The document discusses various topics related to acids and bases: 1. It defines acids and bases, and classifies common substances as acidic, basic, or neutral based on their hydrogen ion concentration in water. Examples of strong acids and bases are provided. 2. Indicators and their uses are described, including litmus, phenolphthalein, methyl orange, and bromothymol blue. Everyday uses include testing soil pH and swimming pool chemistry. 3. Acidic oxides such as SO2, SO3, and CO2 are identified, and their position in the periodic table is linked to their acidity. Most transition metal oxides are basic.

1. The Acidic Environment

2. Classify common substance
as acidic, basic of neutral
• Alkali: Bases that dissolve in water
• All of the substances we use are acidic, basic or neutral. This classification is
usually based on the concentration of hydrogen ions produced when the
substance is dissolved in water.
• Acids:
• Common acids:
• Hydrochloric acid (HCl), sulfuric acid (H2S04) and nitric acid (HN03). vinegar,
lemon juice, aspirin Bases:
• Strong bases completely ionize in aqueous solutions to give OH - and a
cation Weak acids do not completely ionize in solution but exist in equilibrium.
Common bases:
• Na0H Ba(0H)2
• K20

3. Identify that indicators such as Litmus,
phenolphthalein, methyl orange and bromothymol
blue can be used to determine the acidic or basic
nature of a material over a range ,is identified by
change in indicator colour.
Indicators: Substances that change colour depending on the concentration
of hydronium ( H O + ) ions in a solution to determine whether a solution is
acidic or basic.
Examples of indicators include litmus, phenolphthalein, methyl orange and
Bromothymol blue.

4. Identify and describe some everyday uses
of indicators including the testing of soil
acidity/basicity
• Everyday uses of indicators include:
• The monitoring of the pH-level of fish-tank environments
• Some marine animals are very sensitive to pH levels of water
Sea water is about pH 8.5
• Testing the acidity or alkalinity of soils
• because some plants need an acidic soil
–E.g. azaleas and camellias

5. • While others need an alkaline soil
–e.g. most annual flowers and vegetables
• Agricultural Lime, calcium carbonate increases the basicity
while Manure increases acidity.
• Testing home swimming pools
•
• These need to be approximately neutral (close to pH: 7.4),
though adding chemicals to sanitise the water can change its
acid-alkali balance. If it is too acidic, it is dangerous to
humans, if it is too basic, then pests grow

6. Perform a first hand investigation to
prepare and test a natural indicator
Aim: to prepare and test a natural indicator
Method: Thinly slice some red cabbage leaves and place them
in a beaker
Add water into the beaker and boil the mixture while carefully
stirring with a rod
Allow the mixture to cool down and then filter out the cabbage
leaves from the solution. Use the solution to test the pH of
chemicals
Colour of indicator: Neutral: Blue/purple Acidic: pink to red
Basic: Green to yellow

7. Identify oxides of non-metals which acts as
a acids an describe the conditions under
which they act as acids
Acidic Oxides: One, which reacts with water to form acids, or one,
which reacts with bases to form salts.
Examples-SO2, SO3,CO2
Basic Oxides: One which reacts with acids to form salts, but does not
react with alkali solutions
Examples- CuO, Fe2O3
Neutral Oxides:
Examples- CO, NO, N2O

8. Analyse the position of these non-metals in the
periodic table and outline the relationship
between position of elements in the periodic table
and acidity/basicity of oxides.
Group 1 oxides are strongly basic. The basicity increases down
the group
Group 7 oxides are strongly acidic. The acidity decreases down
the group.
Generally, the acidity of oxides increases across the periodic
table and decreases down the periodic table. In group 3, the
trend down the group is that it starts from highly acidic, then
to amphoteric, then basic.
Most transition metal oxides are basic although some are
amphoteric.

9. Define Le Chatelier's principle
Chemical Equilibrium: When the forward rate of reaction
equals the reverse rate and the concentrations of all the
species in the system are constant.
Due to the equilibrium, the reactants are forming products at
the same rate at which the products are being converted back
to reactant, therefore the concentration of species are
constant.
Le Chatelier's Principle: If a system at equilibrium is
disturbed (i.e. Changes to equilibrium are made e.g. Adding
more reactants), then the system adjusts itself so as to
minimize the disturbance, re-establishing equilibrium.

10. Identify factors which can affect the
equilibrium in a reversible reaction
Factors which affect the equilibrium in a reversible
reaction:
Change in concentration
Change in temperature- If the temperature is lowered, the
exothermic reaction is favoured .
If the temperature is increased, the endothermic reaction is
favoured.
Change in total gas pressure
An increase in pressure leads to decrease in volume leading to
increase in concentration
A decrease in pressure leads to an increase in volume leading to
decrease in concentration.
By increasing the total gas pressure on the equilibrium, the
equilibrium shifts to counteract the increase in pressure

11. Describe the solubility carbon dioxide in water
under various conditions as an equilibrium
process and explain in terms of Le Chatelier's
principle
Carbon dioxide gas is soluble in water. The following
equation describes its equilibrium:
CO2 (g) CO2 (aq) Carbon dioxide gas dissolves in water.
CO2 (aq) +H 2O(l) H+(aq)+HCO3- (aq)
When there is more carbon dioxide, this reaction increases
concentration of hydrogen ions in water (Le Chatelier's
principle); therefore the water becomes more acidic and tastes
sharp
Example- Soda water contains carbon dioxide in water

12. -Identify natural and industrial source of sulpher
dioxide and oxides of nitrogen.
-Descriptive using equations, examples of
chemical reactions which release sulpher dioxide
and chemical reactions which release oxides on
nitrogen.
Natural sources of Sulphur dioxide ( SO2 ):
Geothermal hot springs
Emissions from volcanoes.
Bushfires
Natural decay of vegetation on land, wetlands and oceans

13. Industrial sources of sulphur dioxide:
The burning of fossil fuels which contain sulphur, such as
coal and oil ( industrial emissions) S (s) + O2 (g) = SO2 (g)
When metals are extracted from sulfide ores.
E.g. 2ZnS(s)+3O2 (g) = 2ZnO(s) + 2SO2 (g)
0ther examples of industries which emit sulphur dioxide
include:
Petroleum refineries Cement manufacturing Metal processing
facilities
0xides of nitrogen: Dinitrogen monoxide (N20), nitric oxide
(N0), nitrogen dioxide (N02)
Natural sources of oxides of nitrogen ( NO, NO2 , N2O ):
Lightning O2 (g)+N2 (g = 2NO(g)

14. Analyse information from secondary
sources to summaries the industrial
origins of sulpher dioxide and oxides of
nitrogen and evaluate reasons for
concerned concern about their release into
the environment
• Reasons for concern about release of SO2 and NOx
• These oxides combine with water to form an acidic solution
in the atmosphere, creating acid rain.
• Quick summary of effects
–Kills marine life
–Kills plants and forests
–Erodes the build environment

15. • Increase in oxides have led to increased production of
photochemical smog, especially in large cities which brings
forth respiratory problems for humans and animals, i.e.
breathing difficulties, tissue damage,

16. Assess the evidence which indicates
increase in atmospheric concentration of
oxides of sulpher and nitrogen
• Evidence of the oxides of sulphur and nitrogen being formed:
• 0xides of sulphur and nitrogen are mainly found in the
atmosphere as a result to a number of natural process such
as:
• Volcanic action, which produces sulphur dioxide
• Natural decay of vegetation on land and sea ( SO2 ) Lightning
which produces nitric oxide
• Due to the industrial revolution and increased human
population, there has been an increase of these oxides.
Manmade sources of the pollutants include

17. Calculate volumes of gases given masses of
some substances in reactions, and
calculate masses of substances given
gaseous volumes, in reactions involving
gases at 0°C and 100 kPa or 25°C and 100
kPa
• At 100kPa (1 bar):
-STP (standard temperature and pressure):
273K (0oC), molar volume of gases is 22.71 L/mol
-SLC (standard laboratory conditions) or RTP (room temp and
pressure):
-298K (25oC), molar volume of gas is 24.79 L/mol

18. At 101.3kPa (1atm):
-STP:
273K (0oC), molar volume of gases is 22.41 L/mol
-SLC or RTP:
298K (25oC), molar volume of gas is 24.47 L/mol
• In a reaction involving gases only, moles ratio can be
considered as volume ratios

19. Explanations formation and effects of
acidic rain
What is Acid Rain-
• Acid rain is rain with a higher concentration of hydrogen ions
( H +) than normal, which causes the rain to be more acidic.
• This rise of acidity is caused by high emission of SO2 and
NOx , which is released into the atmosphere by the burning of
fossil fuels.
• The major cause of acid rain is the burning of fossil fuels
from power plants, which produces electricity, and
automobiles.
Formation of acid rain:
• Pollutants such as smoke and fumes released from burning
fossil fuels rise up into the atmosphere and combine with
moisture (water) to form acid rain

20. • The main chemicals in these pollutants are SO2 and
nitrogen oxides ( NOx ) The sulfur dioxides and nitrogen
oxides react with water in the atmosphere to produce mild
acidic solutions of sulfurous acid and nitric acid.
• Rainwater, snow, fog and other forms of precipitation
containing these mild acidic solutions fall into the earth as
acid rain
• Effects of acid rain:
• Destruction of plants and forests. The change of pH levels in
soil due to the acid levels in acid rain creates difficulty for
plants to absorb sufficient minerals and kills important
micro-organisms which help sustain life. Also leaves of trees
are damaged and lost, due acidic rain

21. Define acids as proton donor and describe
the ionization of acids in water
• When acids dissolve in water, it can donate a proton, ie a
Hydrogen ion ( H +) to a water molecule.
• The proton ion in water solution may be represented as
H+(aq) or H3O+(aq ) (hydronium)
•
• For example, the ionization of HCl

22. Identify acids including acetic(ethanoic),
citric (2- hydroxypropane -1,2,3-
tricarboxylic), hydrochloric and sulfuric
acid
• Acetic/Ethanoic acid ( CH3COOH ):
• Acetic acid is a natural acid and is a component in vinegar
• Citric Acid ( C6 H8O7 ):
• Citric acid is also a naturally occurring acid, found in almost all
living things 0ccurs in citrus fruits such as lemons and oranges
• Hydrochloric acid ( HCl )
• Another natural acid (it is present in the stomach for digestion)
• Sulfuric acid ( H 2 SO4 )
• Non-natural acid, produced for many industrial processes
• Major uses include: fertilizers, manufacture of chemicals, electrolyte
in batteries

23. Describe the use of pH scale in comparing
acids and bases
The pH scale is a measure of the concentration of hydrogen
ions in solution. Low pH means a high level of proton
(hydrogen ions) concentration and thus very acidic and a high
pH means a low level of proton concentration, so it is basic. "

24. Describe acids and their solutions with the
appropriate use of the terms strong, weak,
concentrated and diluted
• Strong acids: ones that disassociate completely or almost
completely in water, forming H
• Weak acids: those that partially ionize in water
• Concentrated acids: are those that have a large amount of
solute in a given amount of solution. It has a high number of
moles per litre
• Dilute acids: those that have low amount of solute ionized in
a given amount of solution.
• Important: The strength of an acid does not depend on the
concentration. Strength depends on the type of acid, eg HCl
can never be called a weak acid

25. Identify pH as - log10 (H+) and explain that
a change in pH of 1 means a ten-fold
change in (H+)
• pH is a measure of the concentration of hydrogen ion, hence:
pH=-log10 [H + ]
• By definition, [H +][OH -] =10^-14 ie. pH +pOH =14
• [H3O+]=10^-pH . So if pH = 7, then the concentration of
H3O+= 10^-7 mol/L (neutral)
• A neutral substance is one where [H+]= [0H-] = 1x10-7molL-1
• A change in pH of 1 means a ten-fold change in [H+]. This
means that a solution of pH 1 has 10 times the concentration
of [H+] in pH 2 and 100 times the concentration in pH 3. The
pH scale is a logarithmic scale,

26. Compare the relative strengths of equal
concentrations of citric, acetic and
hydrochloric acids and explain in terms of
the degree of ionization of their molecules
• Citric and acetic acid are both very weak acids. Acetic acid is
a weaker acid with a pH of around 2.9 while citric acid has a
pH of around 2.1
• Hydrochloric acid is a much stronger acid, around pH 1. This
means that HCl has a very high degree of ionization, so the
forward reaction in:
HCl(aq) +H 2O(l) =H 3O(aq)+Cl(aq) is much more favoured,
producing more ions. Acetic acid on the other hand has a low
degree of ionization, and mainly exists in water as acetic acid
molecules rather than ions.

27. Describe the difference between a strong
and a weak acid in terms of an equilibrium
between the intact molecule and its ions"
• Strong acids disassociate completely or almost completely in water,
releasing H ions in aqueous solution. Since the production of
ions is much, much more favoured, therefore the equilibrium lies
well to the right. For example:
HCl(aq) +H 2O(l) = H3O(aq)+Cl-(aq)
• Weak acids only ionize partially in water, resulting in an equilibrium
between its intact molecules and ions

28. Outline the historical development of
ideas about acids including those of -
Lavoisier - Davy - Arrhenius
Lavoisier:
• Proposed that an acid must contain oxygen
• Non metal oxides reacted with water to form acidic solution
Davy:
• Redefined an acid as a substance containing hydrogen
(rather than oxygen), which could be partially or totally
replaced by metals when it reacted together, to produce a
salt. He realized this by discovering that HCl had acidic
properties.

29. Arrhenius:
Suggested that acids ionized in solution to produce hydrogen
ions as the only positive ion in solution
Bases produced hydroxide ions as the only negative ion in
solution Completely ionized acids were strong acids, weak if
partially ionized Arrhenius proposed these ideas only when the
acid was in an aqueous state

30. Outline the Bronsted-Lowry theory of acids
and bases
• Bronsted-Lowry acid: Any molecule or ion that acts as a
proton donor (hydrogen ions)
• Bronsted-Lowry base: Any molecule of ion that can act as a
proton acceptor
• The Bronsted-Lowry theory of acids and bases states that an
acid is a substance which in solution donates a proton, while
a base is a substance that accepts a proton
• HCl (g) +H 2O(l) =H 3O +Cl-
• HCl (acid)is the proton donor, since it gives a H+ ion Water
(base) is the proton acceptor, accepting the H+ ion
• Monoprotic acids: Acids that ionize and lose one protons
Diprotic acids: Acids that ionize and lose two protons

31. Describe the relationship between an acid
and its conjugate base and a base and its
conjugate acid
• Acids donate a proton to form its conjugate base, while bases
accept a proton to form its conjugate acid.
• Acid +Base=Conjugate base (ofacid )+Conjugate acid (of base)
eg. HCl (aq) +H O(l) = Cl- +H3 O+
• Cl - is the conjugate base of HCl (acid)
• H3 O + is the conjugate acid of H O (base)

32. Identify a range of salts which form
acidic, basic or neutral solutions and
explain their acidic neutral or basic
nature

33. Identify conjugate acid/base pairs
Acid +Base=Conjugate base (of acid )+Conjugate acid (of
base)
Strength Acid Conjugate Base
Strong H2SO4
HCl
HSO −
4
Cl −
Weak CH 3COOH CH3COO-−
H2CO3 HCO3 −
NH 4+ NH
HSO4− SO4 (2−)
4 SO4
Very Weak H 2O OH−

34. Identify amphiprotic substances and
construct equations to describe their
behaviour in acidic and basic solutions
Amphiprotic : substance that can act either as a proton donor
(acid) or a proton acceptor (base)
Examples include: H 2O

35. Some More points described in detail in the Notes
“The Acidic Environment” are listed Below
• Identify neutralization as a proton transfer reaction which is
exothermic
• Describe the correct technique for conducting titrations and
preparation of standard solutions"
• Qualitatively describe the effect of buffers with reference to a
specific example in a natural system"
• "Analyse information from secondary sources to assess the
use of neutralization reactions as a safety measure or to
minimize damage in accidents or chemical spills
• "Describe the differences between the alkanols and alkanoic
acid functional groups in carbon compounds

36. • Identify the IUPAC nomenclature for describing the esters produced
by reactions of straight chained alkanoic acids from C1 to CB and
straight- chained primary alkanols from C1 to CB
• Explain the difference in melting point and boiling point cause by
straight-chained alkanoic acid and straight-chained primary alkanols
structures“
• Identify Esterification as the reaction between an acid and an
alkanols and describe using equations, examples of esterification
• Describe the purpose of using acid in esterification for catalysis
• Explain the need for refluxing during esterification
• Identify data, plan, select equipment and perform a first hand
investigation to prepare an ester using reflux
• Outline some examples of the occurrence, production and uses of
ester
• Process information from secondary sources to identify and describe
the uses of esters as flavours and perfumes in processed foods and
cosmetics

37. Thanks for viewing and visiting
Reference and acknowledgment
- Compiled Notes of Isaac Seunglee Suh