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C4 - Chemical Changes and Electrolysis - Coggle Diagram
C4 - Chemical Changes and Electrolysis
L1 - L3 - Oxidation and Reduction
Metals react with oxygen
to produce
metal oxides
; the production of a metal oxide is an
oxidation reaction
(in terms of the metal involved) because the
metal loses electrons
and gains oxygen.
The
higher a metal is on the reactivity series
, the
more reactive it is
and the
more substances it can displace in their compound.
For example,
potassium is very reactive
and can displace any substance under it in the reactivity series. A
more reactive metal forms a positive ion more easily
than a less reactive metal.
Unreactive metals
are found in the Earth in their
pure, unreacted form
(such as gold),
more reactive metals
will often be
found as metal oxides(ores)
. Metals that are
less reactive than carbon
can be
extracted from their ores using carbon reduction
(loss of oxygen,
gain of electrons
).
A reaction that involves
both oxidation and reduction
is known as a
redox reaction
.
L4- Reaction of Metals With Acid and Water
Metal + Acid ---> Salt + Hydrogen
When writing an
ionic equation
, break down
compounds into their constituent elements
, express state symbols and charges. When forming a
net ionic equation, spectator ions should be removed
.
L5 - Neutralisation and the pH Scale
The
pH of a solution
can be tested through the use of a
pH probe
, a
pH meter
or with indicators such as
universal indicator
,
litmus paper
,
methyl orang
e or
phenolpthalein
.
Acid + Alkali ---> Salt + Water
All
acids
contain
hydrogen
(released as a H+ ion), all
alkalis
contain
hydroxide ions
(OH-).
The pH scale spans from 0 to 14, with values
lower than 7 being acidic
and those
larger than 7 being alkali
; water has a n
eutral pH (7).
L6 - Strong and Weak Acids
When
acids are added to a solution
, their
ions disassociate(separate)
, this is what
makes the solution acidic
.
Strong acids completely ionise
in solutions; the
greater the tendency for the acid to ionise, the lower the pH
.
Weak acids only partially ionise in solutions
; these reactions are often
reversible
.
pH
is a
measure of the concentration of hydrogen ions
in a solution; the
more hydrogen ions there is
, the
more acidic the solution is
. As
pH decreases by one unit, the hydrogen ion concentration increases by a factor of 10.
L7 - Reactions of Acids with Metal Carbonates
Acid + Metal Carbonate ---> Salt + Water + Carbon Dioxide
Acid + Metal Hydroxide ---> Salt + Water
A soluble base
(e.g sodium hydroxide) is
known as an alkali
, an
insoluble metal hydroxide
(copper hydroxide) is referred to as a
base only.
Acid + Metal Oxide ---> Salt + Water
Bases
are
chemicals that can neutralise acids.
The test for carbon dioxide is the
cloudy limewater test
; carbon dioxide will turn
limewater cloudy.
L8 - Electrolysis Part 1
Electrolysis
is a
process used to separate a compound
using
electrically charged
carbon
electrodes
.
An
electrolyte is a fluid
that can
conduct electricity.
This can be a
liquid, solution or even a gas
(under certain pressures).
An
anode
is a
positively charged electrode
(attracts negatively charged ions/anions).
A cathode is a negatively charged electrode
(attracts positively charged ions/cations).
General method:
Melt or dissolve an ionic compound
to
allow charge to flow
.
-
Place
two, carbon electrodes
on
either side of the vat.
-
Begin to
run a current through the electrodes.
-
The
electrodes should attract their complementary ions.
With
molten ionic compounds
,
oxidation takes place at the anode
;
reduction takes place at the cathode
. This is because at the anode, negatively charged ions lose electrons and at the cathode, positively charged ions gain electrons.
During the
electrolysis of a molten ionic compound
, a
solid or molten metal
will
form at the cathode
. At the
anode, a halide gas can be produced.
L9 - Electrolysis Part 2
Cryolite is often used to decrease the melting point of a substance
, this henceforth
reduces the energy required
and
therefore the cost.
During the
electrolysis of aluminium oxide
,
solid aluminium is formed at the negative electrode
.
Oxygen
is separated from the oxide and is attracted to the positive electrode, here it
reacts with the carbon electrode and produces carbon dioxide
. This means that the
electrode has to be replaced regularly
.
During the
electrolysis of aluminium oxide(oxidation)
, the
oxygen loses electrons
and the
aluminium gains electrons(reduction)
.
Aluminium
is a
lightweight, non-corrosive metal
, it is a
good thermal and electrical conductor
and is
found naturally as bauxite
(which is mainly aluminium oxide).
L10 - General Rules of Electrolysis
[
Of aqueous solutions
] :
Oxygen forms at the positive electrode
unless there are halide ions present.
[
Of aqueous solutions]
:
Diatomic hydrogen is formed
unless the metal involved is less reactive than hydrogen
.
C5 - Energy Changes
L1 - Exothermic and Endothermic Reactions
During an exothermic reaction, a system loses energy to the surroundings, heating the surroundings and cooling the system. When this is graphed, a negative energy change for the system should be seen. Examples of exothermic reactions include oxidation, neutralisation and combustion.
During an endothermic reaction, a system gain energy; cooling the surroundings. This cools the surroundings and heat the system. A positive energy change for the system should be observed. Examples of endothermic reactions are photosynthesis and self cooling ice packs.
On a reaction profile, a value for the reactants energy being greater than the value for the products energy means the reaction is exothermic.
Activation energy is described as the minimal amount of energy needed for colliding particles to begin a reaction.
Required Practical:
Use a 50cm^3 measuring cylinder to obtain 30cm^3 of a dilute hydrochloric acid, pour this acid into a polystyrene cup.
-
Stand the cup inside a beaker(increases stability).
-
Use a thermometer to measure the temperature of the acid, record this.
-
Measure and pour 5cm^3 of a sodium hydroxide solution into the cup.
-
Put a lid on the cup and gently stir the solution; when the temperature stops changing, record it.
-
Add an extra 5cm^3 of sodium hydroxide solution to the cup, lifting the lid for a minimum period of time, when the temp no longer changes, record the value.
-
Continue the previous step until 40cm^3 of sodium hydroxide has been added in total.
-
Calculate the mean maximum value for the temperature reached.
L2 - Energy Change Calculations
Energy Change = Bonds Broken - Bonds Made
If the energy needed to break the bonds is greater than the energy transferred to the surroundings when the products are made, the reaction will be endothermic
L3 - Cells and Batteries
A simple cell can be constructed with two metals of differing reactivity in contact with an electrolyte. The greater the difference in reactivity between these metals, the greater the voltage produced.
In non-rechargeable batteries and cells; the chemical reactions stop when one of the reactants has been used up. Alkaline batteries are non-rechargeable.
In rechargeable batteries, the chemical reaction within can be reversed if an external current is applied.
L4 - Fuel Cells
Hydrogen is being developed as a fuel because it produces no pollutants and burns well.
As fuel (hydrogen) enters the cell, it becomes oxidised by the negative electrode. The electrons which are lost are passed through a wire to the opposite electrode. The hydrogen atoms then more through the electrolyte to the positive cathode. Here they react with oxygen (oxygen is gaining electrons / reduction) to form water.
At the negative electrode: 2H2 + 4OH ---> 4H2O + 4e-
-
At the positive electrode: O2 + 2H2O + 4e- ---> 4OH
-
Overall: 2H2 + O2 ---> 2H2O
Pro's
: Only water vapour produced, relatively simple devices (last longer than batteries), only require hydrogen and oxygen.
Con's
: Requires more space to store, hydrogen is also explosive, to make the hydrogen fuel fossil fuels are often used to produce the required energy