1. Unit 1 - Chemical Changes
and Structure

2. Reaction Rates

3. Reaction Rates
During the course of a chemical reaction, reactants are being converted into products.
Measurement of the rate of reaction involves measuring the ‘change in the amount’ of a reactant or product in a certain time.
The rate of reaction changes as it progresses, being relatively fast at the start and slowing towards the end.
What is being measured is the average rate over the time interval chosen.
Reactions can be followed by measuring changes in concentration, mass and volume.

4. Where property = mass/volume/concentration
Nat 5
Where property = mass/volume/concentration
The above is used when there is no change in
mass/volume/concentration measured, for example
during a colour change reaction.
Higher

5. Calculate the reaction rate; During the first 5 seconds
Time (s)
Volume (cm3)
0.00 5
7.00 10
10.50 15
12.00 20
12.75 25
13.00 30
Calculate the reaction rate;
During the first 5 seconds
Between seconds
Between seconds
Of the whole reaction
(7-0/5-0 = 7/5 = 1.4cm3/s)
( /15-10 = 1.5/5 = 0.3cm3/s)
( /30-20 = 0.25/10 = 0.025cm3/s)
(13-0/30-0 = 13/30 = 0.43cm3/s)

6. Collision Theory A chemical reaction can only occur if there is a
successful collision between reactant molecules.
From S3 we know that we can speed up a
chemical reaction by;
Decreasing particle size (increasing surface area)
Increasing concentration (of reactant)
Increasing temperature
Adding a catalyst

7. Collision Theory – Particle Size
The smaller the particle size, the higher the surface area.
The higher the surface area, the greater the number of collisions that can occur at any one time.
The greater the number of collisions, the faster the reaction.
Therefore the smaller the particle size, the faster the reaction rate.

8. Collision Theory – Concentration
The higher the concentration, the higher the number of particles.
The higher the number of particles, the greater the chance of collisions that can occur.
The greater the number of collisions, the faster the reaction.
Therefore the higher the concentration, the faster the reaction rate.

9. Collision Theory – Temperature
The higher the temperature, the higher the energy the particles have.
The higher the energy, the faster the particles move.
The faster the particles move, the greater the chance that they can collide
The greater the number of collisions, the faster the reaction.
Therefore the higher the temperature, the faster the reaction rate.

10. Catalysts
A catalyst speeds up a chemical reaction without getting used up or changed itself.
Catalytic converters are used in exhaust systems to turn harmful gases into less harmful gases. Platinum, rhodium (expensive transition metals) are used as the catalyst used in catalytic converters.

11. Atomic Structure
and Bonding

12. Modern Day Model of the atom
The atom is made up of a dense centre called the nucleus, which contains protons and neutrons
Modern Day Model
of the atom
The electrons are very light and are found in a space around the nucleus called the electron shell.

13. Particle Mass Charge Location Proton Neutron Electron1
+ve
(positive)
Neutron
(no charge)
Electron
zero
(1/1850)
-ve
(negative)

14. Mass Number 16 O
Symbol 8
Atomic Number

15. Important protons electrons neutrons protons atomic number
Atomic Number = No. of ____________
No. of Protons = No. of ____________
(when the atom is neutral!)
Mass Number =
No. of __________ + No. of _________
No. of Neutrons =
Mass Number - ____________ __________
protons
electrons
neutrons
protons
atomic
number

16. No. Of Protons No. Of Neutrons Atomic Number Mass Number
Symbol
No. Of Protons
No. Of Neutrons
Atomic Number
Mass Number
No. Of Electrons 24
12 Mg 12 80
35 Br 45 40
19 K 19
9 F 9

17. No. Of Protons No. Of Neutrons Atomic Number Mass Number
Symbol
No. Of Protons
No. Of Neutrons
Atomic Number
Mass Number
No. Of Electrons 16
8 O 31
15 P 24 12 S

18. What if the element isn’t neutral?
Symbol
No. Of Protons
No. Of Neutrons
Atomic Number
Mass Number
No. Of Electrons 24
12 Mg 2+ 80
35 Br - 40
19 K + 16
8 O 2-
when there is a charge the no of protons and electrons are
no longer equal. (i.e. the no of electrons change)

19. Isotopes Protons Electrons Neutrons 28 14 Si 25 31 63 29 Cu 65

20. Many elements exist as 2 or more isotopes.
Isotopes are atoms of the same element
(same number of protons) but have
different number of neutrons.
This means for isotopes, the atomic number
stays the same but the mass number
changes.
Many elements exist as 2 or more isotopes.

21. relative atomic mass The relative atomic mass (R.A.M.) of an
element is the average mass number for
a sample of that element.
R.A.M is related to isotopes.

22. example from whiteboard (think number line)63 65 Cu Cu 29 29
The relative atomic mass of Copper (Cu) is 63.5
What does this tell you about the proportion of the two types of isotopes in a sample of copper?
example from whiteboard (think number line)

23. Electron Arrangements
Electrons are arranged in shells (or energy levels.)
Lithium has the electron arrangement 2,1 so there
are two electrons in the electron shell closest to
the nucleus, and one in the next shell: Li X

24. The 2,8,8,2 Rule
For the first 20 elements (hydrogen to calcium) we follow the 2, 8, 8, 2 rule.
A maximum of 2 electrons are allowed in the 1st electron shell, 8 electrons in the second, 8 electrons in the 3rd and 2 electrons in the 4th.
Each electron shell must be full before the next one is started.

25. Coincidence? Maybe. Maybe Not?
The number of ______ ________ affect
the way that the atom reacts.
In other words…
Elements with the same number of outer
electrons (elements in the same group) have
similar chemical properties.
electrons
outer

26. Important groups to remember;
Group 1 – the alkali metals; all very reactive soft metals. (two examples are ________ and ________)
Group 2 – alkaline earth metals; similar to group 1 metals but not as soft or reactive. (two examples are ________ and ________)
Group 7 – the halogens; very reactive non-metals. (two examples are ________ and ________)
Group 8 – the noble gases; very unreactive non-metals (two examples are ________ and ______)
In-between groups 2/3 – transition metals. (two examples are ________ and ________)

28. Remember to label what each colour represents.
solids
Remember to label what each colour represents.
Colour in all of the solid elements one colour
(easiest thing to do is leave them white!)
Colour in all of the liquids one colour
(Bromine and Mercury)
Colour in all of the gases one colour.
(elements 1, 7, 8, 9, 17 and all of group 0)
liquids
gases

29. Carbon Magnesium Potassium Fluorine Oxygen Nitrogen Beryllium
For each of the elements (right)
write down the following;
Name –
Symbol -
Atomic number –
Mass number -
Number of protons –
Number of electrons –
Number of neutrons –
Electron Arrangement –
Group number –
Metal or non metal –
Solid, liquid or gas -
Carbon
Magnesium
Potassium
Fluorine
Oxygen
Nitrogen
Beryllium
Aluminium
Neon
Bromine

30. BONDING

31. Why do atoms bond? Noble gases have a complete outer electron shell.
This arrangement of electrons is very stable and
therefore other elements want to be like the noble
gases.
The number of electrons the element needs to lose or
gain to be like the noble gases is called the valency.
The atoms of other elements can collide together and
combine to achieve the full outer electron shell.

32. The Covalent Bond step 1 Two atoms get close enough to
each other to collide.
step 1 O X O X
Oxygen with red crosses – 2,6
Oxygen with purple crosses – 2,6

33. The Covalent Bond step 2 The two atoms are attracted
to one another through the positive nucleus
of one and the negative electrons of the
other.
step 2 O X O X

34. The Covalent Bond step 3 O molecule
The two atoms combine and share enough outer electrons for each of them to become stable (full outer shell)
step 3 X O
Oxzygen with red crosses – 2,8
Oxygen with purple crosses – 2,8
O molecule 2

35. Important !! A covalent bond is a shared pair of electrons
between non-metal atoms.
A covalent bond is held by the attraction of
the positive nucleus and negative outer
electrons of the different atoms.

36. Representing a Covalent Bond
We can represent one shared pair of electrons
(covalent bond) by “ ”.
Therefore hydrogen bonded to hydrogen would
be;
This molecule is made up of 2 atoms it is classed as diatomic.
Mono- means _______
Di- means ________
Tri- means ________
Tetra- means _______ H

37. Naming Covalent Compounds
If a compound name ends in ‘-ide’ then that
compound only contains two elements – e.g. carbon
nitride contains carbon and nitrogen only.
Sometimes prefixes are used in naming compounds –
e.g. silicon dioxide
Mono – one
Di – two
Tri – three
Tetra - four

38. Draw the following covalent elements and compounds
using both the lines and the overlapping electron
shells (circles.)
You only have to draw the outermost shell electrons.
Hydrogen Chloride (HCl)
Phosphorus Trichloride (PCl3)
Water (hydrogen oxide) (H2O)
Sulphur Fluoride (SF2)
Ammonia (nitrogen trihydride) (NH3)
Carbon Dioxide (CO2) *** tricky

39. Working out formula for covalent compounds
symbol
valency
cross
cancel
formula

40. Working out formula Examples
Hydrogen Sulphide
Hydrogen Chloride
Phosphorus Oxide
Carbon Sulphide
Hydrogen Fluoride
Carbon Chloride
Silicon Oxide
Carbon Hydride
Nitrogen Hydride
Carbon Nitride
symbol
valency
cross
cancel
formula

41. Shapes of Molecules tetrahedral Formula = CH4
Name = Carbon Tetrahydride;
Structure is drawn like C H
tetrahedral

42. NH H2O HCl
pyramidal
planar / bent
linear

43. Complete this task on paper – only use jotter as last resort
Challenge
use a pencil
Complete this task on paper – only use jotter as last resort
Your challenge is to;
Work out the chemical formula
Name the compound (use prefixes if necessary)
Draw the line drawing +
Draw these molecules using the overlapping circles methods
Identify the shape of the molecule
Selenium and Iodine
Hydrogen and Fluorine
Carbon and Chlorine
Carbon and Hydrogen
Nitrogen and Hydrogen
Phosphorus and Bromine
Carbon and Sulphur
Silicon and Oxygen

44. Properties of Covalent MoleculesH
Methane molecule (CH4)
This is a covalent molecule. Other examples include
water (H2O), oxygen (O2) and candle wax (C8H18).
Covalent molecules;
have low melting and boiling points
can be solid, liquid or gas at room temperature.
never conduct electricity (in any state)

45. Covalent Network Substances
Some covalent substances do not have individual
molecules.
Diamond is an example of a covalent network
structure. Sand (silicon dioxide) is another example.

46. This is a diamond structure (carbon atoms only.)

47. Properties of Covalent Network
Covalent networks;
have extremely high melting and boiling points.
are always solid at room temperature.
never conduct electricity
graphite is the exception to this rule as
although it is a covalent network it will
conduct electricity.

48. Weak intermolecular Bond
Strong Covalent Bond
(hard to break)
Weak intermolecular Bond
(easy to break)

49. Covalent Molecular Vs Covalent Network
Although the covalent bonds within covalent molecules are strong, the force of attraction between the molecules are weak.
These weak forces of attraction don’t require a lot of energy to break and therefore covalent molecules have low melting/boiling pts.
All bonds in a covalent networks are very strong covalent bonds and it takes a lot of energy to break these bonds – i.e. very high melting/boiling points.

50. Type of covalent substances
Molecular
Network
Melting/Boiling Pts
Type of bond broken
Strength of bonds broken
Low
Very high
Intermolecular
Covalent
Weak
Very strong

51. Typical Past Paper Questions - Bonding
Which line in the table shows the properties
of a covalent molecular compound?

52. Which of the diagrams (left)
show the structure of a
diatomic compound? A B C D

53. Which diagram (left) shows
the structure of a covalent
network structure? A B C D

54. Ionic Bonding
metal
When a __________ atom wants to become stable and bond with a non-metal atom, it gives away or loses electrons.
Metals lose electrons to get a full outer shell.
These ‘lost’ electrons don’t just disappear – instead;
Non metals gain electrons to get a full outer shell.

55. ion
ion
An ionic bond is between a metal and non metal. An ionic bond is held together via the attraction between the positive metal ion and the negative non metal ion.

56. Ionic compounds don’t form molecules like covalent substances
Ionic compounds don’t form molecules like covalent substances. They are arranged in large lattice structures.
The ________ metals ions are attracted to the _______ non metals ions.
This attraction is very strong making an ionic lattice a strong stable substance – because of this ionic compounds are always solid at room temperature.
Ionic compound conduct electricity when molten (melted) and when in solution but NOT as a solid. This is because the ions are free to move.
Ionic compounds have high melting points and boiling points normally in the range; 400oC to 1400oC.
diagram goes here

58. Making Ions Examples For each of the following elements;
Draw the electron arrangement of the atom.
Draw the electron arrangement of the ion.
Work out the charge of the ion.
Name the noble gas that the ion similar to.
a) Aluminium f) Oxygen
b) Magnesium g) Chlorine
c) Lithium h) Nitrogen
d) Calcium i) Sulphur
e) Potassium j) Bromine

59. Example; Beryllium Atom Ion
The Beryllium ion (charge ____) has the same electron arrangement as the noble gas _________________.

60. Atom
Ion
The ___________ ion (charge ____) has the same electron arrangement as the noble gas _________________.

61. Working out formulae of Ionic Compounds
symbol
valency
cross
cancel
formula
example from whiteboard

62. Working out formula with complex ions symbol valency (charge) cross
Complex ions are ions which contain more than one element.
We can identify complex ions from their names as they end in ‘ite’ or ‘ate’. (the exceptions to this rule are hydroxide and ammonium)
Sodium Sulphate
symbol
valency
cross
cancel
formula
(charge) Na
(SO42-) 1 2 2 1
Na2 (SO42-)

63. Lithium Chloride Manganese (III) Chloride
Work out the chemical formula for the following ionic compounds… Remember the roman numerals indicate the valency of transition metals. (Be careful some contain complex ions – see data booklet)
Lithium Chloride Manganese (III) Chloride
Aluminium Oxide Copper (III) Carbonate
Beryllium Bromide Rhodium (I) Oxide
Calcium Carbide Vanadium (II) Sulphate
Iron (II) Oxide Barium Carbonate
Copper (III) Iodide Cadmium (II) Hydroxide
Zinc (I) Sulphite Ammonium Phosphate
Calcium Hydroxide Ammonium Dichromate
Sodium Carbonate Zinc (II) Sulphide
Lithium Nitrate Hafnium (I) Permanganate
Hydrogen Sulphate Aluminium Ethanoate
Sodium Carbonate
Calcium Sulphate
Beryllium Fluoride
Palladium (II) Bromide
Nickel (III) Iodide
Titanium (II) Sulphite
Caesium Selenide
Barium Chromate
Potassium Phosphate
Silver (I) Oxide
Iron (III) Phosphide 63

64. example from whiteboard
Potassium iodide reacts with lead (II) nitrate to
form a yellow solid and a clear solution. The
products are thought to be lead (II) iodide and
potassium nitrate.
Word equation
Chemical equation
example from whiteboard

65. example from whiteboard
Molten iron is used to join steel railway lines together.
Molten iron is produced when aluminium reacts with iron (III) oxide. Another product is thought to be aluminium oxide.
Word Equation
Chemical Equation
example from whiteboard

66. example from whiteboard
Calcium carbonate reacts with hydrogen chloride to
form a calcium chloride, water and a gas. The gas
was tested and is turned limewater from colourless
to milky.
Write the chemical equation for the above
reaction.
example from whiteboard

67. Balancing Equation Examples
Not only do we always have to have the same elements on both sides of a ‘reaction arrow’ but we also need to have the same amounts of them too.
In order for this to happen we need to balance the equation. 2 2
H = H = 2
O = O = 1 4 4 2

68. Balancing Equation Practice
H2 + Cl2 HCl
Al + Cl2 AlCl3
C3H8 + O2 CO2 + H2O
Fe2O3 + CO Fe + CO2 ***
NaOH + H2SO4 Na2SO4 + H2O
NH3 + O2 NO + H2O ***
Mg(OH)2 + HCl MgCl2 + H2O ***
*** = tricky
Now try the sheet – do not write on it.

69. Calculations

70. Formula Mass The formula mass (or gram formula mass - gfm) of a
substance is obtained by adding the relative atomic
masses of all of the elements in a compound together.
In other words, the formula mass is the total mass
of a compound.
The formula mass has NO units.

71. example from whiteboard
Worked Example 1
Calculate the formula mass of calcium chloride.
Formula;
Formula Mass;
Calcium =
Chlorine =
(Total) =
example from whiteboard

72. example from whiteboard
Worked Example 2
Calculate the formula mass of hydrogen sulphite.
Formula;
Formula Mass;
Hydrogen;
Sulphur;
Oxygen;
(total)
example from whiteboard

73. example from whiteboard
Worked Example 3
Calculate the formula mass of magnesium nitrate.
Formula;
Formula Mass;
Magnesium =
Nitrogen =
Oxygen =
(Total) =
example from whiteboard

74. Work out chemical formula and formula mass of...
Hydrogen oxide Calcium carbonate
Carbon dioxide Lithium phosphate
Nitrogen fluoride Hydrogen nitrate
Aluminium phosphide Ammonium sulphide
Magnesium bromide Gold (I) sulphate
Copper (II) iodide Silver (I) ethanoate
Zinc (III) oxide Potassium dichromate
Iron (II) chloride Nickel (III) chromate
Lead (II) nitride Tin (II) oxide
Mercury (I) bromide Platinum (II) sulphite

75. The Mole One mole of a substance is the formula mass
but with the units grams.
For example;
One mole of calcium chloride = 111g
One mole of hydrogen sulphite =
One mole of magnesium nitrate =
use
previous
examples

76. Important !!! m n fm m = mass of substance n = no. of moles
fm = formula mass
Important !!!
(grams) m n fm
(moles)

77. Rearranging the equationm n fm
mass =
no. of mole x formula mass
m = n X fm
no of mole =
mass / formula mass
n = m fm
formula mass =
mass / no of moles
fm = m n

78. Calculate the number of moles in 10.1g of potassium nitrate (KNO3)
example from whiteboard m n fm

79. example from whiteboard
0.75moles of a compound X weighs 14g. Calculate the formula mass of the substance.
example from whiteboard m n fm

80. What mass of sodium carbonate (Na2CO3) is present in 0.5 moles?
example from whiteboard m n fm

81. Questions Calculate the mass of 3 moles of copper (I) bromide. CuBr
show all working
Calculate the mass of 3 moles of copper (I) bromide. CuBr
Calculate the mass of 4 moles of calcium nitrate. Ca(NO3)2
How many moles are there in 4.23g of magnesium carbonate? MgCO3
How many moles are there in 1kg of aluminium chloride? AlCl3
0.75 moles of compound Z weigh g, calculate the formula mass.
1.25 moles of compound Y weigh 69.40g, calculate the formula mass.

82. How many moles of each substance in;
A – 14g of Nitrogen gas (N2)
B – 84.5g of Magnesium carbonate
C – 400g of Copper (II) oxide
D – 321g of Iron (III) hydroxide
What is the mass of;
A – 1 mole of aluminium
B – 2.5 Moles of Oxygen gas (O2)
C – 0.5 moles of Lithium sulphate
D – 0.1 moles of ethane (C2H6)

83. Calculation using Balanced Equation
Balanced equations can be used to calculate masses
of substances involved in chemical reactions.
When we write the balanced chemical equation, the
number in front of each formula represents the
number of moles of the substance.
e.g. Ca + 2HNO Ca (NO3) H2

84. example from whiteboard
What mass of calcium chloride would be produced when 10g of
calcium reacts fully with hydrogen chloride?
Ca HCl CaCl H2
example from whiteboard

85. CH O CO H2O
What mass of carbon dioxide is formed when 64g of methane
are burned completely in air?

86. example from whiteboard
Calculate the mass of iron that would be produced
from 2 moles of iron (lll) oxide.
Fe2O H2 2Fe H2O
example from whiteboard

87. example from whiteboard
C9H2O O CO H2O
Calculate the mass of water produced when 6.4g of
nonane (C9H2O) is burned.
example from whiteboard

88. Acids and Bases

89. The pH scale… 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Acids
Alkalis
Neutral

90. What does it really mean?
Red very acidic
Orange - slightly acidic
Yellow - very slightly acidic
Green - neutral
Green/Blue – very slightly alkaline
Blue slightly alkaline
Purple - very alkaline 90

91. Oxide Acid formed Other name of acid Acid formula Ions in the acid
Carbon Dioxide
Carbonic Acid
Hydrogen Carbonate
Sulphur trioxide
Sulphuric Acid
Hydrogen Sulphate
Sulphur dioxide
Sulphurous Acid
Hydrogen Sulphite
Nitrogen Dioxide
Nitric Acid
Hydrochloric Acid
Hydrogen Chloride
Ethanoic Acid
Hydrogen Ethanoate

92. What conclusion can you come to..?
ALL acids contain H ions

93. Acids have an excess of hydrogen ions
An acid is a solution which has a greater
_____________________________
______________ than pure water.
i.e.
Acids have an excess of hydrogen ions
(H+ ions)

94. Oxide
Name of alkali formed
Formula of alkali
Ions in the alkali
Sodium Oxide
Sodium Hydroxide
Potassium Oxide
Calcium Oxide
Ammonia (NH3)
Ammonium Hydroxide

95. What conclusion can you come to..?
ALL alkalis contain OH ions

96. Alkalis have an excess of hydroxide ions (OH- ions)
An alkali is a solution which has a greater
______________________________
than pure water.
i.e.
Alkalis have an excess of hydroxide ions (OH- ions)

97. equation from white board
Neutral Substances
equation from white board
Water has equal amounts of H+ and OH- ions so the overall pH is neutral (pH 7)
During neutralisation reactions the H+ ion (from the acid) and the OH- ion (from the alkali) react and produce water.
The H+ and OH- ions being able to move is the reason why water conducts electricity.

98. Important Summary All solutions contain both H+ and OH- ions.
Acids have more H+ than OH- ions
Alkalis have more OH- than H+ ions
Neutral substances have equal amounts of H+ and OH- ions

99. Diluting Acids or Alkalis
Diluting an acid or alkali is similar to what happens with diluting juice. As an acid is diluted it becomes ‘weaker’ as there are less H+ ions. As a result the pH increases towards pH 7 until it eventually becomes neutral. As an alkali is diluted it becomes ‘weaker’ as there are less OH- ions. This means the pH decreases towards pH 7 (neutral)

100. Neutralisation Reactions
Acids can be neutralised using bases.
Worked Example
General Equation 1
ACID + Metal Hydroxide SALT + WATER
Sodium Hydroxide + Hydrochloric Acid Sodium Chloride + Water

101. Bases Bases are another word for alkalis. Examples of bases are;
metal hydroxides,
metal oxides and
metal carbonates.
Bases neutralise acids to form water.

102. Salts Salts are _______ _______ compounds
(contain a positive and negative charge).
We can name the salt produced by looking at
the acid and base that have been used.
soluble
ionic

103. The first part of the salt’s name comes from the base used;
e.g. Sodium Hydroxide gives – sodium …….
Calcium Oxide gives – calcium …….
Ammonia gives – ammonium…….
The second part of the name comes from the type of acid used during the neutralisation;
Hydrochloric acid produces chloride salts
Sulphuric acid produces sulphate salts
Nitric acid produces nitrate salts
Ethanoic acid produces ……………… ethanoate salts

104. questions 1. Complete the names of the salts produced;
a) Calcium Hydroxide + Hydrochloric Acid
b) Sodium oxide + Nitric Acid
c) Potassium carbonate + Sulphuric Acid
2. Which acid has been used to make calcium sulphate?
3. Write the word equation for magnesium hydroxide
reacting with an acid of your choice.

105. Metal Oxide + Acid reaction
General Equation
ACID + METAL OXIDE SALT + WATER
Example (from your experiment)
Word Equation
Chemical Equation (balance if necessary)

106. Add 10cm3 of acid and 10cm3 of water to a 100cm3 beaker.
Heat the acid until ALMOST boiling using a slightly blue flame.
When the acid is hot enough, use a spatula to add small amounts of copper(II) oxide (1g in total) to the beaker.
Stir the mixture gently for up to half a minute after each addition.
When all the copper(II) oxide has been added, continue to heat gently for 1 to 2min to ensure reaction is finished.
Allow the beaker to cool while you set up the filtration.

107. Metal Carbonate + Acid reactions
Calcium carbonate (limestone) is often used for building materials.
Metal carbonates react with acid, so acid rain can damage limestone buildings.
General Equation
ACID + METAL CARBONATE SALT + WATER + CARBON
DIOXIDE

108. _________________ and _________________
questions
1. Complete the names of the salts produced;
a) Calcium Hydroxide + Hydrochloric Acid
b) Sodium Hydroxide + Nitric Acid
c) Potassium Hydroxide + Sulphuric Acid
2. Which acid and alkali can be used to make Calcium
Sulphate?
_________________ and _________________
3. Write the word equation for magnesium hydroxide
reacting with an acid- you choose the acid.

109. Ionic Formula Writing ionic formula is simply writing the chemical
formula of a compound but showing the charges of
the ions.
Metals form positive/negative ions
Non metals form positive/negative ions.
Chemical Formula – NaOH
Ionic Formula – Na+ OH-

110. Spectator Ions When acids and alkalis are added to water, they
dissociate to form ions.
Some of these ions react, and some do not. The
ions that don’t react are called spectator ions.
Ionic equations make it easier to identify spectator
ions. Rather than write the chemical formula of
reactants and products, we can write the ions
present in aqueous solution:

111. Sodium Hydroxide + Hydrochloric Acid Worked Example
Word Equation;
Formula Equation (with state symbols);
Ionic Equation (with state symbols);
Spectator ions appear on reactant and product sides of the
reaction and remain unchanged, so we can cross them out.
Rewrite equation without (omitting) spectator ions

112. Mole and Solution Calculation
For solutions we use a different triangle.
This allows us to work out the number of
moles of solute (substance dissolved), volume
(in litres), or concentration of solution.
1 ml = 1 cm3
1000ml = 1000cm3
1000cm3 = 1 litre
1cm3 = 1/1000 = litres
200cm3 = 200/1000 = 0.2 litres
useful

113. Important !!! n c V n = no of moles (moles) c = concentration
(mol/l or moll-1)
v = volume
(litres)
Important !!! n c V

114. Rearranging the equationc V
no of moles =
concentration x volume
n = c X v
concentration =
no of moles / volume
c = n v
volume =
no of moles / concentration v = n c

115. example from whiteboard
Calculate the no. of moles of potassium hydroxide which must be dissolved to make 200 cm3 0.5 mol/l solution
example from whiteboard n c V

116. example from whiteboard
Calculate the concentration of the following solution: 0.1 mole of hydrogen chloride dissolved to make 500 cm3 of solution.
example from whiteboard n c V

117. example from whiteboard
Calculate the volume (in litres) of the following solution: 5 mol/l solution containing 2 moles of sodium hydroxide.
example from whiteboard n c V

118. Calculate the mass, in grams, of sulphuric acid (H2SO4) present in 500cm3 of a 5 mol/l solution

119. Titration Calculation
Step 1: Write the formula:
(C x V x P)acid = (C x V x P)alkali
C = Concentration of acid or alkali
V = Volume of acid or alkali – Always make sure the
units of volume are the same on each side!!
P = Number of H+ (for acid)
Number of OH- (for alkali)

120. Step 2:
Find out the number of H+ in the formula of
acid. (e.g. HCl= 1 , H2SO4= 2) OR
Find out the number of OH- in the formula
of alkali. (e.g. NaOH= 1, Mg(OH)2= 2)
Step 3: Put all of the known values into the
equation shown above.
Step 4: Complete the calculation.

121. Examples of Power Work out the powers of either the H+ (acids) or OH-
(alkalis) in the following;
NaOH
Ca(OH)2
H3PO4
HCl
H2SO4
Al(OH)3
H5S2

122. (C x V x P)acid = (C x V x P)alkali
Worked Example
In a titration, 10 cm3 of 2 mol/l sodium hydroxide (NaOH)
solution was neutralised by 25cm3 of dilute hydrochloric
acid (HCl).
Calculate the concentration of the acid in mol/l.
(C x V x P)acid = (C x V x P)alkali
example from whiteboard

123. example from whiteboard
In a titration, 25 cm3 of 2 mol/l sodium hydroxide (NaOH) solution was neutralised by 28.7cm3 of sulphuric acid (H2SO4). Calculate the concentration of the acid in mol/l.
example from whiteboard

124. example from whiteboard
In a titration, 20 cm3 of potassium hydroxide (KOH) solution was neutralised by 42.6cm3 of 0.5 mol/l hyrdochloric acid (HCl). Calculate the concentration of the base in mol/l.
example from whiteboard

125. example from whiteboard
In a titration, 2 mol/l sodium hydroxide (NaOH) solution was neutralised by 22cm3 of 0.1mol/l sulphuric acid (H2SO4). Calculate the volume of the base in cm3 that was neutralised.
example from whiteboard

126. example from whiteboard
In a titration, 0.05 mol/l potassium hydroxide (KOH) solution was neutralised by 17.1cm3 of 0.25mol/l phosphic acid (H3PO4). Calculate the volume of the base in cm3 that was neutralised.
example from whiteboard

127. Never (EVER) use 'rough' values in calculations