Chapter 7 Reactions and Solutions Copyright

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Chapter 7
Reactions and
Solutions
7.1 Writing Chemical Reactions
• We will learn to identify the following patterns of
chemical reactions:
–
–
–
–
combination
decomposition
single replacement
double replacement
• Recognizing the pattern will help you write and
understand reactions.
Combination Reactions
• The joining of two or more elements or compounds.
A + B  AB
• Example: MgO(s) + CO2(g)  MgCO3(s)
Decomposition Reactions
• produce two or more products from a single reactant
AB  A + B
• Example: CaCO3(s)  CaO(s) + CO2(g)
Replacement Reactions (2 types)
1. Single replacement
A + BC  B + AC
• Example:
Cu(s)+2AgNO3(aq)  2Ag(s)+Cu(NO3)2(aq)
2. Double Replacement
AB + CD  AD + CB
• Example:
HCl(aq)+NaOH(aq) NaCl(aq)+H2O(l)
7.2 Types of Chemical Reactions
Precipitation Reactions
• Chemical change in a solution that results in one or more 2
insoluble products.
• To predict if a precipitation reaction can occur it is helpful
to know the solubilitie of ionic compounds (Table 7.1).
Predicting Whether Precipitation Will Occur
• Have the ionic compounds exchange partners.
• Look at the new compounds found and determine if
any are insoluble according to the rules (Table 7.1)
• The insoluble salt will be the precipitate.
• Example: Pb(NO3)2(aq) + NaCl(aq)  PbCl2(s) + NaNO3(aq)
Reactions with Oxygen
• Reactions with oxygen generally release energy.
• Example: combustion of natural gas
• Organic compounds (carbon containing cp.) CO2
and H2O usually the products.
• Example: CH4+2O2CO2+2H2O
• Rusting (corrosion of iron) is an inorganic
example.
4Fe + 3O2  2Fe2O3
Acid-Base Reactions
• These reactions involve the transfer of a
hydrogen ion (H+) from one reactant to another.
• Example:
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
The hydrogen ion on HCl was
transferred to the oxygen in
OH-, giving H2O
Oxidation-Reduction Reactions
• Reaction involves the transfer of one or more
electrons from one reactant to another.
• Example:
Zn(s) + Cu2+(aq) Cu(s) + Zn2+(aq)
Two electrons are
transferred from Zn to
Cu2+
7.3 Properties of Solutions
• Solution - homogeneous mixture
• Solute – that which is dissolved, the substance
present in the lesser quantity
• Solvent – that which does the dissolving, the
substance present in the largest quantity
• Aqueous Solution - solution where the solvent is
water
• Solutions can be liquids as well as solids and gases.
General Properties of Liquid Solutions
• Clear, though may have color
• Remember that electrolytes are compounds that
dissociate when they dissolve in water.
NaCl( s) 
 Na (aq)  Cl (aq)
H 2O

• Nonelectrolytes do not dissociate.
• Volumes are not additive.
–1 L of ethanol + 1 L water does not give 2 L of solution.
• Solutions vs. Colloids vs. Suspensions
4
• Colloidal suspension - contains solute particles which
are not uniformly distributed.
–
Due to larger size of particles (1nm - 200 nm)
–
Smaller than 1 nm, have solution.
–
Larger than 200 nm, have a precipitate.
• Tyndall effect - the ability of a colloidal suspension to
scatter light.
–
See this as a haze when shining light through the mixture.
–
Solutions: light passes right through without scattering.
• Suspension:
–
mixture contains particles much larger than a colloidal
suspension
–
particles easily settle out and form a precipitate
Degree of Solubility
• Solubility - how much solute can dissolve in a solvent.
• Factors which affect solubility:
1 Polarity of solute and solvent
• The more they are different, the less soluble.
2 Temperature
• increase in temp. usually increases solubility.
3 Pressure
• for gas solutes, increased pressure of gas
increases solubility
• Unsaturated solution - does not contain all the solute
that can be dissolved at a particular temperature.
• Saturated solution - contains all the solute that can be
dissolved at a particular temperature.
• Supersaturated solution - contains more solute than
can be dissolved at that temperature.
• How is this done?
•
Heat solvent, saturate it with solute then cool slowly.
•
Sometimes the excess will precipitate out.
•
If it doesn’t precipitate, the solution will be supersaturated.
Solubility of Gases: Henry’s Law
• Henry’s Law - the number of moles of a gas dissolved
in a liquid is proportional to the partial pressure of the
gas above the liquid. (At constant temperature)
• Solubility and Equilibrium – a saturated solution is an
example of a dynamic equilibrium.
7.4 Concentration of Solutions: Percentage
• Concentration - amount of solute dissolved in a given
amount of solution.
amount of solute
concentrat ion 
amount of solution
• We will learn to express concentration in 3 ways:
Weight/Volume Percent, Weight/Weight Percent and Molarity
6
Weight/Volume Percent
W
grams of solute
%

100%
V milliliter s of solution
Weight/Weight Percent
W
grams solute
%

100%
W
grams solutions
Used most often with solid solutions.
Note: there is an error with the above formula on
page 186 of your book.
7.5 Concentration of Solutions: Moles and
Equivalents
Molarity:
moles solute
M
L solution
• A dilution is required if..
–
you have a solution in stock and wish to make a
solution of less concentration.
–
Simply add the correct amount of water.
moles solute
M

• Since:
L solution
• Then: M x V = moles of solute
7
• In a dilution will the moles of solute change? No.
• So, since M x V=moles
M1V1 = M2V2
• You may use %W/V
concentration units as
well
This is the dilution formula, DO NOT USE IT FOR
ANYTHING ELSE!
Representation of Concentration of Ions in
Solution
•
Two common ways of expressing concentration of
ions in solution:
1. Moles per liter (molarity)
2. Equivalents per liter (emphasizes charge) We
will not cover this.
•
Let’s look at Na3PO4.
•
If we dissolved 1 mol of sodium phosphate into 1 L
of water, what would be the concentration (M) of the
solution? 1 M
9
• Now let’s consider the individual ions. What would the
concentration of Na+ be? 3 M
• What would the concentration of PO43- ions be? 1 M
• Equivalent is defined by the charge. One way of
looking at it: it is the mole of charges:
• 1 mol Na+ = 1 equivalent Na+
• 1 mol PO43- = 3 equivalents PO43• Okay, back to the 1 M Na3PO4 solution. What are the
equivalents of Na+?
• 3 mol Na+ = 3 equivalents of Na+.
• What are the equivalents of PO43-?
• 1 mol PO43- = 3 equivalents of PO43-.
• Do you see where the word equivalent might come
form?
• So the following are things we can say about a 1 M
solution of Na3PO4.
3 mol/L Na+
1 mol/L PO433 equivalents/L Na+
3 equivalents/L PO43• One Equivalent of an ion is the number of grams
of the ion corresponding to Avogadro’s number of
electrical charges.
molar mass of ion (g)
One equivalent of an ion 
number of charges on ion
7.6 Concentration-Dependent Solution
Properties
• Colligative properties- properties of solutions that
depend on the concentration of the solute particles,
rather than the identity of the solute.
• There are four:
–
–
–
–
vapor pressure lowering
boiling point elevation
freezing point depression
osmotic pressure
Vapor Pressure Lowering
• Raoult’s Law - when a nonvolatile solute is
added to a solvent, the vapor pressure of the
solvent decreases in proportion to the
concentration of the solute.
Freezing Point Depression and
Boiling Point Elevation
• Before we can discuss the above colligative properties
we must learn a new concentration unit.
• Molality (m) = moles of solute per kg of solvent.
moles solute
Molality 
kg solvent
• Note: the denominator is in kg solvent, not in kg
solution.
Freezing point depression (DTf) is proportional to the
number of solute particles
DTf=kf m
freezing point depression constant
Each solvent has it’s own value.
• Note, it is solute particles not just solute.
• If the solute is an electrolyte, what will it do? It will
break apart into ions.
• Therefore the same concentration of NaCl will affect the
freezing point twice as much as glucose (a
nonelectrolyte)
Boiling point elevation (DTb) is proportional to the
number of solute particles
DTb=kb m
boiling point elevation constant
Each solvent has it’s own value.
• Again, an electrolyte will affect the boiling point to a
greater degree than a nonelectrolyte with the same
concentration.
Osmotic Pressure
• Semipermeable membranes- allow
solvent but not solute to diffuse from one
side to another.
• Osmosis - the movement of solvent from
a dilute solution to a more concentrated
solution through the membrane.
• Osmotic pressure (Π) - the amount of
pressure required to stop the flow.
• Again, M is molarity of particles. Called
osmolarity.
P=MRT
Hypertonic solution = a solution has a greater
concentration of solutes than another across a
semipermeable membrane
Hypotonic solution = a solution has a lesser
concentration of solutes than another across a
membrane
Isotonic solution = both solutions have the same
concentration of solutes and are in equilibrium
7.7 Water as a Solvent
• Water often referred to as the “universal solvent”.
• Most abundant liquid on earth.
• ~ 80% of the human body is water.
– transports ions, nutrients and waste into and out of cells.
– solvent for biochemical reactions in cells and digestive tract
– reactant or product in some biochemical processes.
• Active transport: Cellular energy must be expended to make
concentration of ions different on each side of the cell membrane.
– This is accomplished via large protein molecules embedded in cell
membranes.
CATIONS IN THE BLOOD and CELLS
• Na+ and K+ two most important cations.
• Danger to body occurs when Na+ and K+ in blood
and in cells becomes too high or low.
- Na+ too low:
- Decrease of urine output
- Drymouth, flushed skin, fever
- Na+ too high:
- Confusion, stupor or coma
- K+ too high or too low:
- Death by heart failure
ANIONS IN THE BLOOD
• Cl- acid/base balance
- maintenance of osmotic pressure
- oxygen transport by hemoglobin
• HCO3-
- Form in which most waste CO2 is carried
out of the body.
PROTEINS IN THE BLOOD
• Blood clotting factors
• antibodies
• albumins (carriers of non-polar substances which
cannot dissolve in water)
• Proteins are transported as a colloidal suspension.
• The blood also transports nutrients and waste
products.