Week 3, Lesson 3
Chapter 13 – Introducing Acids &
Bases
Acids
• Acids are commonly used in our homes.
• Many foods contain acids to enhance flavour or
as a preservative.
• The sour taste of some food substances is due to
the presence of acid.
• In industry solutions of acids are used extensively
to produce a wide range of products, such as
fertilisers, drugs, explosives and plastics.
• Acids are also used to clean metal surfaces before
use.
Acids cont…
• The three most commonly used acids in the
lab are hydrochloric acid, sulfuric acid and
nitric acid.
• Acids can cause severe problems if misused.
• Eg, acid rain, acid contamination.
Properties of Acids
•
•
•
•
•
Change the colour of some indicators.
Tend to be corrosive
Taste sour
React with bases
Have a relatively low pH.
Bases
• Acids react with bases.
• In solution, bases sometimes taste bitter and have a
slippery feel.
• They react with some plant extracts to counteract the
effect of acids.
• If a base is continually added to a solution of an acid,
the properties of the acid slowly disappear.
• The same goes for if an acid is added to a basic
solution.
• The acid and base are said to have neutralised each
other.
• Bases are effective cleaners because they react with
fats and oils to produce water-soluble soaps.
Properties of Bases
•
•
•
•
•
Turn litmus blue
Have a slippery feel
Are caustic
React with acids
Have a relatively high pH
Safety with Acids and Bases
• Acids and bases should be treated with caution and
you should avoid these solutions coming in contact
with your skin and eyes.
• You should always wear safety goggles and lab coats.
• Concentrated sulfuric acid is a viscous liquid and
accidents can happen if water is added to the acid in
order to produce dilute solutions.
• The ionisation/hydrolysis of sulfuric acid is an energyreleasing reaction.
• A common method of neutralising an acid spill is to
add sodium hydrogen carbonate powder.
Indicators
• One of the characteristic properties of acids and bases is their
ability to change to colour of certain plant extracts.
• Litmus is a dye obtained from lichen.
• In the presence of acid, litmus turns red.
• Such plant extracts are called indicators.
• Indicators are often extracted from plant dyes and are themselves
acids or bases.
• They change from one colour in acid to another in basic solution.
• Common indicators include methyl orange, phenolpthalein and
litmus.
• Universal indicator is a mixture of many indicators and changes
through a range of colours to easily establish the pH of a solution.
• The indicators that undergo a single colour change are also used for
many analyses.
• pH meters are often used in laboratories for more accurate
determination of solution acidity and are not affected by coloured
solutions.
Indicators and pH Ranges
Reactions Involving Acids and Bases
• Acids react with many metals particularly those found in the main groups
of the periodic table, although they also react with several transition
metals.
• These reactions typically produce a solution of a metal salt and hydrogen
gas.
• Salts are compounds usually made up of a metal cation and a non-metal
anion.
• The salt produced will depend on the acid used in the reaction.
• Acids also react with many compounds such as metal hydroxides and
metal carbonates.
• A salt is again produced with each of these, together with water.
• In the case of metal carbonates, carbon dioxide is also formed.
General Reaction Types involving
Acids
• There are numerous ways in which acids and
bases react.
• It is possible to group some reactions together
on the basis of the reactants involved and the
products formed.
• The following are six of the more common
reaction types.
Acid + Reactive Metal  Salt +
Hydrogen
• Reactive metals include Ca, Mg, K but not Cu,
Ag or Au.
• When dilute acids are added to main group
metals, and some transition metals, bubbles
of hydrogen gas are released, and a salt is
formed.
• For example, the reaction between dilute
hydrochloric acid and zinc metal can be
represented by the equation:
2HCl(aq) + Zn(s)  ZnCl2(aq) + H2(g)
Cont…
• This reaction can also be represented by an
ionic equation.
• In an aqueous solution the hydrochloric acid is
ionised and the zinc chloride is dissociated.
• The equation can therefore be written as:
2H+(aq) + 2Cl-(aq) + Zn(s)  Zn2+(aq) + 2Cl-(aq)
+H2(g)
However the chloride is the only spectator ion.
The ionic equation is therefore:
2H+(aq) + Zn(s)  Zn2+(aq) + H2(g)
Acid + Metal Hydroxide  Salt +
Water
• Metal hydroxides include NaOH, Ca(OH)2 and
Mg(OH)2.
• The hydroxide ions from metal hydroxides
readily react with the H+ ion from acid.
• The products are a salt and water.
• For example, the reaction between solutions
of sulfuric acid and sodium hydroxide can be
represented as:
H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + H2O(l)
Cont…
• The sulfuric acid is ionised in solution and both
sodium hydroxide and sodium sulfate are ionic, and
therefore dissociate.
• Water however is a covalent molecular substance
that does not ionise to a significant extent.
• So the equation becomes:
2H+(aq) + SO42-(aq) + 2Na+(aq) + 2OH-(aq) 2Na+(aq) +
SO42-(aq) + 2H2O(l)
The ionic equation is therefore:
H+(aq) + OH-(aq)  H2O(l)
Acid + Metal Oxide  Salt + Water
• Metal oxides include Na2O, MgO, CaO and
ZnO.
• Metal oxides are usually basic since they
contain the oxide ion.
• Water soluble oxides tend to form the
hydroxide ion.
• The reaction between dilute nitric acid and
solid calcium oxide can be represented by the
equation:
2HNO3(aq) + CaO(s)  Ca(NO3)2(aq) + H2O(l)
Cont…
• The calcium oxide is solid, so the ions are not
dissociated. The nitric acid is ionised in
solution and therefore dissociates.
• The equation becomes:
2H+(aq) + 2NO3-(aq) + CaO(s)  Ca2+(aq) + 2NO3-(aq) +
H2O(l)
The nitrate ions are the spectator ions so the
ionic equation is:
2H+(aq) + CaO(s)  Ca2+(aq) + H2O(l)
Acid + Metal Carbonate  Salt +
Water + Carbon Dioxide
• Metal carbonates include Na2CO3, MgCO3 and CaCO3.
• Acids reacting with metal carbonates produce carbon
dioxide gas with a salt and water.
• For example the reaction between a solution of nitric
acid and solid magnesium carbonate can be
represented by the equation:
2HNO3(aq) + MgCO3(s)  Mg(NO3)2(aq) + H2O(l) + CO2(g)
The nitrate ions are the spectator ions so the ionic
equation is:
2H+(aq) + MgCO3(s)  Mg2+(aq) + H2O(l) + CO2(g)
Acid + Metal Hydrogen Carbonate
 Salt + Water + Carbon Dioxide
• Metal hydrogen carbonates include NaHCO3, KHCO3 and
Ca(HCO3)2.
• Acids added to metal hydrogen carbonates (aka
bicarbonates) produce a salt, water and carbon dioxide.
• For example, the reaction between solutions of
hydrochloric acid and sodium hydrogen carbonate can be
represented by the equation:
HCl(aq) + NaHCO3(aq)  NaCl(aq) + H2O(l) + CO2(g)
The sodium and chloride are spectator ions in this reaction
so the ionic equation is:
H+(aq) + HCO3-(aq)  H2O(l) + CO2(g)
Acidic Oxide (non-metal oxide) +
Base  Salt + Water
• Acidic oxides include SO2, SO3, P4O10 and CO2.
• When oxides are added to water they form
acidic solution.
• The reaction of these with bases produce a
salt and water.
• For example the reaction between carbon
dioxide and a solution of calcium hydroxide
can be represented by the equation:
CO2(aq) + Ca(OH)2(aq)  CaCO3(s) + H2O(l)
Cont…
• The table below lists some common acidic
oxides and the anions they produce in the
reaction.
Acidic Oxide
Acid Formed when
oxide is added to
water
Anion produced in
reactions with
bases
CO2
Carbonic Acid
(H2CO3)
Carbonate (CO32-)
SO2
Sulfurous Acid
(H2SO3)
Sulfite (SO32-)
SO3
Sulfuric Acid
(H2SO4)
Sulfate (SO42-)
P4O10
Phosphoric Acid
(H3PO4)
Phosphate (PO43-)