Chemistry 1-2
Mr. Chumbley
Modern Chemistry: Chapter 13
p. 410 – 419
Compounds in Aqueous
Solutions
Section 1
p. 411 – 419
Compounds in Solution
 We
have identified two types of
compounds:


Ionic Compounds
Molecular Compounds
 When
dissolved in water, ionic and
molecular compounds behave differently
Ions in Solution
 Ions
separate from each other when ionic
compounds are dissolved in water
 Dissociation
is the separation of ions that
occurs when an ionic compound
dissolves
Dissociation
 We
can use a chemical equation to
indicate dissociation of ions in solution
NaCl 𝑠
CaCl2 𝑠
H2 O
H2 O
Na+ (𝑎𝑞) + Cl− (𝑎𝑞)
Ca2+ (𝑎𝑞) + 2Cl− (𝑎𝑞)
Sample Problem 13A
Write the equation for the dissolution of
aluminum sulfate, Al2(SO4)3, in water.
A. How many moles of aluminum ions and
sulfate ions are produced by dissolving 1
mol of aluminum sulfate?
B. What is the total number of moles of ions
produced by dissolving 1 mol of
aluminum sulfate?
Molecular Compounds in
Solution
A
molecular compound ionizes in a polar
solvent
 Ionization
occurs when ions are formed
from solute molecules by the action of the
solvent
 Ionization
is different from dissociation
Hydronium Ion

Many molecular compounds have a
hydrogen atom bonded by a polar covalent
bond

The ionization of these compound is enough
to transfer the H+ ion to the water molecule
making it H3O+

The H3O+ ion is known as the hydronium ion
Colligative Properties of
Solutions
Section 2
p. 422 - 432
Colligative Properties
 The
amount of solute sometimes affects
the properties of solutions
 Colligative
properties are properties that
depend on the concentration of solute
particles but not on their identity
Nonelectrolytes in Solution

Nonelectrolytes affect how a solvent undergoes phase changes

A solution with a nonelectrolyte solute will have a lower freezing point than
the pure solvent

Freezing point depression is the difference between the freezing points of the
pure solvent and a solution of nonelectrolyte in the solvent

A solution with a nonelectrolyte solute will have a higher boiling point than a
pure solvent

The concentration of nonelectrolytes lowers the vapor pressure of the solvent

Boiling point elevation is the difference between the boiling points of the pure
solvent and a nonelectrolyte solution of that solvent

Both freezing point depression and boiling point elevation are directly
proportional to the concentration of the solution
Electrolytes in Solution
 Electrolytes
in solution also cause freezing
point depression and boiling point
elevation
 Electrolytes
are much more effective at
altering properties than nonelectrolytes
Acids and
Bases
Chemistry 1-2
Mr. Chumbley
Modern Chemistry: Chapter 14
p. 440 – 469
Properties of Acids and
Bases
Section 1
p. 441 – 450
Common Acids and Bases
Properties of Acids

Acids are defined by their properties:
1.
Aqueous solutions of acids have a sour taste
2.
Acids change the color of acid-base indicators
3.
Some acids react with active metals to release
hydrogen gas
4.
Acids react with bases to produce salts and water
5.
Acids conduct electric current
Naming Acids
 There
are two types of acids:

Binary acids are composed of only two
different elements: hydrogen and a more
electronegative element

Oxyacids are composed of hydrogen,
oxygen, and a third element, usually a
nonmetal
Naming Binary Acids

Binary acid nomenclature
follows the following rules:
1.
The name of the binary
acid begins with the
prefix hydro-
2.
3.
The root of the name of
the second element
follows the prefix
The name ends with the
suffix –ic
Formula
Acid Name
Molecule
Name
HF
hydrofluoric
acid
hydrogen
fluoride
HCl
hydrochloric
acid
hydrogen
chloride
HBr
hydrobromic
acid
hydrogen
bromide
HI
hydriodic
acid
hydrogen
iodide
H2S
hydrosulfuric
acid
hydrogen
sulfide
Naming Oxyacids


Oxyacids only form from
polyatomic ions and
hydrogen
To name an oxyacid,
simply take the name of
the polyatomic ion and
change the suffix
 -ate changes to -ic
 -ite changes to -ous
Formula
Acid Name
Anion
CH3COOH
acetic acid
acetone
(CH3COO-)
HNO2
nitrous acid
nitrite (NO2-)
HNO3
nitric acid
nitrate (NO3-)
HClO
hypochlorous
acid
hypochlorite
(ClO-)
HClO4
perchloric
acid
perchlorate
(ClO4-)
H2CO3
carbonic
acid
carbonate
(CO32-)
H3PO3
phosphorous
acid
phosphite
(PO33-)
Common Acids used in
Industry

Sulfuric acid (H2SO4)





Used to produce many
material goods
found in batteries
Used as a desiccant
Hydrochloric acid (HCl)



Nitric acid (HNO3)


Used in explosives
Used in production of rubbers,
plastics, and pharmaceuticals

Acetic acid (CH3COOH or
HC2H3O2)


Phosphoric acid (H3PO4)


Most commonly found in
fertilizers
Flavor additive to beverages
Stomach acid
Used in metal production,
food processing, oil wells
Sold commercially as muriatic
acid to maintain swimming
pools and masonry


Vinegar contain acetic acid
Used to manufacture plastics,
and to produce food
supplements
Used as a fungicide
Properties of Bases

Bases differ from acids
1.
Aqueous solutions of bases taste bitter
2.
Bases change the color of acid-base indicators
3.
Dilute aqueous solutions of bases feel slippery
4.
Bases react with acids to produce salts and water
5.
Bases conduct electric current
Arrhenius Acids and Bases

When certain acids and bases form in
solution, they produce ions

Arrhenius acids increase the concentration of
hydrogen ions, H+, in aqueous solution

Arrhenius bases increase the concentration of
hydroxide ions, OH-, in aqueous solution
Acid-Base Theories
Section 2
p. 452 – 457
Brønsted-Lowry

The Arrhenius acid-base definition was expanded by Johannes
Brønsted and Thomas Lowry

They redefined acids and bases as proton donors or acceptors

A Brønsted-Lowry acid is a molecule or ion that is a proton donor

A Brønsted-Lowry base is a molecule or ion that is a proton
acceptor

In a Brønsted-Lowry acid-base reaction, protons are transferred
fromone reactant (the acid) to another (the base)
Number of Protons

Acids can be classified by the number of
protons they donate

Monoprotic acids donate only one proton per
molecule

Polyprotic acids donate multiple protons per
molecule


Diprotic acids donate two protons
Triprotic acids donate three protons
Acid-Base Reactions
Section 3
p. 457 – 463
Brønsted-Lowry Reactions

Brønsted-Lowry involve conjugate acid-base
pairs

A conjugate base is the substance that
remains after a Brønsted-Lowry acid has given
up a proton

A conjugate acid is the substance that
remains after a Brønsted-Lowry base gains a
proton
Neutralization Reactions

When acids and bases react with each other,
they produce water and a salt

In aqueous solutions, neutralization is the
reaction of hydronium ions and hydroxide ions
to form water molecules

A salt is an ionic compound composed of a
cation from a base and an anion from an
acid
Acid-Base
Titration and pH
Chemistry 1-2
Mr. Chumbley
Modern Chemistry: Chapter 15
p. 470 – 499
To neutralize 17.3 mL of an unknown solution
of sulfuric acid, 25.0 mL of o.oo15 M
calcium hydroxide is added. What is the
concentration of the hydrochloric acid?
In a titration, 27.4 mL of 0.0154 M Ba(OH)2 is
added to a 20.0 mL sample of HCl solution
of unknown concentration until the
equivalent point is reached. What is the
concentration of the acid solution?
Self-Ionization of Water
 In
the self-ionization of water, two water
molecules produce a hydronium ion and
a hydroxide ion by transfer of a proton
 The
molar concentration of both
hydronium and hydroxide ions is 1.0 × 10-7 M
 This
value is known as the ionization
constant of water, Kw
pH and pOH

The concentrations of hydronium and
hydroxide ions determine pH and pOH

pH is defined as the negative common
logarithm of the hydronium ion concentration

pOH is defined as the negative common
logarithm of the hydroxide ion concentration

The sum of a solution’s pH and pOH is always
14
The pH Scale

The relative strength of acids and bases is
shown using the pH scale

Strong acids have a very high concentration
of hydronium ions and have a low pH value

Strong bases have a very low concentration
of hydronium ions and have a very high pH
value