Download Int2 Unit Three Summary

INTERMEDIATE 2 CHEMISTRY
SUMMARY NOTES
Unit Three: Acids, Bases & Metals
1. Acids & Alkalis
pH Scale
pH is measured using Universal indicator, pH paper or a pH
meter.
The scale can go from -1 to 15 but most substances lie between
1 and 14.
Acids have a pH < 7
Alkalis have a pH > 7
Neutral substances have a pH = 7.
The pH scale is dependent on the concentration of H+(aq) ions
and is based on the fact that water is partially ionised.
H2O(l)
H+(aq) + OH-(aq)
Acids and alkalis affect the pH when added to water, so
change the H+ ion concentration.
In water:
Concentration of H+ = Concentration of OHIn acids:
Concentration of H+ > Concentration of OHIn alkalis:
Concentration of H+ < Concentration of OHSoluble metal oxide form alkalis.
Soluble non-metal oxides form acids.
All fizzy drinks are acidic because of the carbon dioxide gas
dissolved in them.
Most cleaning compounds tend to be alkaline.
Strong acids are those which are completely dissociated into
ions (ie. Every molecule is broken into ions).
Weak acids are those which are only partially dissociated into
ions.
All alkanoic acids (eg. Ethanoic) are weak acids. Hydrochloric,
sulphuric and nitric acids are all strong.
Strong alkalis are those which are completely dissociated into
ions (ie. Every molecule is broken into ions).
Weak alkalis are those which are only partially dissociated into
ions.
Ammonia solution (ammonium hydroxide) is a weak alkali.
Ammonia is also very water-soluble and has a pungent smell
(wet nappies). Sodium, potassium and calcium hydroxides are all
strong.
Strong vs Weak
Strong acids have a higher ion concentration than a weak acid
of the same concentration. This means they:
 React faster
 Have a higher electrical conductivity
 Have a lower pH
However, they are neutralised by the same amount of a base.
Effect of Dilution
Diluting an acid increases its pH towards 7.
Diluting an alkali decreases its pH towards 7.
Diluting both reduces their electrical conductivity.
Concentration and The Mole
For solutions:
N = cV
where
N = number of moles (mol)
c = concentration in moles per litre
(mol l-1)
V = volume in litres (l)
2. Salt Preparation
Reactions of Acids
Neutralisation reaction: moves the pH of an acid to 7,
producing water.
Base: a substance which neutralises an acid.
Alkali: a soluble base.
Salt: ionic compound formed from the negative ion of the acid
and the positive ion from the neutraliser (usually metal).
Acid
Type of salt
Hydrochloric
… chloride
Nitric
… nitrate
Sulphuric
… sulphate
Ethanoic
… ethanoate
1. acid + alkali (metal hydroxide) →
salt + water
2. acid + metal oxide
→
salt + water
3. acid + metal carbonate → salt + water + carbon dioxide
4. acid + metal (MAZINTL) →
salt + hydrogen
1 → 3 are neutralisations, while 4 is a displacement.
Spectator ions: Ions which do not change during the reaction
so don’t actually take part in the reaction.
Acid rain: caused by sulphur dioxide, nitrogen dioxide and
carbon dioxide. Affect buildings, structures and wildlife.
Farmers use lime to neutralise the effects on soils and lakes.
Fertilisers: water-soluble compounds containing one of the 3
essential elements needed for healthy plants – nitrogen,
phosphorus and potassium (NPK). Can be made by neutralisation
of acids by ammonia.
Precipitation: formation of an insoluble salt from 2 solutions.
How you know a precipitation reaction will take place:
1. Use the solubility data on page 5 of the databook.
2. Take the names of the 2 soluble chemicals you start with:
For example: lead nitrate + potassium iodide
3. Swap the 2 positive ions to make 2 new substances:
That is: lead iodide + potassium nitrate
4. If one of these compounds is insoluble, then a
precipitation reaction will occur!
How to spot a precipitation reaction equation:
1. Check the state symbols.
2. The 2 reactants must both be (aq) and one product must
be (s).
Example: Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq)
Titration Calculations
At the end point (neutralisation) of a titration:
Number of H+ ions = Number of OH- ions
So the calculation is:
pcV (acid) = pcV (alkali)
where
p(acid) = number of H ions in formula
p(alkali) = number of OH ions in formula
c = concentration and V = volume
3. Metals
A cell is the simplest form of battery and consists of 2
different metals joined in a circuit, separated by an
electrolyte.
An electrolyte is used to complete the circuit and is usually a
solution of an ionic compound. This allows charge to flow.
The voltage produced in a cell varies as follows:
1. The further apart the 2 metals are in the Electrochemical
Series, the higher the voltage.
2. The higher the concentration of the electrolyte, the higher
the voltage.
Electrons flow through the wires of the circuit from the metal
higher up the Electrochemical Series to the one lower down,
eg. Zn → Cu.
Displacement Reactions
These occur if the solid metal is higher up the
Electrochemical Series than the metal in the salt solution. The
metals will then swap places. The metal higher up would rather
be in the ionic form so gives its electrons to the metal in
solution.
Oxidation and Reduction
OILRIG
Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)
A Redox reaction is one where reduction and oxidation occur at
the same time, eg. Displacements, electrolysis and making
electricity in a cell.
Redox Equations
The number of electrons in the 2 ion-electron equations must
be the same, so that they cancel each other out.
1
2
Zn(s)
+
Ag (aq) + e-
→ Zn2+(aq) + 2e→ Ag(s)
Equation 2 has only one electron moving so we must multiply
this by 2 before we can add the 2 equations together to get
the redox equation.
1
2x2
add
Zn(s)
2Ag+(aq) + 2eZn(s) + 2Ag+(aq)
→ Zn2+(aq) + 2e→ 2Ag(s)
→ Zn2+(aq) + 2Ag(s)
Reactions of Metals
With Oxygen
With Water
Magnesium
Potassium
Zinc
Sodium
Iron
Calcium
Copper
Magnesium (when hot)
With Acid
Magnesium
Aluminium
Zinc
Iron
Nickel
Tin
Lead
Extraction of Metals
Metals are extracted from their ores. This is a reduction
process.
Metal
Method of Extraction
potassium
Electrolysis
aluminium
zinc
copper
silver
Heating with
carbon or
carbon monoxide
Heat
alone
Extraction of iron is done in a blast furnace. As well as the
iron ore, carbon is added (turned into carbon monoxide gas for
more efficient extraction) and limestone removes impurities.
Corrosion
This is a reaction where the surface of a metal is turned into a
compound. This is an oxidation reaction of the metal.
Rusting is the corrosion of iron.
This process can be investigated using ferroxyl indicator. This
turns blue in the presence of Fe2+ ions (the first stage of
rusting)
2 things are needed for rusting to occur:
1. Oxygen (from air)
2. Water
Rusting is speeded up by:
1. Salt
2. Acid
3. Attaching a metal lower in the Electrochemical Series
Preventing Corrosion:
1. Physical protection (paint, plastic, metal stops air and
water getting to the metal)
2. Sacrificial protection (a metal higher in the
Electrochemical Series is attached)
3. Direct Electrical Protection (attached to a negative
terminal supplying electrons)
In sacrificial protection, the metal higher up gives its
electrons up to the iron. It is corroded to save the iron.
Iron can be covered by another metal by electroplating. It is
attached to the negative terminal and dipped into a solution
containing ions of the other metal.
Galvanising is the covering of iron by zinc (can be done by
electroplating or just dipping in liquid zinc). This offers BOTH
physical and sacrificial protection as zinc is higher in the
Electrochemical Series.
On the other hand, tin-plating offers only physical protection
and if cracked would then speed up rusting as tin is below iron.