Year 11 Chemistry ~ Unit 2
Area of Study 1
Water: Essential to
Life



Water is the most abundant liquid on Earth, covering
over 70% of the planet.
The Earth’s position in the solar system means that it
gains the maximum benefit from the sun’s heat.
Water on Earth will neither boil or freeze.
Water: Essential to
Life
Figure 10.2 For the space
scientist, the presence of
gaseous or frozen water in outer
space is not an indicator of life,
but the presence of liquid water
would be considered as a
positive sign. So far, however,
none has been detected in outer
space, which makes our home
planet a unique place in the
solar neighbourhood.
The Water Cycle



The water on
Earth exist in
solid, liquid
and gas and
can readily
change from
one state to
another.
The water on
earth changes
state in order
to transport it
self around the
planet.
This movement
is called the
water cycle.
Water and Living Things Photosynthesis


Water is involved in the reactions of life:
photosynthesis and respiration.
Photosynthesis is a reaction between water
and carbon dioxide that forms glucose (energy
source of all living things).
6H 2O(l )  6CO2( g )  C6 H12O6( s )  6O2( g )
chlorophyll
Respiration


As living things require energy, the glucose will
react with oxygen in a process called respiration.
When respiration takes place, water that was
removed from photosynthesis is returned to the
water cycle.
C6 H12O6( s )  6O2( g )  6 H 2O(l )  6CO2( g )
Water and Living
Things

Water fulfils several other functions in
plants and animals.



Providing a transport system for nutrients and
soluble wastes.
It provides a system to transfer heat.
It provides a cooling system to the body.
The Properties
of Water

Water exists as a liquid over a temperature range commonly found on Earth

Hydrogen bonds between water molecules - relatively strong compared to
other intermolecular forces.
Strong hydrogen bonds give water a relatively high melting temperature
compared to other molecules of similar size.



Polar molecule - difference in electronegativity between oxygen and
hydrogen.
Polarity means it can dissolve many substances – Universal solvent (it such a
good solvent….is not found pure in nature… LIKE DISSOLVES LIKE

Oils are non-polar which means they are insoluble – float on water.

Most gases have low solubility (decrease temp increase gas solubility)
High Melting and Boiling
Temperatures
Water Expands on
Freezing




When water freezes,
the molecules form a
lattice where each
water molecule forms
hydrogen bonds with
four other molecules.
The hydrogen bonds
take up a set amount of
space.
As was melts, the water
molecules are able to
move closer together.
As such, the liquid will
occupy less space than
the solid.
High Latent Heat
Values




Latent heat measures the energy needed to change the state of a
substance at its melting and boiling temperature.
Latent heat of fusion is the energy needed to change a fixed
amount of water from a solid to a liquid at 0°C.
Latent heat of vaporisation is the energy needed to change a
fixed amount of water from a liquid to a gas at 100°C.
The higher latent heat means that more heat is needed to melt
or boil water than for equivalent amounts of other substances.
High Latent Heat
Values
High Heat Capacity




Water has the capacity to store a large
amount of heat energy.
The specific heat capacity of a substance
is the amount of heat energy needed to
raise 1 gram of the substance by 1°C.
Water has a specific heat energy of
4.2Jg-1°C-1.
When calculating the heat energy (Joules,
J) needed to heat a particular mass of a
substance, the following formula can be
used:
Heat energy = specific heat capacity x mass (g) x temperature change
= SHC x m x ΔT
The Process of
Dissolving

When one substance
dissolves in another, the
following processes
occur:



Solute particles are
separated.
Solvent particles separate
from each other.
Solvent and solute
particles attract each
other.
Dissolve, Dissociate or Ionise??

Dissolves
A substance will dissolve in water if is polar or can form ionic particles.
Polar covalent compound that can
form hydrogen bonds with water
(s) or (l) (aq)
2O
C2 H5OH (l ) H

C2 H5OH ( aq)
Ionisation
Polar covalent molecule.
So highly polarised due to
electronegativity that the covalent
bond breaks
The dipole–dipole attraction between the molecules of water and
hydrogen chloride leads to the breaking of the polar covalent
bond between the hydrogen and chlorine atoms.
Dissociation
Ionic solids in a 3D ionic lattice
Positive cations and negative anions
separate from each other
+ and – ions are more attracted to
water due to the ion-dipole
attraction so water pulls them out of
their bonds.
Not all ionic compounds are soluble
because the energy required to break
the bonds within the lattice is much
greater than the energy that the
water molecules can provide.
2O
K 2CO3( s ) H

2K(aq)  CO3( aq)
Chapter 11
Solubility Guide
Measuring Solubility

The solubility of a substance refers to the maximum
amount that can be dissolved in a given amount of
solvent at a certain temperature.

A solution in which no more solute can be dissolved is
said to be saturated.

Solubilities can be measured and compared by
determining the mass of a solute that will dissolve in
100g of solvent at a given temperature.
Solubility Curves

The relationship between
solubility and temperature
can be represented by a
solubility curve.

Each point on the curve
represents saturated
solution.

Every point below the curve
represents an unsaturated
solution.

Every point above the curve
represents a supersaturated
solution.
Crystallisation

As a solution cools, some of the solute will no longer
stay dissolved and small crystals will form.

This is known as crystallisation.

Lower temperatures will decrease the solubility of a
solute.
Solubility of Gases


Gases are much less soluble in water than most
solids but their ability to be soluble is essential for
the survival of aquatic life. Fish breathe the
dissolved oxygen in water and plants rely on
dissolved carbon dioxide.
The solubility of a gas depends on temperature
of the liquid and the pressure of the gas.


As temperature increases, the gas becomes less
soluble. This is evident by the small bubbles that
form as you heat water.
Gas becomes more soluble as the pressure is
increased.
Solubility of Gases
Temperature increase = Less soluble
Concentration of
Solutions



The concentration of a solution is the amount of
solute present compared to the amount of solution.
High amount of solute = concentrated.
Low amount of solute = dilute.
Concentration of
Solutions
Calculating
Concentration


The concentration of a solution can be
calculated by comparing the ratio of solute
compared to solvent.
The most commonly used units for
concentration are:


Mass of solute per litre of solution
Amount, in mol, of solute per litre of solution.
Mass per Litre

When calculating the concentration of a
solution in mass per litre we use:
mass _ of _ solute
concentration 
volume _ of _ solvent ( L)
Amount per Litre



Measuring solubility in moles per litre is the
most common method used by chemists.
This is also known as the molarity of a
solution.
The units used are moles per litre or molar
(M).
Amount per Litre

The molarity of a solution can be calculated using
the following formula:
n=cxV
Volume in L
Amount
in mol
Concentration
in molL-1 or M

To convert to grams per litre, multiply by the molar
mass.
Dilution



When more solvent is added to a solution,
the action is described as dilution.
Even though the volume of the solution will
change, the amount of solute remains the
same.
Dilution will change the concentration of the
solution.
Dilution

Because the number of moles of solute
remains the same for both the original
solution and the diluted solution, the
concentration can be calculated using the
following relationship:
c1V1 = c2V2
Concentration and
volume of the
original solution
Concentration and
volume of the
diluted solution
Removing Dissolved
Solutes



As mentioned earlier, water is an extremely
effective solvent.
As such, it is very difficult to find pure water
naturally.
There are a range of different ways we can
purify water.
Chapter 12
Precipitation
Reactions




Sometimes when two solutions are mixed a solid will
form and separate from the solution.
This solid is called a precipitate.
A precipitate is easily separated from a solution
through filtering.
If a known impurity is in the water, it can be
removed by adding a solution that will form a
precipitation reaction and the precipitate can then
be filtered off.
Precipitation
Reactions



When a substance dissolves in water it will
dissociate into ions.
When two solutions are mixed together the
ions from both solutions will also mix
together.
At times, some of the cations and anions will
attract each other to form an insoluble
compound.
Precipitation
Reactions
Precipitation
Reactions
Figure 12.3 Possible exchange of ions in solution.
Balanced Chemical
Equations

1.
2.
Reactions can be represented by writing
balanced chemical equations.
Identify possible products of the reaction
by swapping the cations and anions of the
reactants.
Use a valency table to write the correct
formulas of the reactants and products.
Balanced Chemical
Equations
3.
4.
Write the equation and balance it so that
equal numbers of all atoms occur on both
sides.
Use the solubility table to deduce which of
the products will form the precipitate and
place the subscript (s) next to it to denote
the state.
Ionic Equations



Because the ions in solution are dissociated, it is
more accurate to write ionic equations.
When writing ionic equations, you can see that only
some ions undergo a change and others remain the
same on both sides of the equation.
These ions are often referred to as spectator ions
and can be left out of the ionic equation.
Treatment of Drinking
Water


Because of the amount of solutes dissolved in
water and for the fact that water provides the
environment for some bacteria to live,
drinking water must go through a number of
steps in order to make it suitable for
drinking.
The first step in this process is flocculation
where solids are removed and acids are
neutralised.
Treatment of Drinking
Water




Lime (Ca(OH)2) is added to the water to
neutralise the acids.
Alum (Al2(SO4)3) is then added to form a
precipitate with the hydroxide ions from the
lime.
The gelatinous precipitate traps solids that
are suspended in the water.
The water is then left for the solids to settle
before they are removed.
Treatment of Drinking
Water





The water is then filtered through a bed of sand and
gravel to remove any other suspended particles.
The final step is chlorination.
Gaseous chlorine is passed through the water to
destroy any bacteria.
Some countries fluoridate their water by adding
fluoride ions.
Fluoridation strengthens the enamel on the teeth
and helps to fight tooth decay.
Treatment of Drinking
Water
Treatment of Drinking
Water
Desalination




Desalination is the practice of collecting sea water
and removing the salts to make it suitable for
drinking.
One method of removing the salts in water is
distillation.
Distillation involves the evaporation of water to
separate it from the salts before condensing the
steam and collecting the pure water.
Distillation is a very expensive process is not
practical to be used to desalinate water on a large
scale.
Desalination
Desalination



Osmosis is the natural tendency for water to move
from a region of low salt concentration to a region
with high salt concentration.
This process can be reversed if the salt water is
placed under high pressure.
During reverse osmosis, the salt water and fresh
water are separated by a semipermiable membrane
that allows the water through but not the dissolved
ions.
Desalination


Reverse osmosis is more suitable for large scale
desalination plants.
The only difficulty being finding a suitable
membrane that will not rupture under pressure.
Chapter 13
Introducing Acids and
Bases



Acids and the related compounds called
bases are very commonly used within our
homes.
Acid comes from the Latin ‘acere’, which
means 'sour’.
In industry, solutions of acids are used
extensively to produce a wide range of
products such as fertilisers, drugs, explosives
and plastics.
Acids
Properties of Acids



Taste sour
Have relatively low pH.
Tend to be corrosive.

Acids change litmus (a blue vegetable dye)
from blue to red.

Can conduct electricity when in solution
(Aqueous).
Bases



Acids react with bases in such a way that
they will remove each other’s properties.
As such, they are said to neutralise each
other.
Bases can be found in many household
cleaning products because they react fats and
oils to form water-soluble soaps.
Bases
Properties of Bases






Taste bitter
Feel slippery or soapy
Have a relatively high pH.
Are Caustic (Corrosive)
Bases don't change the color of litmus; but they can
turn red (acidified) litmus back to blue.
Can conduct electricity when in solution.
Safety with Acids and
Bases






Acids and Bases need to be treated with
caution:
Avoid contact with skin and eyes,
Wear safety glasses and lab coat,
Label all bottles and containers,
When diluting acids, add the acid to the
water (not the water to the acid),
Notify the teacher if a spill occurs.
Indicators



Indicators are often extracted from plant dyes and
their colour changes with the pH of a solution.
A Universal indicator is a mixture of many
indicators and changes through a range of colours
from red to green to violet.
A pH meter can be used to determine the pH of
solutions that do not change colour.
Indicators
Indicators
Indicators
Figure 13.3 The pigment
extracted from red cabbage acts
as an acid–base indicator. Here
(left to right) it has been added
to hydrochloric acid, sodium
hydroxide solution and water.
Reactions of Acids and
Bases



Acids and bases were originally grouped
together because they have similar chemical
behaviours.
When acids react they will usually form a salt
(a metal cation and non-metal anion) and
other products.
The salt will depend on what anion is formed
from the acid.
Reactions of Acids and
Bases
General Reaction Types
Involving Acids – Type 1
General Reaction Types
Involving Acids – Type 2
General Reaction Types
Involving Acids – Type 3
General Reaction Types
Involving Acids – Type 4
General Reaction Types
Involving Acids – Type 5
General Reaction Types
Involving Acids – Type 6
General Reaction Types
Involving Acids –
Summary
BrØnsted-Lowry Acids
and Bases




Johannes BrØnsted and Thomas Lowry
described the reactions of acids as involving
the donation of protons (H+)
A proton is a hydrogen atom that has lost its
electron.
A substance is an ACID if it DONATES a
proton (H+).
A substance is a BASE if it ACCEPTS a proton
(H+).
BrØnsted-Lowry Acids
and Bases
• It is a reciprocal relationship as the acid
donates a proton to the base who accepts it.
• This is why acids and bases react together.
+
H
HCl (g) + H2O (g)
Acid
Base
+
H3O
(aq)
+ Cl
--
(aq)
Acid-Base Conjugate
Pairs

If substances can be formed from each other
by donating or accepting a proton they are
said to be a conjugate acid-base pair.
HCl (g) + H2O (g)
Acid
Base
+
H3O
(aq)
Conjugate
Acid
+ Cl
--
(aq)
Conjugate
Base
Acid-Base Conjugate
Pairs
The Hydrogen Ion


The Hydrogen Ion H+ (aq) (Proton) in a
solution is also represented as H3O+ (aq).
This is called a hydronium ion and reacts in
exactly the same way as the H+ ion.
Amphiprotic
Substances


Some substances can act as acids or bases
depending on what they are reacting with.
This means they can Accept AND Donate
Protons.
Amphiprotic
Substances
Base
Acid
Acid
Base
Therefore HCO3- is Amphiprotic
Amphiprotic
Substances
Acid and Base
Strength


Different acid solutions of the same
concentrations do not have the same
strength.
The tendency of an acid or base to accept or
donate protons from water can be used as a
measure of their strength.
Strong Acids





A Strong acid very readily donates a proton.
If an acid readily donates all its protons (H+’s) this
means it completely ionises.
Acids that completely ionise in solution are called
Strong Acids.
Solutions containing strong acids would contain ions
only with virtually no unreacted acids remaining.
Strong acids are HCl, H2SO4, HNO3
Weak Acids




A Weak Acid only donates some of its
protons.
This means it only partially ionises (or
Dissociates) – or in other words, not every
molecule breaks apart.
Acids that partially ionise in solution are
called Weak Acids.
Weak acids are CH3COOH, HF
Strong and Weak Acids
Figure 14.6 (a) In a 1 M solution, hydrochloric acid is virtually
completely ionised in water.
(b) However, in a 1 M solution of ethanoic acid, only a small proportion
of ethanoic acid molecules are ionised.
Strong and Weak
Bases




Strong bases easily accept protons (or an H+).
Strong bases include OH-, HCO3Weak bases can accept protons (or an H+) but
do not do so readily.
Weak Bases include ammonia (NH3).
Polyprotic Acids





Polyprotic Acids are acids that can donate more than
one proton.
The number of protons an acid can donate depends
on how many hydrogens it has to begin with.
Monoprotic acids donate only one proton.
Diprotic acids donate two protons.
Triprotic acids donate three protons.
Strength vs
Concentration



It is important not to confuse the terms
Strong and Weak with Concentrated and
Dilute.
Strength is a measure of the ionisation of an
Acid or Base.
Concentration is how much of an Acid or Base
actually dissolves in a solution.
Strength vs
Concentration
Acidic, Basic and
Neutral Solutions


The acidity of a solution is a measure of the
concentration of Hydrogen Ions (H+) present.
The higher the concentration of H+ ions the
more acidic the solution.
Acidic, Basic and
Neutral Solutions


Pure water will ionise slightly by acting as a
very weak acid and a very weak base:
H2O(l) + H2O(l)
+
H3O
(aq)
--
+ OH (aq)
Pure water is said to be neutral because the
concentration of the H+ and the OH- ions are
both exactly the same – 10-7M
Acidic, Basic and
Neutral Solutions



Acidic Solutions contain greater
concentrations of H3O+ than OH-.
Neutral Solutions contain equal
concentrations of H3O+ and OH-.
Basic Solutions contain lower concentrations
of H3O+ than OH-.
Measuring Acidity



Experiments have shown that all aqueous
solutions (aq) contain both H3O+ and OHions.
Their product is always 10-14.
[H3O+] x [OH-] = 10-14
If either [H3O+] or [OH-] are increased, the
other must decrease proportionally.
Measuring Acidity

A solution is:
if [H3O+] > 10-7 and [OH-] < 10-7
 Neutral if [H3O+] = 10-7 and [OH-] = 10-7
 Basic if [H3O+] < 10-7 and [OH-] > 10-7
 Acidic
The pH Scale



The pH scale is a convenient way of
indicating the acidity of a solution.
The strength of an acid is based on the
[H3O+] but these values are often very small
numbers.
The pH can be calculated by:
pH = -log[H+]
The pH Scale
The pH Scale
The pH Scale
Calculations Involving
Reactions


A balanced chemical equation tells us the
proportion of reactants that must react
together to complete a chemical reaction.
The coefficients in a balanced equation give
us the ratio of amounts of reactants and
products.
Mole Ratios

For the reaction:
2H2(g) + O2(g)


2H2O(g)
2 moles of H2 react with 1 mole of O2 to form
2 moles of water.
The amount of oxygen gas used will always be
equal to half the amount of hydrogen used
and half the amount of water produced.
Stoichiometry


Stoichiometry is the term used for using
mole ratios to calculate relative amounts of
reactants used or products formed.
If a quantity of either a reactant or product is
known, the quantities of all other reactants
and products can be calculated using
stoichiometry.
Stoichiometry





The following steps outline the stoichiometric
process:
Step 1 – Write a balanced equation.
Step 2 – Calculate the amount (in mol) of the
substance with the known quantity.
Step 3 – Use the mole ratio to calculate the amount
of the required substance.
Step 4 – Calculate the quantity required.
Excess Reactants




In most situations the amounts of reactants will not
be exactly the amount required to react completely.
It is more than likely that one of the reactants will be
used up before the other.
The reaction stops when one of the reactants (the
limiting reagent) is used up even though some of
the other substance is unreacted.
The other substance is said to be in excess.
Excess Reactants


It is important to identify the limiting
reagent and the reactant that is in excess
when doing calculations on a reaction.
All calculations must be centred around the
limiting reagent.
Volumetric Analysis


The concentration of solutions of acids and
bases can be determined accurately by a
technique called volumetric analysis.
This involves reacting a solution with
unknown concentration with a solution of
accurately known concentration (a standard
solution).
Volumetric Analysis

1.
2.
A volumetric analysis involves the following
steps:
A pipette is used to dispense a known
volume (aliquot) of the solution to be
analysed into a conical flask.
An indicator is added to show when the
equivalence point of the acid-base
reaction is reached.
Volumetric Analysis
3.
4.
5.
6.
The standard solution is placed into a burette to
be titrated (reacted) with the unknown solution.
The standard solution is titrated carefully with the
unknown solution until the indicator changes
colour and equivalence point is reached.
The volume (titre) is recorded and the process is
repeated a number of times to ensure accuracy.
Stoichiometry is used to determine the unknown
concentration.
Volumetric Analysis

Figure 15.7 The equipment needed to conduct a volumetric
analysis: (a) volumetric flask; (b) pipette; (c) burette.
Oxidation and
Reduction


Many of the chemical reactions that play a
significant role in maintaining our
environment are oxidation-reduction
reactions or redox reactions.
Examples of redox reactions are the reactions
that corrode metals, combustion reactions
and photosynthesis and respiration.
Redox Reactions





Originally, oxidation was described as a reaction
with oxygen.
When oxygen reacts with a substance, the substance
is said to be oxidised.
A transfer of oxygen is referred to as redox reaction.
The substance that has gained oxygen has oxidised.
The substance that has lost oxygen has reduced.
Redox Reactions


Oxidation and reduction occur
simultaneously during a redox reaction.
Consider the redox reaction:
Oxidation –
gain of oxygen
Fe2O3(s) + 3CO(g)
2Fe(s) + 3CO2(g)
Reduction – loss
of oxygen
Electron Transfer




Not all redox reactions involve the transfer of
oxygen.
A more accurate description of a redox
reaction is the transfer of electrons.
Oxidation is the loss of electrons.
Reduction is the gain of electrons.
OIL RIG
Oxidation Is Loss
Reduction Is Gain
Half Equations
Consider the Redox reaction:
2Mg ( s )  O2 ( g )  2MgO( s )
The magnesium will undergo oxidation by losing
electrons to form magnesium ions:
Mg  Mg 2  2e 
The oxygen gas will undergo reduction by gaining
electrons to form oxide ions:

O2  4e  2O
2
Writing Overall Redox
Equations




When writing equations for redox reactions we
usually write the half equations first.
The number of electrons must be balanced so that
both half equations contain the same amount of
electrons.
All the reactants from both half equations are
written on the left hand side of the equation while
all the products are written on the right.
The electrons are cancelled out.
Copper and Silver
Nitrate
Figure 16.4 Copper wire placed in a silver nitrate
solution forms deposits of silver crystals.
Oxidants and
Reductants


An oxidant (or oxidising agent) is a
substance that causes another to be oxidised,
and is itself reduced.
A reductant (or reducing agent) is a
substance that causes another to be reduced,
and is itself oxidised.
Oxidants and
Reductants
Classifying Redox
Reactions


It is relatively easy to identify whether half
equations are either oxidation or reduction
reactions but sometimes it is not so easy to
identify full equations as redox reactions.
Redox reactions can be identified by
observing a change in Oxidation Numbers
of the substances involved.
Oxidation Number
Rules

1.
2.
3.
Oxidation numbers are determined using the
following rules:
Free elements have an oxidation number equal to
0. Eg Na(s), C(s), Cl2(g).
In ionic compounds the oxidation number s equal
to the charge on the ion. Eg CaCl2: Ca2+=+2 and Cl= -1.
Oxygen usually has an oxidation number of -2
and hydrogen has +1 but there are a few
exceptions.
Oxidation Number
Rules
4. The sum of oxidation numbers in a
neutral compound is 0 and in a
polyatomic ion is equal to the charge of
the ion.
Using Oxidation Numbers
to Identify Redox
Reactions



By defining oxidation numbers for the atoms
involved in a reaction we can look for
increases and decreases in oxidation
numbers.
An increase in oxidation number means the
element has undergone oxidation.
A decrease in oxidation number means the
element has undergone reduction.
Example
For the reaction:
+2 -2
0
+4
-2
2CO( g )  O2 ( g )  2CO2 ( g )
Carbon in the carbon monoxide has gone from +2
to +4 which means that oxidation has occurred.
Oxygen in the oxygen gas has gone from 0 to -2
which means that reduction has occurred.
As both oxidation and reduction have occurred the
reaction is a redox reaction.
Writing Half
Equations


Although most half equations are quite easy to
write, some involving polyatomic ions can be
more difficult.
The following steps will make balancing these
half equations easier:
1.
2.
3.
4.
Balance all elements except O and H in the half
equation.
Balance the O atoms by adding water.
Balance the H atoms by adding H+ ions.
Balance the charge by adding electrons (e) and then
add states.
Example
A green solution containing Fe2+ ions is mixed with
a purple solution containing MnO4- ions. Fe3+ and
Mn2+ ions are formed.
Write a balanced equation for this reaction.
The half equation involving the iron ions is quite
simple:
Fe2+(aq)
Fe3+
+ e-
Example
The half equation involving the manganese is a little more
difficult:
Step 1: Balance all elements except for O and H
MnO4Mn2+
Step 2: Balance O by adding water
MnO4Mn2+ + 4H2O
Step 3: Balance H atom by adding H+
MnO4- + 8H+
Mn2+ + 4H2O
Step 4: Balance the charge with electrons
MnO4- + 8H+ + 5eMn2+ + 4H2O
Example
To complete the full balanced equation we must
balance the electrons in each half equation:
Fe2+
Fe3+ + e(X5)
MnO4- + 8H+ + 5eMn2+ + 4H
2O
5Fe2+
5Fe3+ + 5eCombine the two equations:
MnO4- + 8H+ + 5e- + 5Fe2+
Mn2+ + 4H2O + 5Fe3+ + 5e-
Cancel out the electrons:
MnO4- + 8H+ + 5Fe2+
Mn2+ + 4H2O + 5Fe3+
Evidence for Electron
Transfer




When a redox reaction takes place, the results can
be visible but it is not always possible to see the
transfer of electrons.
To show the transfer of electrons the half reactions
must be separated and joined by an external circuit.
A galvanometer can be placed within the circuit to
measure the flow of electrons.
Such an experiment is called a galvanic cell.
Galvanic Cells
Figure 16.7 The apparatus used to demonstrate electron
flow during oxidation–reduction reactions.
Galvanic Cells





Galvanic cells consist of two half cells.
Each half cell must consist of an electrode to
conduct the electrons and an electrolyte in which
ions are free to move through the solution.
Oxidation will occur in one half cell and reduction
will occur in the other.
The electrode at which oxidation occurs is called the
anode.
The electrode at which reduction occurs is called the
cathode.
Galvanic Cells




The half cells are also connected by a salt bridge.
A salt bridge contains an ionic compound that
allows ions to flow between the solutions to
complete the circuit and prevent an over
accumulation of charge.
Anions flow into the anode to balance out the
positive charge formed from oxidation.
Cations flow into the cathode to balance out the
negative charge formed from reduction.
Galvanic Cells
The Electrochemical
Series




Chemists have constructed a table of half equations
for redox reactions that can be formed in half cells
in the order of their reactability or their ability to
Oxidise. This is called the Electrochemical Series.
The elements that reduce most readily are at the top
of the series.
The elements that are lower in the series are more
likely to oxidise.
The electrochemical series is only valid for the
conditions from which it was formed (standard
conditions).
The Electrochemical
Series
Predicting Redox
Reactions


The series can be used to predict what will
happen when two specific half cells are
combined to form a galvanic cell.
The half cell that is higher in the series will
reduce and the one lower will oxidise.
Predicting Redox
Reactions
Green Chemistry




Many aspects of our life have been enhanced by
chemistry and the chemical industry.
The production of new and improved products have
extended life expectancies and made living more
comfortable.
However, these new have sometimes come at a cost
to the environment and to human health.
Laws and processes have been put in place to
identify hazardous substances and to replace them
with safer chemicals that will do the same job.
Green Chemistry
Figure 17.2 The stepwise process that leads to the scientific
development, evaluation and possible replacement of a product
designed to fill a particular need.
Development of CFCs



Chlorofluorocarbons (CFCs) have been identified as
a group of compounds that are responsible for the
deterioration of the ozone layer.
The ozone layer acts as a filter that prevents some
UV light from reaching the Earth.
CFCs were used as a cheaper and safer alternative for
refrigeration, air conditioning and as propellants for
aerosol cans.
Development of CFCs



It was after some time that scientists
discovered that in the presence of UV light,
CFCs would react with and breakdown ozone
in the ozone layer.
As a result, the frequency of skin cancer in
some areas increased.
The use of CFCs became outlawed and the
hunt for a more environmentally friendly
alternative began.
Development of CFCs


Research determined that the chlorine atom
in CFCs as the main factor in the reaction
with ozone.
Hydrofluorocarbons (HFCs) were developed
as an alternative and were found to be just as
effective without the environmental impact.
Green Chemistry



Laws and treaties that were enacted to reduce global
pollution were aimed at dealing with wastes after
the had been produced.
Green Chemistry is a set of principles that
evaluates the environmental impact of a chemical
process.
The green approach is that the best way to minimise
hazardous waste is not to produce it in the first
place.
Principles of Green
Chemistry