Dr Mo Int 2 Chemistry Unit 3 Acids, Bases + Metals Flashcards

Terms Definitions
pH below 7
Acidic
pH above 7
Alkaline
pH = 7
Neutral
pH pure water
Neutral pH = 7
Soluble non-metal oxides
acidic
Soluble metal oxides
alkaline
Soluble hydroxides
alkaline
Ammonia dissloves in water to produce
an alkali
Vinegar
Weak acid
Bleach
Strong alkali
Sodium hydroxide
strong alkali
Ammonium hydroxide
Weak alkali
Sulphuric acid
Strong acid
Hydrochloric acid
Strong acid
Acidic solution
contains more hydrogen ions than hydroxide ions.
Alkaline solution
contains more hydroxide ions than hydrogen ions.
In water and neutral solutions
concentration of hydrogen ions EQUALS concentration of hydroxide ions.
Diluting an acid:
pH rises towards 7
concentration of hydrogen ions decreases
concentration of hydroxide ions increases
Diluting an alkali
pH falls towards 7
concentration of hydroxide ions decreases
concentration of hydrogen ions increases
Water equilibrium:
Water molecules dissociate to form hydrogen ions and hydroxide ions. This is reversible.
When a reversible reaction is in equilibrium
concentrations of reactants and products are constant, but may be different from each other.
Moles (concentration)
moles = concentration x volume
volume must be in litres
In aqueous solution
strong acids
completely dissociate into ions
In aqueous solution
weak acids
partially dissociate into ions
In aqueous solution
strong bases
completely dissociate into ions
In aqueous solution
weak bases
partially dissociate into ions
Nitric acid
Strong acid
Ethanoic acid
Weak acid
Compared to a weak acid, a strong acid (of equal concentration) has
lower pH
higher conductivity
faster rate of reaction
Compared to a weak base, a strong base (of equal concentration) has
higher pH
higher conductivity
faster rate of reaction
Neutralisation
reaction of acids with bases
salt and water are always produced
Titration equation
Pa Ca Va = Pb Cb Vb
P = power - No. hydrogen or hyroxide ions
C = concentration, V = volume
a = acid, b = base
Examples of bases:
metal oxide
metal hydroxides
metal carbonates
Acid + metal hydroxide →
salt + water
hydrogen ions and hydroxide ions react to form water
Acid + metal oxide →
salt + water
hydrogen ions and oxide ions react to form water
Acid + metal carbonate →
salt + water + carbon dioxide
hydrogen ions and carbonate ions react to form water and carbon dioxide
Water soluble bases form
Alkalis
Everyday examples of neutralisation
treatment of acid indigestion
adding lime to acidic lochs and soil.
Acid + metal →
salt + hydrogen
only works with metals above hydrogen in the Electrochemical series (databook pg 7)
Test for hydrogen gas
hydrogen burns with a pop!
Acid rain caused by
sulphur dioxide (from burning fossil fuels) and nitrogen dioxide (from sparking air in car engines) dissolving in water in atmosphere.
Acid rain has damaging effects on:
buildings made from carbonate (marble) rock
structures made of iron and steel
soils and plant and animal life
A salt is:
compound in which the hydrogen ions of an acid have been replaced by metal ions (or ammonium ions).
Hydrochloric acid salts are
chlorides
Nitric acid salts are
nitrates
Sulphuric acid salts are
sulphates
Some salts containing nitrogen e.g. ammonium nitrate, ammonium sulphate and potassium nitrate are used as
fertilisers
These salts are soluble in water
Precipitation
reaction of two solutions to form a solid
insoluble salts can be prepared by precipitation
(use pg 5 of databk)
Spectator ions
do not take part in the reaction
their state and charge remains the same.
Electricity can be produced in a simple cell by
connecting two different metals with an electrolyte between them.
Electrochemical series
Series produced by measuring the voltage produced between different metal pairs.
Displacement
a metal is added to a solution containing ions of a metal lower in the electrochmical series.
higher metal forms ions in solution. Lower metal forms atoms (solid at bottom of beaker).
Position of hydrogen in the electrochemical series
is worked out by reacting different metals with acid.
Direction of electron flow in a cell
electrons flow through wires from metal higher up in the electrochemical series to the one lower down.
Ion bridge
allows the movement of IONS to complete the circuit.
Oxidation
loss of electrons
Reduction
gain of electrons
Redox
oxidation and reduction occur together
metal ion → metal atom
reduction
metal atom → metal ion
oxidation
ion-electron equations
used to show oxidation and reduction
find them pg 7 of databooklet.
Metal + water →
metal hydroxide + hydrogen
only magnesium and above will react
Metal + oxygen →
metal oxide
only copper and above will react
Metal ore
naturally occuring metal compound
Metals found uncombined in Earth crust are
less reactive.
Gold, silver, copper
Blast Furnace
used to extract iron from iron ore
Blast furnace key reactions:
C + O2 → CO2
CO2 + C → CO
Fe2O3 + CO→ Fe + CO2 (these equations must be balanced)
Aluminium and other reactive metals are extracted from ores using
electrolysis
Corrosion
chemical reaction which involves the surface of a metal changing from an element to compound.
Rusting
corrosion of iron
Substances needed for rusting are
water and oxygen
Stages in rusting are:
1) iron atoms lose 2 e- to form iron (II) ions
2) Iron (II) ions lose 1 e- to form iron (III) ions
e- lost by iron go to water and oxygen to form hydroxide ions
Ferroxyl indicator turns blue in presence of
iron (II) ions
Ferroxyl indicator turns pink in presence of
hydroxide ions
Rate of corrosion increased by
acid rain
salt
attaching metal lower in electrochemical series (unless iron surface is completely covered)
Corrosion stopped by
Physical protection - painting etc
Sacrificial protection
Connecting iron to negative terminal of battery
Physical protection
barrier to water and oxygen
painting, greasing, electroplating, galvanising, tin-plating, coating with plastic
Sacrificial protection
Fe attached to metal higher in E.C.S.
e- flow from higher metal to Fe. Other metal corrodes in place of Fe.
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Term:
Definition:
Definition:

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