ACIDS BASES AND SALTS
1. ACIDS
An acid is a substance that ionizes/ dissociates or break down into ions
to give hydrogen ion (H+) AS THE ONLY POSITIVE ION, in solution.
An acid is also defined as a proton (hydrogen ion,H+) donor.
Eg HCl
+
H2 O
→
H+
+
Cl-
Physical properties of acids
-
Have a sour taste
Turns blue litmus red
Have a pH less than 7
Strong acids are corrosive
Strong acids (man-made acids)
These are acids which completely ionize/dissociate in water. They
produce a very high number of hydrogen ions (H+) when dissolved in
water. They have a Ph between 1 and 3.
Examples
Hydrochloric acid – HCl
Sulphuric acid – H2SO4
Nitric acid
- HNO3
Phosphoric acid- H3PO4
Examples of complete ionization
Space for two examples (2 equations)
Weak acids (natural acids/organic acids)
These are acids which partially/partly ionize/dissociate in water.
They produce very few hydrogen ions (H+) in solution. The have a Ph
between 4 and 7.
Examples
Ethanoic acid (found in vinegar)
Ascorbic acid (found in vitamin C)
Lactic acid (found in sour milk)
Citric acid (found in citrius fruits)
Example of partial ionization
CHEMICAL PROPERTIES OF ACIDS
1. Acids react with metals to form a salt and hydrogen gas.
NB: only the MAZIT metals (magnesium, aluminium, zinc, iron and
tin) metals are used.
Acid + metal → salt + hydrogen
HCl + Mg → MgCl2 + H2
Test for hydrogen gas
- Use a burning/lighted splint
- Result: pop sound is produced
2. Acids react with carbonates to give a salt, carbon dioxide and
water.
Acid + carbonate → salt + carbon dioxide + water
2HCl + CaCO3 → CaCl2 + CO2 + H2O
Test for Carbon dioxide
- Bubble/pass the gas in lime water
- Result: limewater turns milky/ white precipitate formed
3. Acids react with bases or alkalis to form salt and water. This
reaction is called a NEUTRALISATION REACTION.
ACID
HCl
+ ALKALI →
+ NaOH →
SALT +
NaCl +
WATER
H 2O
4. Acids react with metal oxides (bases) to form salt and water. This
is also a neutralization reaction.
Acid + metal oxide → salt + water
H2SO4 + CuO
→
CuSO4 + H2O
BASES
Bases are the opposite of the acids.
A base is defined as a proton acceptor.
A BASE WHICH IS SOLUBLE IN WATER IS CALLED AN ALKALI
NB: not all bases are alkalis but all alkalis are bases.
Some alkalis include NaOH, KOH, NH4OH
When dissolved in water , alkalis produce the common ion as
hydroxide ion (OH-).
An alkali can also be defined as a substance that ionizes/dissociate
in water to give hydroxide (OH-) AS THE ONLY NEGATIVE ION.
STRONG ALKALIS
These are alkalis which completely ionizes/dissociate in water.
They produce a high number of hydroxide ions (OH-) in solution.
Have a very high Ph range between 12 and 14.
Eg NaOH + H2O
→
Na+
+
OH-
WEAK ALKALIS
These alkalis partially/partly ionizes/dissociate in water. They
produce very few hydroxide ions when dissolved in solution. They
have a Ph between 8 and 11.
Eg NH4OH + H2O → NH4+
+ OH-
PHYSICAL PROPERTIES OF ALKALIS
-
They have a bitter taste
Turn red litmus blue
Have a Ph more than 7
Feel soapy to the touch ie they are slippery when felt
Chemical properties
1. Alkalis and bases react with acids forming a salt and water.
REMEMBER- This is called a neutralization reaction.
NaOH
+
HCl
→ NaCl
+ H 2O
2. Alkalis and bases react with ammonium salts to form a salt,
ammonia gas and water.
Eg NaOH + NH4Cl
→
NaCl + NH3 + H2O
Test for ammonia gas
- Use a wet/damp/moist red litmus paper
- Result: moist/damp/wet red litmus paper turn blue.
THE DIFFERENCE BETWEEN STRENGTH AND CONCENTRATION
Concentration is the amount of acid or alkali that is dissolved in a
litre of water.
NB; A concentrated solution of an acid/alkali contains a large
amount of the acid/alkali in a little amount of water.
A dilute solution contains little amount of the acid/alkali in a large
amount of water.
Strength of an acid/ alkali refers to the extent to which the
acid/alkali ionizes in solution (ie degree of ionization).
Strength therefore refers to whether the substance is strong or
weak.
INDICATORS
An indicator is a substance used to show whether another substance
is acid or alkaline. This is shown by change in colours to acid and
alkalis.
Examples of indicators
indicator
phenolpthalein
Methyl orange
Methyl red
Blue litmus
Red litmus
Universal
indicator
Colour of
indicator
colourless
orange
red
blue
red
green
Colour in acid
Colourless
pink
red
Red
red
Colour in
alkali
Pink
Yellow
Yellow
Blue
blue
NB: universal indicator shows different colours in strong acids and
weak acids. It also shows different colours in strong and weak alkalis
depending on the strength.
pH (potential hydrogen)
This is the measure of the acidity or alkalinity of a solution. The pH is
related to the number of hydrogen ions (H+) and hydroxide ions (OH) in solution.
NB: Low pH means high number of H+ ions and low number of OHions.
High pH means low number of H+ ions and high number of OHions.
pH is measured using a pH scale or pH meter.
The scale/meter ranges from 1 to 14, and has no units.
NB; THE pH METER/scale IS ONLY USED WITH RESULTS FROM THE
UNIVERSAL INDICATOR.
pH scale
diagram
APPLICATION OF ACID-BASE REACTIONS IN DAILY LIFE
- Brushing teeth with toothpaste: toothpaste is a base, and will
neutralize acid formed by bacteria in the mouth.
- Treatment of acidic soils:
Acidic soils are treated by adding quicklime (CaO) or slaked lime
[Ca(OH)2]
- Baking with baking powder:
Baking powder contains tartaric acid and bicarbonate of soda
(sodium hydrogen carbonate). The two solid react together when
dissolved in water produce carbon dioxide gases that raise dough.
- Treatment of indigestion;
Excess acid in the stomach leads to indigestion, this can be
neutralized by milk of magnesia, baking soda (NaHCO3).
- Treatment of stings;
Bees inject acidic liquids into the skin (methanoic/formic acid).
This is treated by calamine lotion (ZnCO3), bicarbonate of soda, or
baking soda.
Wasp stings are alkaline and are treated with vinegar.
- Descaling;
Removing scales or fur on electric kettles (caused by hard water)
using vinegar.
OXIDES
An oxide is formed when a metal or non-metal react with oxygen.
Types of oxides
Oxides are classfified into:
# Acidic oxide
# Basic oxides
# Neutral oxides
# Amphoteric oxides
-
1. Acidic oxides
These are non-metallic oxides which dissolve in water forming
acidic solutions.
Examples:
Carbon dioxide (CO2)
Sulphur dioxide (SO2)
Nitrogen dioxide (NO2)
Sulphur trioxide (SO3)
CO2 + H2O →
H2CO3 (Carbonic acid)
SO2 + H2O → H2SO3 (Sulphurous acid)
SO3 + H2O → H2SO4 (sulphuric acid)
When acidic oxides react with alkali, they are neutralized forming a salt
and water solution.
Example:
CO2
SO3
+
+
NaOH → Na2CO3
NaOH
→ Na2SO4
+ H2O
+ H2O
2. Basic oxides (they are alkalis)
These are metallic oxides which dissolve in water and form a
basic/alkaline solution
Examples
- Sodium oxide (Na2O)
- Calcium oxide (CaO)
- Potassium oxide (K2O)
Na2O + H2O → NaOH
CaO
+ H2O → Ca(OH)2
Basic oxides react with acids forming salt and water solution. Ie
they neutralize acids.
CaO + HCl
→ CaCl2 + H2O
3. Neutral oxides
These are the non-metallic oxides which dissolve in water
forming a neutral solution.
NB: these oxides do not react with acids and alkalis.
Examples:
- Carbon monoxide (CO)
- Nitrogen monoxide (NO)
- Hydrogen oxide (water) H2O
4. Amphoteric oxides
These are metallic oxides which have/show both acid and basic
properties. They react with both acids and bases in the same
way, forming a salt and water solution.
NB: they can act as both acids and bases
Examples:
- Zinc oxide (ZnO)
- Aluminium oxide (Al2O3)
- Lead II oxide (PbO)
Reaction with acid (act as a base)
ZnO + HCl
→
Al2O3 + HCl →
ZnCl2 + H2O
AlCl3 +
H2 O
Reaction with alkalis (act as an acid)
ZnO + NaOH
→
Na2ZnO2
[Sodium zincate]
+
H2 O
Al2O3 + NaOH →
NaAlO2
+
[Sodium aluminate]
H2 O
METHODS OF COLLECTING GASES
1. Collecting over water/downward displacement of water
The method is used to collect gases which are insoluble in
water, or gases that are slightly soluble in water ie where only
a small amount of the gas dissolve in water.
Examples:
H2, O2, CO2, Cl2, N2
Diagram
2. Upward displacement of air/ downward delivery
The method is used for dense gases (high density gases). The
gas displaces the air in the gas jar.
Examples:
Cl2, CO2, SO2
Diagram
3. Downward displacement of air/ upward delivery
The method is used for gases of low density (gases less denser
than air)
Examples:
Ammonia (NH3), H2
Diagram
4. Collecting using a gas syringe
This method is used for any gas, where the volume of gas
collected is to be measured.
Diagram
SALTS
A salt is formed when the hydrogen ion (H+) of an acid is
replaced by a metal ion.
NB: Salts are named from the acid in which they are obtained
from.
Salts made from hydrochloric acid are the CHLORIDES
Salts made from sulphuric acid are the SULPHATES
Salts made from nitric acid are the NITRATES
Salts made from phosphoric acid are the PHOSPHATES
A good number of salts are soluble in water while other salts
are insoluble.
NB: knowledge of the solubility rules of salts is important when
preparing salts
METHODS OF PREPARING SOLUBLE SALTS
1. REACTING ACIDS WITH SOLIDS [Metal, metal carbonate,
metal oxide (insoluble base)]
NB: only MAZIT metals are used. It is very dangerous to use
very reactive metals like potassium and sodium.
(a) Reacting an acid with a metal
Acid + metal →
H2SO4 + Mg →
(b) Reacting an acid with a carbonate
Acid + carbonate →
HCl + CaCO3
→
(c) Reacting an acid with an insoluble base
Acid + base →
HNO3 + CuO →
STEP 1: REACTION. Add excess solid (metal, carbonate or
insoluble base) to the acid. Excess solid ensures that all
the acid react to give a neutral salt solution.
NB; when using an insoluble base, the acid is first warmed
so as to speed up the reaction.
STEP 2: FILTRATION. The unreacted solid is removed from
the salt solution by filtration.
STEP 3: EVAPORATION. The salt solution is heated to
evaporate excess water, leaving the salt saturated.
STEP 4: CRYSTALLISATION. When saturated, the salt
solution is left to cool and crystallise at room
temperature.
STEP 5: WASH AND DRY. The crystals are washed with
small amounts of cold distilled water and dried with clean
tissues or filter papers.
(Diagram)
2. ACIDS ARE REACTED WITH ALKALIS
,
Acid + alkali → salt + water
HCl + NaOH → NaCl + H2O
In this method, the amount of acid needed to react with an
alkali is added with the help of an indicator such as methyl
orange. The acid is added slowly until an immediate
permanent ORANGE color is obtained indicating neutral
products. This is the END POINT. This method is known as
TITRATION.
END POINT; A point where an acid has completely
neutralized the alkali.
STEP 1: REACTION. The acid is titrated against the alkali to
obtain the salt solution.
NB; the salt solution may be filtered through charcoal to
remove the indicator.
-
TITRATION
Apparatus and materials used;
Burette
Pipette
Conical flasks (250 ml)
Beakers (250 ml)
Methyl orange
Retort stand
White tile
Funnel
Acid and alkali
Procedure
1.
2.
3.
4.
Fill the burette with the acid
Pipette 25 cm3 of the alkali into a conical flask
Add 2-3 drops of the methyl orange indicator
Add the acid slowly (drop by drop) into the conical flask with the
alkali until an immediate permanent ORANGE colour is obtained.
5. Record results in the table
6. Repeat steps 1-5 until consistent results are obtained
Titration
Final burette reading
(cm3)
Initial burette reading
(cm3)
Volume of acid used
(cm3)
Tick ( ) best results
1
2
3
4
Step 2: FILTRATION, the solution obtained from titration is then filtered
through charcoal to remove the indicator.
Step 3: EVAPORATION, the salt solution is then heated to evaporate
excess water, until the solution is saturated.
Step 4: CRYSTALLISATION, when saturated the salt solution is then left
to cool and crystallise at room temperature.
PREPARATION OF INSOLUBLE SALTS
Insoluble salts are generally prepared by the PRECIPITATION method.
This involves mixing two solutions of soluble salts. When these are
mixed a reaction takes place in which the salt is formed as a
PRECIPITATE (insoluble).
Examples:
1. preparation of Barium Sulphate (BaSO4)
We need a soluble Barium salt eg, BaCl2, Ba(NO3)2 etc. we also need a
soluble sulphate salt eg, Na2SO4, (NH4)2SO4 etc
SOLUBLE + SOLUBLE →
INSOLUBLE + SOLUBLE
barium nitrate + Sodium sulphate
Ba(NO3)2 + Na2SO4
→
→
2. Preparation of silver chloride (AgCl)
(skip two lines)
3. Preparation of lead (II) iodide (PbI2)
(skip 2 lines)
TEST FOR IONS (Qualitative analysis)
These are the tests conducted on unknown samples to determine
the positive ions (cations) and the negative ions (anions) in the
sample hence the name and formula of an unknown compound.
1. TESTING FOR NEGATIVE IONS (ANIONS): Each negative ion is
tested for using a different test reagent.
anion
test
Carbonate
Add dilute acid
Chloride
(in solution)
Acidify with
dilute nitric acid,
then add
Result
Bubbles/effervescence
Carbon dioxide gas
produced
White precipitate
Sulphate
(in solution)
Nitrate
(in solution)
Iodide
(in solution)
aqueous silver
nitrate
Acidify with
dilute nitric acid,
then add
aqueous barium
nitrate or barium
chloride
Add aqueous
sodium
hydroxide, then
alumnium foil;
warm carefully
Acidify with
dilute nitric acid,
then add lead (II)
nitrate
White precipitate
Ammonia gas
produced
Yellow precipitate
2. TEST FOR POSITIVE IONS (CATIONS): ONLY two test reagents
are use to test for positive ions;
(i) Aqueous Sodium hydroxide
AND/OR
(ii)
Aqueous ammonia (ammonium hydroxide)
cation
Aluminium
Zinc
Calcium
Copper (II)
Iron (II)
Effect of
aqueous
sodium
hydroxide
White ppt,
soluble in
excess giving a
colourless
solution
White ppt,
soluble in
excess giving a
clourless
solution
White ppt,
insoluble in
excess
Light blue ppt,
insoluble in
excess
Effect of aqueous
ammonia/ammonium
hydroxide
GREEN ppt,
insoluble in
excess
Green ppt, insoluble
in excess
White ppt, insoluble n
excess
White ppt, soluble in
excess giving a
colourless solution
No ppt, or very slight
ppt
Light blue pp, soluble
in excess giving a dark
blue solution
Iron (III)
ammonium
cation
Cu+2
Zn+2
Ca+2
Al+3
Fe+2
Fe+3
RED-BROWN
ppt, insoluble
in excess
Ammonia gas
produced on
warming
Result of
adding
NaOH(aq)
Light blue
ppt, insoluble
in excess
Red-brown ppt,
insoluble in excess
Result of
adding
NH4OH/NH3(aq)
Light blue ppt,
soluble in
excess giving a
dark blue
solution
Name and
formula of
precipitate
Copper (II)
Hydroxide