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chemistry
CH3001
CORE CONCEPTS IN CHEMISTRY
NCEA LEVEL 3
2013/1
chemistry
ncea level 3
Expected time to complete work
This work will take you about 10 hours to complete.
This booklet gives you an overview of core NCEA Level 2 chemistry concepts so you can work
towards all of the Level 3 chemistry achievement standards.
In this topic you will focus on the following learning outcome:
•• demonstrating understanding of core chemistry concepts.
This will involve:
•• describing the structure of the atom
•• writing chemical formulae and balancing chemical equations
•• using the mole concept in calculations.
Copyright © 2013 Board of Trustees of Te Aho o Te Kura Pounamu, Private Bag 39992, Wellington Mail Centre, Lower Hutt 5045,
New Zealand. All rights reserved. No part of this publication may be reproduced or transmitted in any form or by any means without
the written permission of Te Aho o Te Kura Pounamu.
contents
1
Elements and the Periodic Table
2
Structure of the atom
3
Electrons and energy levels
4
Bonding
5
Valency and formulae
6
Naming compounds
7
Chemical reactions
8
Balancing symbol equations
9
The mole
10
Mole calculations
11
Teacher-marked assignment
12
Glossary
13
Answer guide
14
Periodic Table
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how to do the work
When you see:
Use the Topic webpage or the Internet.
1A Complete the activity.
Check your answers.
Contact your teacher.
You will need:
•• a pen or pencil and a ruler
•• a computer with Internet access and access to the Topic webpage.
Resource overview
This topic consists of 10 lessons covering the fundamental concepts of curriculum level 7
chemistry. It is recommended that you complete this booklet to revise these concepts. If you feel
confident that you have understood the concepts of a lesson, you can skip the activities. You are
expected to complete the teacher-marked assignment CH3001A.
The theme of this resource is how chemists use the Periodic Table as a tool to help them make
sense of chemistry. If you have difficulty with these concepts, it is advisable to enrol in CH2000
first, before continuing with CH3000. Keep this booklet as a handy reference.
You will get the most out of your studies if you use this booklet alongside a computer with an
Internet connection, using the Topic webpage. It is possible to study this topic using just the
booklet if you read the explanations and the answers very carefully.
Complete all the activities. Make sure that you attempt all the written activities. Mark your own
answers, using the Answer guide. Try to think critically about the chemistry involved.
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1
elements and the periodic table
learning intention
In this lesson you will learn to:
•• write the symbols for elements and use the Periodic Table as a tool to sort the elements.
making sense of chemistry
Although there are millions of different substances in the universe, they are all composed from
just over 100 known elements.
An element is a substance which cannot be broken down further by chemical means. You will
come across another definition in lesson 2.
More than 100 elements is still a lot of chemistry to learn and make sense of. Fortunately, since
the early twentieth century we have been able to sort them very easily using a Periodic Table,
based on their atomic number (see lesson 2) and their properties. You will investigate some of
these periodic properties in detail in the booklet Periodic trends (CH3041).
The Periodic Table consists of 18 groups (columns) and 7 periods (rows). The extra inserts called
the lanthanide and actinide series have not been shown.
1
1
18
1 H
2
3
2 Li
The Periodic Table
2
13
14
15
16
17
4
5
6
7
8
9
10
He
Be
B
C
N
O
F
Ne
11
12
13
14
15
16
17
18
Al
Si
19
31
32
3 Na Mg
3
4
5
6
7
8
9
10
11
12
P
S
Cl
Ar
4 K
20
21
22
23
24
25
26
27
28
29
30
33
34
35
36
Ca
Sc
Ti
V
Cr Mn Fe Co
Ni
Cu Zn Ga Ge As
Se
Br
Kr
37
38
39
40
41
42
46
52
53
54
5 Rb
43
44
45
47
48
49
50
Sr
Y
Zr Nb Mo Tc
Ru Rh Pd Ag Cd
55
56
57
72
76
87
88
6 Cs Ba
7 Fr
73
74
75
In
Sn Sb
Te
I
Xe
77
78
80
81
82
83
84
85
86
Tl
Pb
Bi
Po
At Rn
79
La
Hf
Ta
W
Re Os
Ir
Pt Au Hg
89
104
105
106
107
109
110
Ra Ac
108
51
111
Rf Db Sg Bh Hs Mt Ds Rg
metal
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metalloid
non-metal
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elements and the periodic table
metals
gold
magnesium
te kura
te kura
te kura
don laing
Elements can be sorted in different ways. For example, the Periodic Table on the previous page
has the elements colour-coded as metals, metalloids and non-metals. A few metals are shown
below.
copper
aluminium
You will notice in the above photos that metals are shiny and can be shaped into different forms.
All metals have the following five properties. Metals are:
•• lustrous (shiny)
•• conductors of electricity
Did you know…
Although most of the
elements are metals, they
make up less than 0.1% of
your body.
•• thermal conductors (conductors of heat)
•• malleable (can be hammered into shapes)
•• ductile (can be drawn into wires).
non-metals
sulfur
fluorine
public domain
Jurii
don laing
public domain
Elements that don’t have these properties are called non-metals. Non-metallic elements have a
range of physical properties. Unlike metals, they are not good conductors of electricity
(exception – graphite) and heat. They do not exhibit the properties of lustre, malleability or
ductility. They may be solids, liquids or gases. Photos of some non-metals are shown below.
bromine
iodine
metalloids
The classification of elements into metals and non-metals is very useful. However, some
elements, such as silicon and germanium, are difficult to classify as either metal or non-metal
(they have properties of both). These are known as metalloids. For example, pure silicon is shiny
but is a very poor conductor of electricity. Metalloids are often semiconductors.
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elements and the periodic table
check your understanding
sodium
magnesium
aluminium
silicon
phosphorus
sulfur
chlorine
te kura
te kura
te kura
wikimedia commons
wikimedia commons
te kura
te kura
The following elements are all from period 3 of the Periodic Table.
te kura
1A
neon
1. Which of the elements in period 3 are metals? Justify your answer.
2. Which of the elements in period 3 are non-metals? Justify your answer.
3. Silicon is a metalloid, as it has properties of both metals and non-metals. List a property of a
metal and a property of a non-metal that silicon has.
Check your answers.
symbols for elements
Chemists use a shorthand version for each element’s name, called a symbol. For example, C is
the symbol that represents carbon, and S is the symbol that represents sulfur.
For many elements the symbol is the first letter of the name of the element. If there are several
elements starting with the same letter, then there are two letters in the symbol. For example, Ca
is the symbol that represents calcium.
There are two rules when writing symbols for elements.
1. The first letter of a symbol is always a capital letter.
2. Where there is a second letter it is always a small letter. For example, the symbol for
magnesium is Mg (not MG).
Sometimes the symbol has been taken from the Latin name. For example, Fe is the symbol for
iron, because iron used to be called ferrum. Potassium (K) and sodium (Na) are given their
symbols based on the Latin names kalium and natrium.
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elements and the periodic table
Below is a list of the names and symbols of common elements. You should know these.
1B
Name
Symbol
Name
Symbol
aluminium
Al
lead
Pb
argon
Ar
lithium
Li
barium
Ba
magnesium
Mg
bromine
Br
mercury
Hg
calcium
Ca
manganese
Mn
carbon
C
neon
Ne
chlorine
Cl
nitrogen
N
chromium
Cr
oxygen
O
copper
Cu
potassium
K
fluorine
F
silicon
Si
helium
He
silver
Ag
hydrogen
H
sodium
Na
iodine
I
sulfur
S
iron
Fe
zinc
Zn
online activity
Visit the Topic webpage to play games to help you learn the names and symbols of the elements.
1C
check your understanding
1. Write down the symbols for the following elements:
a.sodium
e.iron
i.lead
b.silicon
f.manganese
j.phosphorus
c.sulfur
g.magnesium
k.fluorine
d.iodine
h.lithium
l.mercury
2. Write down the names of the elements to which the following symbols refer.
a.Cr
d.Cu
g.N
b.Cl
e.Ca
h.Na
c.C
f.Co
i.Ne
Check your answers.
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elements and the periodic table
1D
check your understanding
The following are incorrect symbols for the given element. Briefly state what is wrong and then
write the correct symbol.
1.
(silver)
5.
(sodium)
2.
(chlorine)
6.
(neon)
3.
(magnesium)
7.
(fluorine)
8.
(aluminium)
4.
(copper)
Check your answers.
key points
•• An element is the simplest kind of pure substance. It cannot be broken down by
chemical means into any simpler substance.
•• Each element is given a symbol. The first letter of the symbol is always upper case and
the second letter (if any) is always lower case.
•• Elements are placed in order of their atomic number in the Periodic Table.
•• The metallic elements are found on the left side of the Periodic Table and the nonmetals on the right side.
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structure of the atom
learning intention
In this lesson you will learn to:
•• describe the structure of the atom and explain the information given by the mass number and
the atomic number.
the atomic model
Elements are made of atoms. The atom is the building block of all chemistry yet is barely visible
under the most powerful microscope. How do we know what its structure is? Are all atoms the
same? Can one atom be changed into another kind of atom?
People have pondered these questions for the last few thousand years, from the alchemists who
have tried to turn lead into gold to the physicists trying to work out what happens when matter is
accelerated at tremendously high speeds.
The last 150 years have seen an enormous increase in our knowledge about this fundamental
building block. As new observations have been made, scientists have modified our idea of what
an atom probably looks like, to explain these observations. These explanations make up the
theory of the atomic model.
An alchemist carrying out
experiments.
gamzis
Public domain
2
Superconducting electromagnets used
for directing proton beams in the Hadron
Collider.
A brief summary of this theory is listed below.
1. All atoms are made up of smaller particles. The three main particles are protons, electrons
and neutrons.
2. Protons are positively charged particles held tightly together in a very small space in the
centre of the atom, called the nucleus.
3. If more than one proton is present, the atom needs neutrons to stop the protons repelling
each other and flying apart. These neutral particles are also found in the nucleus.
4. Electrons are negatively charged particles and are much lighter than protons or neutrons
(think of the mass of a bird compared with a horse).
5. Electrons take up most of the space of the atom surrounding the nucleus, due to the
repulsion between their negative charges.
6. The number of protons in an atom determines the type of element it is.
7. A neutral atom will have an equal number of protons and electrons.
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structure of the atom
Two representations of the atomic model for the element lithium are given below. These are not
to scale as they show the nucleus as relatively large. A better idea of the relative size of the
nucleus compared with the atom is to think of a cricket ball (nucleus) in the middle of a cricket
stadium (atom).
neutron
proton
nucleus
PhET
electron
This older model shows the
electrons in orbits around the
central nucleus. It is still very useful
to see these as energy levels.
online activity
Go to the Topic webpage for the link or download the Build an atom simulation (just Google
‘PhET build an atom’). This is a free download. If you cannot access the Internet, contact your
teacher to ask for the PhET CD.
Follow the instructions below before answering the questions. Sufficient information is given in
the diagrams to enable you to answer the questions that follow these instructions.
1. When you have downloaded the simulation, you
should see something like the diagram on the
right. The Build Atom tab on the top left should be
highlighted.
2. Click on the next to Symbol, Mass Number, and Net
Charge. Leave the model on ‘orbit’ and not ‘cloud’. Make
sure that the element name, neutral/ion and stable/
unstable are all ticked.
3. Now drag a proton from the proton basket into the
nucleus in the centre of the atom. Make a note of the
symbol, mass number and net charge.
4. Now drag an electron from the electron basket to
the inner orbit (blue dashed line). Make a note of the
symbol, mass number and net charge.
5. Now add a neutron from the neutron basket into the
nucleus at the centre of the atom. Make a note of the
symbol, mass number and net charge.
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PhET
2A
This model shows the electrons in
clouds around the central nucleus.
The position of the electrons cannot
be given exactly.
Information from step 3.
Information from step 4.
CH3001
9
structure of the atom
7. Keep on adding protons, neutrons and
electrons to the atom to see what happens.
When you feel confident, have a go at racing
yourself against the clock by clicking on the
Game tab.
PhET
6. Add a second proton from the neutron basket
into the nucleus at the centre of the atom.
Make a note of the symbol, mass number and
net charge.
Information from step 5.
Information from step 6.
QUESTIONS
1. Circle the subatomic particles which affect the charge on an atom.
electronsneutronsprotons
2. Circle the subatomic particles which affect the mass number of an atom.
electronsneutronsprotons
3. Circle the subatomic particles which affect the kind of element the atom belongs to.
electronsneutronsprotons
4. Look at the symbols shown for steps 3–6. What do you think the following parts represent?
a. Top left number
b. Top right number
c. Bottom left number
d. Central letter(s)
Check your answers.
mass number (a)
The activity above should have shown you that the mass number
gives the total number of protons and neutrons in an atom. This is
because both protons and neutrons are both comparatively heavy
and will affect the mass of the atom.
The mass number is represented by the symbol A which is written
on the top left side of the symbol for the element. The mass
number must be given to work out the number of neutrons in an
atom. Similarly, the number of neutrons must be given in order to
work out the mass number of an atom.
The mass number
is NOT the same as
atomic mass. The mass
number is always a
whole number as you
can’t get fractions of
protons or neutrons.
The mass number
is NOT given on the
Periodic Table, as it
does not relate to any
particular element.
The mass number always refers to an atom (or isotope). It never refers directly to an element, as
an element can have different mass numbers. For example, carbon can have a mass number of
12, 13 or 14. A mass number of 14 could refer to both 14C and 14N.
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structure of the atom
atomic number (z)
The number of protons always determines what kind of element
it is. The number of protons is also known as the atomic number
and is represented by the symbol Z. This number is written on the
bottom left of the symbol for the element.
So if alchemists can
remove three protons
from lead (82Pb), they
will finally have formed
gold (79Au)!
A given atomic number can only refer to one kind of element. For example, an atomic number
of 14 will only refer to silicon (14Si). If you change the number of protons, you will change the
element.
In summary, the mass number and the atomic number are written by convention in symbol form
as follows:
mass number atomic number
A
Z
E
symbol of element
Note: Elements are ordered in the Periodic Table according to increasing atomic number. It
is easy to get confused by the position of the atomic number in the conventional form and its
position in the Periodic Table. The Periodic Table does not follow convention, as it does not show
the mass number (it usually shows the relative atomic mass or molar mass, which is something
completely different). The atomic number is usually listed at the top or top left for each element
in the Periodic Table.
For example, compare the following:
32
16
S
Mass number never
has a decimal point.
Atomic number.
Conventional representation.
S Cl Ar
16
17
18
32.1
35.5
39.9
Atomic number.
Not the mass number,
as you can’t get 0.9 of
a neutron or proton.
Part of the Periodic Table (doesn’t use the conventional representation).
isotopes
Although an element is made of atoms having the same atomic number, atoms in an element are
NOT necessarily identical. This is because the number of neutrons in atoms of an element can
vary. Isotopes exist when atoms have the same atomic number but have different mass numbers.
For example, the element carbon, which has six protons (6C), can have six, seven or eight
neutrons in its atoms, giving the isotopes 126C, 136C and 146C.
Isotopes can be written in a number of ways. Carbon-12, C-12 and 126C all refer to the same isotope.
Notice how in each case both the element and the mass number are given.
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structure of the atom
online activity
Go to the Topic webpage for the link or download the simulation Isotopes and atomic mass (just
Google ‘PhET isotopes’). This is a free download. If you cannot access the Internet, contact your
teacher to ask for the PhET CD. You will use this simulation again in activity 9A.
Follow the instructions before answering the
questions.
1. When you have downloaded the simulation,
you should see something like the diagram
on the right. The Make Isotopes tab on the
top left should be highlighted.
2. Click on the next to Symbol, and
Abundance in Nature. Make sure that the
Mass Number is ticked (and not Atomic
Mass).
3. The default image shows 11H. See how many
protons and neutrons are in the nucleus and
check whether this isotope is stable. Find
out what percentage of hydrogen exists as
1
1H by looking at the Abundance in Nature
section.
4. Drag a neutron from the neutrons basket
to the nucleus of the hydrogen atom.
Make a note whether the isotope is stable
or unstable. Find out what percentage of
hydrogen exists as hydrogen-2.
5. Now drag another neutron to the nucleus.
Make a note whether the isotope is stable
or unstable. Find out what percentage
of hydrogen exists as hydrogen-3. What
happens when you add another neutron?
Information from step 3.
Information from step 4.
Information from step 5.
Information from step 6.
6. Now click on the element He on top left
of the Periodic Table on the top left of the
interactive. Make a note of whether the
default isotope is stable or unstable. What
happens when you add neutrons? What
happens when you remove neutrons by
dragging a neutron out of the nucleus back
into the basket?
PhET
2B
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structure of the atom
7. Find out how many isotopes each of the
first ten elements have. Click on each
element in turn and look at the Abundance
in Nature section. If it states 100%, then
only one isotope is found in nature. If it less
than 100%, find out what other isotopes
are present for each element by adding or
removing neutrons.
QUESTIONS
1. Which of the first 10 elements have only one
isotope?
2. Write the symbols for all the isotopes of
lithium.
PhET
8. Make a note of which elements have
naturally occurring unstable isotopes.
These isotopes are radioactive.
Some of the information you can find out in step 7.
3. Write the percentage abundance for each of the different isotopes of oxygen.
4. Which of the first 10 elements have unstable (radioactive) isotopes?
5. What do you notice about the relationship between the stability of the atom and the ratio of
protons to neutrons?
Check your answers.
In summary, the number of protons, electrons and neutrons in an atom can be calculated as
follows:
•• number of protons
= atomic number (Z)
•• number of electrons
= atomic number (for a neutral atom)
•• number of neutrons
= mass number (A) – atomic number (Z)
key points
•• The atom consists of positively charged protons, negatively charged electrons and
neutral neutrons.
•• The atomic number Z gives the number of protons in an atom. It is written as a subscript
on the left of the element symbol.
•• The mass number A gives the number of protons and neutrons in an atom. It is written
as a superscript on the left of an element symbol.
•• The Periodic Table does not show the mass number for an element. Note that the
atomic number is shown above the element in the Periodic Table, and not as a
subscript. (The molar masses, given on the Periodic Table, are explained in lesson 9.)
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3
electrons and energy levels
learning intention
In this lesson you will learn to:
•• write the electron configuration for an element and work out the number of valence electrons
for an element.
evidence for energy levels
In the previous lesson, you looked at an overview of the structure of the atom. Energy levels were
briefly mentioned. What are energy levels? How do we know they exist?
The evidence comes from atomic spectra. When light emitted from an element such as hydrogen
is looked at through a spectrometer, coloured lines like in the example shown below are seen.
Atomic spectrum of hydrogen.
Scientists wondered what caused these lines. To explain these lines, Niels Bohr eventually came
up with the idea of electrons being in fixed energy levels. His theory was modified further by
many other scientists, including Erwin Schrödinger, who introduced quantum mechanics.
The following activity is optional. It shows how our idea of the atom was changed to explain
atomic spectra.
optional online activity
Go to the Topic webpage to download the interactive spectrometer. Then download the PhET
simulation Isotopes and atomic mass from the Topic webpage. These are both free downloads.
(Or, just Google ‘PhET hydrogen atom’.) If you cannot access the Internet, contact your teacher
and ask for the PhET CD.
1. First use the interactive spectrometer to look at the
atomic spectrum for hydrogen. Do the lines match the
example given above?
2. Then download the PhET simulation on the models of
the hydrogen atom. You should see something similar
to the image shown on the right.
PhET
3A
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electrons and energy levels
3. Make sure the dial in the top left corner is
turned to Experiment and white is selected
for light controls. Move the slider at the
bottom to fast to minimise waiting times.
Click the red button above light controls to
insert the tube of hydrogen and then start
the experiment by clicking the button.
PhET
PhET
4. Make sure the spectrometer is showing and
click the start button. You should see the
lines starting to form. Does this match the
atomic spectrum for hydrogen shown on the
previous page?
Bohr model.
Schrödinger model.
5. Now turn the top left dial to Prediction. A list of the different atomic models should then be
shown. The Billiard ball model at the top is the oldest. The bottom model by Schrödinger is
the most recent. See what happens with each of these models.
QUESTIONS
1. Which models did not explain the atomic spectrum of hydrogen at all?
2. Which models gave a similar atomic spectrum to the experimental version?
3. Which model most closely matched its atomic spectrum to the one obtained by experiment?
Check your answers.
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electrons and energy levels
bohr’s theory
Niels Bohr explained the spectral lines by proposing that electrons could only have particular
energies. When an electron gained energy (for example, from a light photon), it jumped to a
higher energy level further from the nucleus. When the electron fell back to a lower energy level
closer to the nucleus, it emitted the extra energy as light. The amount of energy lost would
determine the wavelength of light.
e–
Electron absorbing energy
moves to a higher energy orbit.
e–
Electron moving to a lower
energy orbit emits energy.
electron configuration
The way the electrons are arranged in these energy levels is called the electron arrangement or
the electron configuration.
Each energy level can only hold up to a certain number of electrons.
The first, innermost energy level can only hold up to two electrons. The second energy level can
hold up to eight electrons. The number of electrons in each energy level can be worked out using
the formula 2n2, where n is the number of the energy level.
The table below compares the energy levels and the maximum number of electrons.
Energy level
Maximum number
of electrons
1st
2(1)2 =
2
2nd
2(2)2 =
8
3rd
2(3)2 =
18
3rd
2nd
1st
nucleus of atom
energy levels of electrons in an atom
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electrons and energy levels
To work out the way electrons are arranged you need to remember the following:
1. The atomic number of an atom tells you how many electrons it has.
2. Electrons fill the lowest energy level available before going into higher energy levels.
3. The first energy level can hold up to two electrons. The second energy level can hold up to
eight electrons. The third energy level can hold up to 18 electrons.
As an example, look at the electron arrangement for carbon.
1. The Periodic Table tells you carbon has an atomic number 6, so you know it
has six electrons.
2. These six electrons arrange by first filling the first energy level (two
electrons). The other four electrons will partly fill the second energy level.
3. You write the electron arrangement of carbon as 2, 4.
Electronic configurations will be explored in more detail in the booklet Periodic
trends (CH3041).
3B
Carbon
(six electrons).
optional online activity
For more help, visit the Topic webpage to use get useful tips on how to write electron
configurations.
3C
check your understanding
Make a copy of the table below and complete it for the first eleven elements. Refer to the Periodic
Table near the end of the booklet to find out the atomic number. Carbon is already filled in.
Element
Atomic number
Electron arrangement
6
2, 4
H
He
Li
Be
B
C
N
O
F
Ne
Na
Check your answers.
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17
electrons and energy levels
valence electrons
Electrons in the highest energy level (outermost shell) of an atom are called valence electrons.
For example, consider hydrogen, carbon and chlorine, which have 1, 4 and 7 valence electrons
respectively.
hydrogen
(one electron)
Hydrogen has one electron in
the first energy level. This is the
outermost or valence energy level,
so hydrogen has one valence
electron.
3D
carbon
(six electrons)
chlorine
(17 electrons)
Carbon has six electrons. Two are
in the first energy level and four are
in the second or valence energy
level, so carbon has four valence
electrons.
Chlorine has 17 electrons. Two
are in the first energy level and
eight are in the second energy
level, leaving seven in the third or
valence energy level. So chlorine
has seven valence electrons.
check your understanding
State the number of valence electrons for each of the following elements. Refer to the Periodic
Table near the end of this booklet to find out the atomic number.
1.hydrogen
2.lithium
3.sodium
4.beryllium
5.magnesium
6.boron
7.aluminium
8.nitrogen
9.phosphorus
10.oxygen
11.sulfur
12.fluorine
Check your answers.
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electrons and energy levels
Did you notice anything when completing activity 3C? You may have spotted that elements in the
same group (column) of the Periodic Table have the same number of valence electrons, except
for helium, which has only two valence electrons. In the next two lessons you will see how the
number of valence electrons affects the reactivity of the elements.
Number of valence electrons
He
12
Group 1 = 1 valence electron
21
Ca
Sc
38
39
Sr
Y
Zr Nb Mo Tc
Ru Rh Pd Ag Cd
56
57
72
76
Cs Ba
La
Hf
Ta
W
87
89
104
105
106
K
Rb
55
Fr
88
B
C
N
13
14
15
Al
20
37
8
Ra Ac
22
23
24
Ti
V
40
41
73
Cr Mn Fe Co
Ni
Cu Zn Ga Ge As
42
46
75
44
27
45
47
48
77
78
Re Os
Ir
Pt Au Hg
107
109
110
108
79
30
111
80
112
31
49
In
81
Tl
113
32
P
29
43
26
Si
28
74
25
50
33
51
Sn Sb
82
Pb
114
83
Bi
115
9
10
O
F
Ne
16
17
18
S
34
Se
52
Te
84
Po
116
Cl
35
Br
53
I
85
Group 18 = 8 valence electrons
Be
11
19
7
Group 17 = 7 valence electrons
Li
Na Mg
6
Group 16 = 6 valence electrons
5
Group 13 = 3 valence electrons
4
Group 2 = 2 valence electrons
3
Group 15 = 5 valence electrons
H
2
Group 14 = 4 valence electrons
1
Ar
36
Kr
54
Xe
86
At Rn
117
118
Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo
key points
•• Electrons are found in energy levels.
•• The first energy level closest to the nucleus can have up to two electrons.
•• The second energy level can have up to eight electrons.
•• The third energy level can have up to 18 electrons.
•• Electrons are added to the lowest energy level first.
•• The arrangement of electrons in the energy levels is called the electron configuration.
•• Valence electrons are found in the outermost energy level.
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4
bonding
learning intention
In this lesson you will learn to:
•• explain the three different ways elements become more stable (lose energy) by obtaining full
energy levels.
energy and stability
A quick examination of the reactivity of the elements will show that the elements in Group 18 are
very unreactive. It is thought that these elements are very stable and have a lower energy due to
having full energy levels.
By getting a full energy level, atoms become more stable as they lose energy. Atoms interact
with each other in order to get a full energy level. These interactions are called bonding.
In the next activity you will investigate what happens to the energy of an atom as it interacts with
another atom.
online activity
Go to the Topic webpage to download the Atomic interactions simulation. This is a free
download. (Or, just Google ‘PhET atomic interactions’.) If you cannot access the Internet, contact
your teacher to ask for the PhET CD.
Follow the instructions below before answering the questions. Sufficient information is given in
the diagrams to enable you to answer the questions that follow each set of instructions.
1. When you have downloaded the simulation, you should
see something like the diagram on the right. The
pinned atom will be neon with a yellow arrow pointing
to a second neon atom. Click on Show component
forces. Drag the second atom about 3 cm to the right
and watch what happens.
You should see the second atom start moving slowly
towards the pinned atom. Its speed increases as it gets
closer, because of the increasing strength of attraction
between the negatively charged electrons of the
second atom and the positively charged protons of the
pinned atom. When the second atom has moved too
close, the electrons from both atoms will repel each
other and the energy increases.
Questions
a. Why does the energy decrease as the atoms
approach each other and increase when the atoms
are very close?
PhET
4A
The o (sigma) sign indicates atom diameter. It
is the closest the atoms can get to each other
without the repulsion forces becoming too great.
b. Compare the component forces. Is the attraction
between the two neon atoms strong?
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bonding
2. Click on Adjustable Attraction. Move Atom Diameter (o) to
Small, and move Interaction Strength (ε) three quarters of
the distance from Weak to Strong. These settings model two
interacting hydrogen atoms.
Question
PhET
a. Why is there a much larger drop in energy between the two
hydrogen atoms when compared to the two helium atoms?
PhET
Information from step 2.
3. Move the moving atom so that the small blue circle is about
halfway up the potential energy curve. Release and observe the
behaviour of the moving atom with respect to the pinned atom. You should notice the atoms
vibrating back and forth.
Information from step 3.
Question
PhET
4. Move the moving atom closer to the pinned atom. What happens
to the potential energy of the H2 molecule? What happens to the
moving atom?
Information from step 4.
a. Why does the energy increase when the second atom is moved too
close to the pinned atom?
5. Drag the moving atom about 6 cm to the right, increasing the distance between the atoms.
What happens?
Question
PhET
a. Why is the moving atom not immediately attracted to the pinned atom?
Information from step 5.
6. Now move Atom Diameter (o) to halfway between Small and Large. Set the Interaction
Strength (ε) a quarter of the distance from Weak to Strong. These settings model two
interacting fluorine atoms.
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21
bonding
7. Drag the movable fluorine atom a little to the right. What do you notice?
Questions
a. Why do you think there is a much larger drop in energy between the two hydrogen atoms
when compared to the two fluorine atoms?
PhET
b. What can you say about the strength of the bond between two hydrogen atoms compared
to two fluorine atoms?
Information from steps 6 and 7.
8. Click on the oxygen-oxygen pinned atom/moving
atom pair. What do you notice about the drop in
energy when two oxygen atoms bond?
Drag the movable oxygen atom about 1 cm to the right.
What do you notice?
Questions
b. Based on what you have observed in this activity,
what do you think determines the strength of the
bond?
PhET
a. Based on the observed potential energy curve, rank
the strength of the oxygen-oxygen bond relative
to the fluorine-fluorine bond and the hydrogenhydrogen bond.
Information from step 8.
Check your answers.
bonding theory
The previous activity showed that atoms gain stability by losing some potential energy when they
are attracted to each other. The strength of these interactions varies from element to element.
Neon atoms form very weak interactions with each other which are not strong enough to be
considered a bond. Hydrogen, fluorine and oxygen, however, form much stronger attractions with
each other. These attractions are so strong that a lot of energy is required to move the atoms apart
again.
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bonding
We say that these atoms have bonded with each other. Bonding theory can be based on the idea that:
Atoms will share, gain or lose electrons to acquire the electron arrangement of a noble gas.
•• For most atoms this means acquiring eight electrons in the highest energy level.
•• Hydrogen, lithium and beryllium become stable with just two electrons in their highest
energy level.
•• Atoms can share, gain or lose electrons by forming chemical bonds.
There are three main types of bonding. These are covalent, ionic and metallic bonding. 1. covalent bonding
Non-metal atoms can share one or more of their valence electrons to get full valence energy
levels. Hydrogen and fluorine will share only one valence electron each, as they both only need
one extra electron to fill the valence energy level. Consider hydrogen below.
Valence energy
level requires one
more electron.
Electrons are shared,
so each energy level
has two electrons.
Oxygen requires two electrons to fill its valence energy level, so will share two of its valence
electrons. The electrostatic attraction between the shared electrons and the nuclei of the atoms
forms the covalent bond.
As no electrons are lost or gained, the molecules are neutral.
2. ionic bonding
Ions are formed when atoms gain or lose electrons. Elements on the left side of the Periodic
Table have fewer valence electrons. It requires less energy to remove these few electrons than it
does to gain many more electrons to get a full energy level.
For example, lithium has one valence electron. It requires less energy to lose this electron,
leaving a full first energy level as its outermost energy level, than to gain seven electrons to fill up
the second energy level.
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bonding
Li+ will now have a charge of +1 as it has three protons and only two electrons. A positively
charged cation (said cat-eye-on) is formed.
Valence electron
is removed.
The first energy level is
now the valence energy
level and is full.
Li
e–+Li+
Similarly, it requires less energy for fluorine to gain one more electron to fill the second energy
level, than to lose seven electrons. F– has a charge of –1 as it has nine protons and 10 electrons. A
negatively charged ion is formed, called an anion (said an-eye-on).
One electron
required to fill
valance energy level.
The second energy
level is now full.
F+e–
F–
As opposite charges attract each other by strong electrostatic
forces of attraction, each anion surrounds itself with cations and
each cation is surrounded by anions in a three-dimensional array or
lattice.
This electrostatic attraction between ions is called ionic bonding.
Ionic bonding occurs between a metal and a non-metal, with the
metal transferring its valence electrons to the non-metal.
Three-dimensional ionic lattice.
A way to remember that cations are positive and anions are negative is to think of the following
two pictures!
+++
Cation.
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Anion.
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bonding
3. metallic bonding
Metal elements have too few valence
electrons to share with each other, such as
occurs in covalent bonding. However, the
metal atoms move very close to each other,
allowing their valence electrons to move
freely from atom to atom. As the electrons
are not tied to any particular atom, they are
said to be delocalised.
e-
e
–
ee
e–
e-
ee
-
e-
e-
e-
-
e-
e-
e-
e
e
e-
e–
-
e–
-
electric field
An electric field applied to a metal lattice will
cause the delocalised electrons to move.
Metallic bonding can be described as an electrostatic attraction between the metal atom cores
and the sea of delocalised electrons.
using the periodic table
You can use the Periodic Table to help you identify what type of bonding will occur between
elements. Non-metals occur on the top right side of the Periodic Table. They will form covalent
bonds with each other and ionic bonds with metals.
Metals are on the left side of the Periodic Table. They will form metallic bonds with each other
and ionic bonds with non-metals.
Metalloids can differ in their bonding. You will examine bonding again in CH3042.
Ways of obtaining full energy levels
3
4
Li
Be
11
12
Na Mg
19
20
21
22
23
24
K
Ca
Sc
Ti
V
Cr Mn Fe Co
25
Ni
40
41
42
46
43
26
44
27
He
5
6
7share
8
9
10
Non-metals
B
C
N
O
F
Ne
electrons when bonding
13 with
14 each
15 other.
16
17
18
Al They
Si gain
P electrons
S Cl Ar
29
30
31
32 when
33
34
35
36
bonding
Cu Zn Ga Ge As toSe
Br
Kr
metals.
37
38
39
52
53
Rb
Sr
Sn Sb
Te
I
55
56
Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In
Metals
lose73electrons
when76bonding
with non-metals.
57
72
74
75
77
78
79
80
81
82
83
84
85
Cs Ba
La
Hf
Ta
W
Re Os
Ir
Pt Au Hg
Pb
Bi
Po
At Rn
87
88
89
104
105
106
107
109
110
Fr
Ra Ac
108
45
28
2
Group 18 elements are unreactive.
1
H
47
111
48
49
Tl
50
51
54
Xe
86
112
Rf Db Sg Bh Hs Mt Ds Rg Cu
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bonding
4B
check your understanding
Identify the type of bonding which would occur between the following elements.
1. hydrogen and oxygen
5. iron and iron
2. carbon and sulfur
6. magnesium and iodine
3. sodium and oxygen
7. calcium and fluorine
4. copper and chlorine
8. silicon and oxygen
Check your answers.
key points
•• Atoms with full outer energy levels are more stable.
•• Atoms can bond with other atoms to obtain full outer energy levels.
•• Atoms lose energy when they form bonds.
•• There are three types of bonding:
–– covalent bonding
–– ionic bonding
–– metallic bonding.
•• Covalent bonding involves the sharing of electrons between non-metal atoms. The
atoms are held together by the electrostatic forces between the shared electrons and
the nuclei of the atoms.
•• Ionic bonding involves the transfer of electrons from a metal to a non-metal atom to
form cations and anions. These ions are held together by electrostatic attractive forces.
•• Metallic bonding involves the attraction between delocalised valence electrons and the
positive core of metal atoms.
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5
valency and formulae
learning intention
In this lesson you will learn to:
•• use the Periodic Table to help work out valencies to write the formulae of compounds.
valency
O
One of the core concepts in chemistry is being able to write the formulae
for substances. For example, water is made up of molecules consisting of
two hydrogen atoms and one oxygen atom and so has a formula of H2O. How
can we work out this ratio? This ratio was originally worked out by careful
experimentation. Through these experiments, chemists were able to work
out the combining power of elements. Oxygen is able to combine with two hydrogen atoms, so
oxygen has twice the combining power of hydrogen. This ‘combining power’ is called valency.
H
H
The valency of an element is a number which represents its combining power with other
elements. Knowing the valency (or being able to work it out) is one of the key requirements to
write the formula of a compound.
valency and the periodic table
Chemists carried out many experiments to work out the combining power or valency of each
element. A pattern was found and so the valency of an element can usually be worked out using
the Periodic Table.
The valency of elements in Groups 1 and 2 is the same as the group number. The valency of
elements in Groups 14 to 17 can generally be worked out by subtracting the group number from
18. Elements in Group 13 usually have a valency of 3.
Group
Metal/non-metal
Valency
How to work out
the valency
Number
of
valence
electrons
1
metals
1
Same as group number = 1
1
2
metals
2
Same as group number = 2
2
13
metals
3
14
metals and non-metals
4
18 – group number = 4
4
15
mainly non-metals
3
18 – group number = 3
5
16
non-metals
2
18 – group number = 2
6
17
non-metals
1
18 – group number = 1
7
3
The number of valence electrons has also been listed in the table above for the sake of
comparison. You should notice that the values for valency and valence electrons is the same for
the first four groups, but differs for the last four groups.
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valency and formulae
elements with variable valencies
Some elements can have variable valencies. These elements tend to be found in the middle of
the Periodic Table, although exceptions do occur. In the names of compounds containing metals
with variable valency, the valency is given in Roman numerals directly after the element.
For example: iron(II) oxide, FeO, and iron(III) oxide, Fe2O3. Fe(II) and Fe(III) are called iron-two
and iron-three. Previously the endings –ous and –ic were used to show the different valencies for
an element. You may still read about ferrous oxide and ferric oxide.
Valency and number of valence electrons
5
6
7
Li
Be
B
C
N
1 valence electron. Valency = 1
11
2 valence electrons. Valency = 2
4
12
13
Na Mg
19
K
37
28
P
31
32
33
21
22
23
24
Sc
Ti
V
41
CrMayMn
Co valencies.
Ni Cu Zn Ga Ge As
haveFevariable
39
40
Y
Zr Nb Mo Tc
Ru Rh Pd Ag Cd
49
50
In
Sn Sb
56
57
72
73
74
76
77
78
Cs Ba
La
Hf
Ta
W
Re Os
Ir
Pt Au Hg
81
82
83
Tl
Pb
Bi
87
89
104
105
106
107
109
110
Fr
88
Ra Ac
75
108
46
30
38
55
45
29
Sr
Rb
44
27
15
20
43
26
14
Si
Ca
42
25
Al
47
48
79
80
51
8
O
16
S
34
Se
52
Te
84
Po
9
F
17
Cl
35
Br
53
I
85
8 valence electrons. Valency = 0
He
3
7 valence electrons. Valency = 1
H
2
6 valence electrons. Valency = 2
1
10
Ne
18
Ar
36
Kr
54
Xe
86
At Rn
111
Rf Db Sg Bh Hs Mt Ds Rg
groups of atoms
Some compounds contain groups of atoms which behave like a single atom of a metal or a nonmetal. Examples of these groups are hydroxide (OH–) and ammonium (NH4+). In a similar way to
metal and non-metal atoms, these groups have particular combining powers or valencies.
The tables below include the elements plus the groups of atoms you need to know. You should
learn all of these.
valencies of metals and groups acting like metals
1
Li
Na
K
Ag
Cu
NH4
28
2
lithium
sodium
potassium
silver
copper(I)
ammonium
CH3001
Mg
Ca
Fe
Cu
Zn
Pb
3
magnesium
calcium
iron(II)
copper(II)
zinc
lead
Al
Fe
aluminium
iron(III)
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valency and formulae
valencies of non-metals and groups acting like non-metals
1
2
H
Cl
Br
I
NO3
OH
HCO3
HSO4
MnO4
hydrogen
chlorine
bromine
iodine
nitrate
hydroxide
hydrogen carbonate
hydrogen sulfate
permanganate
O
S
CO3
SO4
CrO4
Cr2O7
3
oxygen
sulfur
carbonate
sulfate
chromate
dichromate
N
P
PO4
4
nitrogen
phosphorus
phosphate
C
Si
carbon
silicon
Note: the common name for hydrogen carbonate is bicarbonate. For example: sodium bicarbonate may be listed as the ingredient in
baking soda. The chemical name would be sodium hydrogen carbonate.
formulae for compounds
Why use formulae? A formula provides a convenient shorthand for representing:
•• the elements present in a compound
•• the relative proportions of these elements.
Writing formulae of compounds is easy if you learn and use the valencies given in the tables
above.
Example: sodium oxide
Sodium has a valency of 1, and oxygen has a valency of 2. This tells you that each oxygen atom will
combine with two sodium atoms and so the formula is Na2O. This is shown stepwise below.
1. You may find it useful to visualise the valency of each element as
a ‘hook’. Draw a circle with one hook to represent Na and a circle
with two hooks to represent O. These are shown on the right.
2. The sodium ion has only one hook. What ion could you add to
the drawing to hook onto the second hook of the oxygen ion?
Did you guess another sodium ion, as shown in the diagram on
the right?
Now check whether the number of hooks is the same on both
sides. Yes, the two hooks of sodium are balanced by the two
hooks of oxygen.
Remove the circles and hooks and the formula is Na2O (sodium
oxide).
Notice where the 2 goes to show two sodium ions. It is written at the
foot of the Na letter symbol on the right side. Notice also that the
formula Na2O has no charge.
© te ah o o t e k ur a p o un a m u
Na
O
Step 1.
Na
O
Na
Step 2.
Na2O
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valency and formulae
helpful hints
•• When a formula contains a metal and a non-metal, the metal is always written first.
•• When a formula contains two non-metals, usually the one with the higher valency is written
first, for example, CO2, SO3, NH3, PCl5. Water (H2O) is one exception here.
•• A useful rule to check a formula is:
number
of atoms
of first
element
Example: Na2O:
×
valency
number
of atoms
of second
element
=
×
valency
(2) × 1 = (1) × 2 (the numbers of atoms are given in brackets)
•• The most common valency for copper is 2. Unless a valency of 1 is indicated, for any copper
compound assume that the metal has a valency of 2. For example, copper sulfate means
copper(II) sulfate, CuSO4.
5A
check your understanding
Write the formulae for compounds formed by the following pairs of elements. To help you do
this, use the valency tables given earlier in the lesson.
1. sodium and chlorine
6. carbon and hydrogen
2. calcium and oxygen
7. hydrogen and sulfur
3. magnesium and bromine
8. aluminium and oxygen
4. phosphorus and hydrogen
9. silicon and oxygen
5. iron(III) and chlorine
10. hydrogen and chlorine
Check your answers.
formulae which use brackets
When writing the formulae of compounds containing groups of atoms (or polyatomic ions), you
may find it necessary to use brackets. An example is the formula for aluminium sulfate.
Work out this formula now, using the valency tables.
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valency and formulae
Your formula should show three sulfate groups combined with one aluminium atom. This is
written Al2(SO4)3.
You may have written Al2SO43 instead. This is incorrect, as it represents:
2 aluminium atoms : 1 sulfur atom : 43 oxygen atoms.
It should be:
2 aluminium atoms : 3 sulfur atoms : 12 oxygen atoms.
In other words, you need a bracket around the SO4 to show 3 × (SO4).
Example: Aluminium sulfate
This is shown stepwise below.
2. The aluminium ion has an extra hook. Add another sulfate. Keep
on doing this until all the hooks have been joined together.
SO 4
Al
1. Draw a circle with three hooks to represent Al and a circle with
two hooks to represent SO4. These are shown on the right.
Step 1.
Remove the circles and hooks and the formula is
SO 4
Al
Al2(SO4)3
SO4
Notice how the 2 showing two aluminium atoms is written at the foot
of the Al letter symbol on the right side. The 3 indicating three
sulfate groups is written outside the brackets on the right side. This
number applies to all the atoms inside the brackets.
The Al stands
for the element
aluminium.
Al
SO
4
Step 2.
The brackets enclose
all the atoms in the
sulfate group.
Al2(SO4)3
The 2 means that
there are two Al
atoms in the formula.
There is no small number
after S. It means that
there is only one sulfur
atom in the group.
The 3 outside the
brackets means
that there are three
groups of SO4.
The 4 means that
there are four oxygen
atoms in the sulfate
polyatomic ion.
Brackets are used whenever more than one of a particular polyatomic ion is used. For example,
brackets are also used in the formula for lead hydroxide, Pb(OH)2.
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valency and formulae
Compare the formulae when the brackets are put in and when they are left out:
Pb(OH)2 one Pb : two O : two hydrogen 
PbOH2 one Pb : one O : two hydrogen  (as incorrect number of oxygen)
5B
check your understanding
Write formulae for compounds containing the following atoms and/or groups. Use the valency
tables to help. Remember to use brackets if you want to show more than one group in the formula.
1. sodium and hydroxide
6. calcium and hydrogen carbonate
2. magnesium and nitrate
7. ammonium and carbonate
3. iron(II) and hydroxide
8. hydrogen and sulfate
4. sodium and sulfate
9. aluminium and sulfate
5. zinc and carbonate
10. calcium and phosphate
Check your answers.
information from formulae of compounds
For molecular substances, the formula tells you the exact number of atoms in one molecule of
the substance.
For other substances, the formula tells you the simplest whole number ratio of atoms of each
element present in a compound. This is sometimes referred to as the ‘empirical’ formula or one
‘formula unit’ of the substance.
For example:
The formula for silica is SiO2. (This occurs commonly as white sand.) The formula represents one
silicon atom combined with two oxygen atoms. It tells you that in any sample of silica, the ratio of
the numbers of atoms is:
atoms of silicon : atoms of oxygen = 1 : 2
One formula unit of silica contains one atom of silicon and two atoms of oxygen.
The formula for magnesium nitrate is Mg(NO3)2. This formula represents one magnesium atom
and two nitrate groups. In any sample of this compound the ratio of atoms is:
atoms of magnesium : atoms of nitrogen : atoms of oxygen = 1 : 2 : 6
One formula unit of magnesium nitrate contains one atom of magnesium, two atoms of nitrogen,
and six atoms of oxygen.
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valency and formulae
5C
check your understanding
For each of the formulae below, write the names of the elements present and give the ratio of the
numbers of their atoms in a sample of the compound.
1.FeCl3
4.Cu(OH)2
2.Al2O3
5.K2O
3.Na2CO3
6.Ca3(PO4)2
Check your answers.
5D
optional online activity
Visit the Topic webpage to use interactive animations to help you learn the groups and write
formulae.
key points
•• The valency of an element is the combining power of the element.
•• The valency can be worked out from the formula of a compound or from the Periodic
Table.
•• Some elements can have variable valencies. The valency is written as a Roman numeral
in brackets after the element name.
•• The formula of a compound gives the ratio in which elements combine.
•• The number of the atoms is written as a subscript on the right of the element symbol.
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6
naming compounds
learning intention
In this lesson you will learn to:
•• name compounds, given their chemical formula.
naming ionic compounds
Only a few rules are needed to name ionic compounds. These are shown in the table below.
Rule
Example
1. The name of the metal or ammonium
group is written first.
Sodium chloride, magnesium oxide,
ammonium nitrate.
2. If the metal has more than one valency,
the appropriate valency is indicated by
a Roman numeral in brackets after the
metal.
Fe(OH)2 is called iron(II) hydroxide (‘iron-two’
hydroxide).
3. The name of the non-metal part of the
compound is written second.
If this is an element (rather than a group)
then the ending of the name is changed
to –ide.
Non-metal groups containing oxygen
usually end in –ate or –ite.
6A
Fe(OH)3 is iron(III) hydroxide.
NaCl is called sodium chloride (not sodium
chlorine).
MgO is called magnesium oxide (not
magnesium oxygen).
NH4NO3 is called ammonium nitrate. The
ending has not changed, because nitrate is a
group.
check your understanding
Write the names for these 10 compounds. Use the tables in lesson 5 if you can’t remember the
names for the symbols.
1. KCl
6. (NH4)2SO4
2. MgI2
7. Na2CO3
3. FeO
8. Ag3PO4
4. FeCl3
9. Mg(HCO3)2
5. PbS
10. CaC2
Check your answers.
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naming compounds
naming covalent compounds and molecules
Non-metals may also show variable valency. A different system of naming is used for compounds
with non-metal atoms only. A prefix is written in front of the element.
Prefix
6B
Example
1. mono- or mon-
as in monoxide, meaning one oxygen
2. di-
as in dioxide, meaning two oxygen atoms
3. tri-
as in trichloride, meaning three chlorine atoms
4. tetra-
as in tetrachloride, meaning four chlorine atoms
5. penta- or pent-
as in pentoxide, meaning five oxygen atoms
check your understanding
Write the names for these eight compounds.
1.NO2
5.SiCl4
2.SO2
6.OF2
3.SO3
7.Cl2O
4.CS2
8.ClO2
Check your answers.
naming guidelines
Here are some more guidelines that can help you make sense of the names of compounds you
may come across.
1. Simple anions (only one element in the anion) end in ‘ide’. The prefix is derived from the name
of the element.
For example: Cl– = chloride (chlorine + ide)
O2– = oxide(oxygen + ide)
S2– = sulfide (sulfur + ide)
Some polyatomic ions (more than one element in the anion) can also end in –ide. You can
recognise these as the prefix is not derived from the name of an element.
For example: OH– = hydroxide CN– = cyanide
© te ah o o t e k ur a p o un a m u
(? + ide)
(? + ide)
CH3001
35
naming compounds
2. Anions ending in ‘ate’ or ‘ite’ contain oxygen. An ‘ate’ group always has more oxygen that an
‘ite’ group.
For example:SO42– = sulfate SO32– = sulfite NO3 = nitrate NO2– = nitrite
3. The prefix ‘per’ is used if more oxygen atoms are present than usual or if the element has a
higher oxidation state (see CH3071). The group name will always end in ‘ate’.
For example: ClO3– = chlorate
MnO42– = manganate
ClO4– = perchlorate
MnO4– = permanganate
4. The prefix ‘hypo’ is used if fewer oxygen atoms are present than usual. The group name will
always end in ‘ite’.
For example: ClO2– = chlorite
ClO– = hypochlorite
5. When oxygen is replaced with sulfur, the word ‘thio’ is used.
For example:SO42– = sulfate S2O32– = thiosulfate
OCN– = cyanate SCN– = thiocyanate
6. When an anion joins with H+ to form an acid, then ‘ide’ or ‘ate’ becomes ‘ic acid’ and ‘ite’
becomes ‘ous acid’.
For example:HNO3 = nitric acidHNO2 = nitrous acid
H2SO4 = sulfuric acid H2SO3 = sulfurous acid
HCl = hydrochloric acid
7. Names of common acids and their anions.
acid tartaric acid aniontartrate
citric acidcitrate
ethanoic acidethanoate
6C
check your understanding
Match the names with the formulae given below.
1.Cl–
perchlorate
6.S2–
sulfate
2.ClO–
chloride
7.SO32–
thiosulfate
3.ClO2–
chlorite
8.SO42–
sulfide
4.ClO3–
hypochlorite
9.S2O32–
sulfite
5.ClO4–
chlorate
Check your answers.
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naming compounds
key points
•• For compounds containing both metal and non-metal
elements:
–– The valency of metal elements having a variable
valency is written as Roman numerals in brackets
after the metal name.
–– If there is only one type of non-metal element, its
name ends in –ide.
Delene Holm
–– The metal part is usually written first.
ATE a dOnut
(One way to remember that a group
ending in –ate contains oxygen!)
–– Non-metal groups containing oxygen usually end in
–ate or –ite.
•• For compounds containing only non-metal elements, the prefixes mon–, di– etc. are
used.
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CH3001
37
chemical reactions
learning intention
In this lesson you will learn to:
•• describe a chemical reaction at both the macroscopic and submicroscopic level.
chemical change
Chemical reactions are happening all around you as well as inside you. Some are slow, some are
fast, but they all involve changes where new substances are formed.
We can view a chemical reaction from two levels: 1. the macroscopic level
2.
the submicroscopic level.
1.
macroscopic level
Describing a reaction from a macroscopic perspective includes our observations of the reaction.
For example, when we burn magnesium ribbon in air we see that the shiny magnesium ribbon
needed some heat to start the reaction. However, once the reaction had started it gave off a very
bright light and left a white powder.
Don Laing
7
Magnesium ribbon.
Magnesium burning in air.
White magnesium oxide formed.
When you are asked to write your observations for a reaction, you should record everything that
happens – what you see, what you hear, what you smell, and how the substances are different at
the end from what you had at the start.
You can then use your observations from the reactions to write a word equation.
word equations
A word equation describes a chemical equation by stating what chemicals reacted and what
products were formed. Whether each chemical was as a solid, liquid or gas can also be included
in the equation.
magnesium
(solid)
+
Reactants
magnesium oxide
(solid)
oxygen
(gas)
The arrow shows
the direction of the
reaction.
Products
The substances on the left of the arrow are called reactants. Substances on the right are called
products. Note that an arrow is used instead of an equals sign. This is because the products are
not the same as the reactants.
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chemical reactions
Here is another example. Consider the reaction between silver nitrate and
sodium iodide solutions.
Colourless solutions of silver nitrate solution and sodium iodide solution were
mixed together to give a pale yellow precipitate (solid) of silver iodide. The
colourless sodium and nitrate ions remained in solution.
The word equation is written as:
silver nitrate
+
sodium iodide
(aqueous solution) (aqueous solution)
Don Laing
Our observations would be as follows:
Products of the
reaction between
silver nitrate and
sodium iodide.
silver iodide
sodium nitrate
(solid)
(aqueous solution)
predicting precipitates
You may be wondering how you can work out what substance has been formed. It is possible to
tell which solutions will give a precipitate when mixed, using the following solubility rules.
Nitrates (NO3–)
Group 1 (Na+, K+)
Ammonium (NH4+)
All soluble.
Chlorides (Cl–)
Bromides (Br–)
Iodides (I–)
All soluble except Ag, Hg(I)
and Pb salts.
Oxides (O2–)
All insoluble except Group 1
and ammonium salts.
Sulfides (S2–)
All insoluble except Group 1
and Group 2 and ammonium
salts.
Sulfates (SO42–)
All soluble except BaSO4,
CaSO4, PbSO4.
Hydroxides (OH–)
Carbonates (CO32–)
All insoluble except Group 1
and ammonium salts.
© te ah o o t e k ur a p o un a m u
Knowing the
solubility rules will
help you in your
study of chemistry.
CH3001
39
chemical reactions
1
4
Be
37
21
22
23
24
25
Ca
Sc
Ti
V
Cr Mn Fe Co
Ni
38
39
40
41
42
Sr
Y
27
28
56
6
7
10
N
INS
O
9
C
8
F
Ne
13
14
15
16
17
18
Ar
35
36
Br
Kr
53
54
Si
INS
P
S Cl
31
32
33
34
Cu Zn Ga Ge As
Se
29
30
45
46
Zr Nb Mo Tc
47
48
57
72
73
74
78
Cs Ba
La
Hf
Ta
79
80
87
88
89
104
105
Fr
Ra Ac
111
55
43
26
5
B
Al
44
Rb
INS
Ru Rh Pd Ag Cd
49
50
In
Sn Sb
Te
I
Xe
76
77
W
Re Os
Ir
81
82
INS
INS
Pt Au Hg Tl Pb Bi
84
85
86
Po
At Rn
106
107
109
110
75
108
51
83
52
Rf Db Sg Bh Hs Mt Ds Rg
check your understanding
The photo shows the results of three chemical reactions. Describe these reactions on the
macroscopic level. Write down (a) your observations and (b) a word equation for each reaction.
Use the solubility rules on the previous page to help you work out the precipitate.
1. Colourless solutions of lead nitrate solution and
sodium iodide solution were mixed together.
2. Colourless solutions of lead nitrate solution and
sodium chloride solution were mixed together.
3. Colourless solutions of lead nitrate solution and
sodium bromide solution were mixed together.
Don Laing
7A
oxides mostly insoluble
sulfides mostly insoluble
insoluble halide
insoluble sulfate
20
mostly insoluble
soluble
K
=
=
=
=
He
mostly soluble
Key:
11
12
Na Mg 3
Li
19
2
Solubility
H
Check your answers.
Left: Reaction between Pb(NO3)2 and NaI.
Middle: Reaction between Pb(NO3)2 and NaCl.
Right: Reaction between Pb(NO3)2 and NaBr.
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chemical reactions
2.
submicroscopic level
Describing a reaction from a submicroscopic perspective includes thinking about what happens
at particle level. Models of atoms and bonding theory are used to try to explain observations that
have taken place at the macroscopic level.
For example, burning magnesium ribbon in air can be described at the particle level as follows.
Mg
Mg
Mg
Mg
Heat released
allows more
oxygen bonds to be
broken (explains
why the reaction
continues).
Mg
Mg
Mg
Mg
O
Mg
O
Electrons are transferred from magnesium
to oxygen (gives a reason why the reaction is
occurring – stability gained from full energy levels).
Mg
Mg
O
Mg
Mg
Mg
O
Mg
Energy is needed to break the
bond in the oygyen molecule
(explains why you need to
light the magnesium ribbon to
start the reaction).
Magnesium oxide is formed (explains the presence
of the white powder).
O
A lot of energy is released when MgO ionic bonds
are formed (explains the bright light).
Another way of describing a chemical reaction at the submicroscopic level is by using a symbol
equation. Symbol equations must be balanced. This means that the number of atoms of
elements on the left-hand side must be the same as the number of atoms on the right-hand side,
as matter cannot be created or destroyed. (How to balance symbol equations is discussed in the
next lesson.) The above reaction can be described as follows:
2Mg(s)
+
O2(g)
2MgO(s)
The (s) in the equation above is the symbol indicating that the substance is a solid. Similarly, (g)
indicates a gas, (l) indicates a liquid and (aq) indicates that the substance is dissolved in water
(aqueous solution).
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CH3001
41
chemical reactions
A symbol equation gives different information to a word equation. A comparison is given below:
Word equation
Symbol equation
•• Names the reactants.
•• Gives the formulae for the reactants.
•• Names the products.
•• Gives the formulae for the products.
•• States can be included.
•• States can be included.
•• Gives the ratio in which the reactants
react.
•• Gives the ratio in which the products are
formed.
We can assume that if we react substances on a macroscopic scale, all the atoms in the
reactants will react in the same way and in the same ratios.
Let’s consider the reaction between silver nitrate and sodium iodide. When each of silver nitrate
and sodium iodide are dissolved in water, the ions are separated from each other and become
surrounded by water molecules. We say that each ion is hydrated.
W
W
W
Ag
W
+
W
W
W
NO3–
W
W
Hydrated silver and nitrate ions.
Key:
42
W
W
W
W
Na+
W
W
W
W
W
W
I–
W
W
Hydrated sodium and iodide ions.
= water molecule
CH3001
© te ah o o te k u ra p ou n a mu
chemical reactions
The diagram below shows what happens when the two solutions are mixed. The water molecules
have been removed for clarity.
AgNO3 is soluble so the ions are separated
and spread out in the solution (explains
why we can’t see the silver nitrate and also
why the reaction is quick as no energy is
required to break ionic bonds).
W
W
Ag+
W
W
W
W
W
NO3_
W
W
W
Ag+
NO3_
I–
NO3_
Na+
Na+
Soluble Na+ and NO3– ions remain spread out
in the solution (explains why we can’t see the
sodium nitrate).
Na+
Insoluble AgI
precipitates out of
solution (explains the
pale yellow precipitate).
NO3_
Ag+
I–
Ag+
I–
Ag+
I–
Na+
NO3_
The symbol equation for the above reaction can be described as follows:
AgNO3(aq)
NaI(aq)
AgI(s)
+
NaNO3(aq)
check your understanding
The photo on the right shows the reaction between a marble chip
(calcium carbonate) and hydrochloric acid to form the calcium chloride
salt*, water and carbon dioxide gas. Use the solubility rules to work out
whether calcium chloride is soluble or not.
don laing
7B
+
*Salt is the general name given to the product between an acid and a base. It is made up of the
metal cation from the base and the non-metal anion from the acid.
1. Describe your observations for this reaction, based on the photo.
2. Write the word equation for this reaction, including states.
3. Describe what is happening at particle level.
4. Write a symbol equation for this reaction, including states.
Check your answers.
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CH3001
43
chemical reactions
7C
check your understanding
An iron nail reacts with copper sulfate, forming iron(II) sulfate and
copper metal. Use the solubility rules to work out whether iron(II)
sulfate is soluble or not.
1. Describe your observations for this reaction, based on the
photos.
don laing
2. Write the word equation for this reaction, including states.
3. Describe what is happening at particle level.
4. Write a symbol equation for this reaction, including states.
Initial.
Later.
Check your answers.
ionic equations
In the example previously given, when silver nitrate was reacted with sodium iodide, the
following symbol equation could be written.
AgNO3(aq)
+
NaI(aq)
AgI(s)
+
NaNO3(aq)
As each of the aqueous species are hydrated and so separated from each other, the equation can
also be written as follows.
Ag+(aq) + NO3–(aq) + Na+(aq) +
I–(aq)
AgI(s) + Na+(aq) + NaNO3–(aq)
Comparing the left and right side of the equation, you should notice that Na+(aq) and NO3–(aq) appear
on both sides of the equation. This means that they are unchanged and do not react. They are said to
be spectator ions. As spectator ions do not take part in the reaction they can be left out.
Ag+(aq) + NO3–(aq) + Na+(aq) +
This leaves only an ionic equation.
7D
I–(aq)
AgI(s) + Na+(aq) + NaNO3–(aq)
Ag+(aq) +I–(aq)
AgI(s)
check your understanding
Write the following reactions as ionic equations. Note that acids dissociate into H+(aq) and its
aqueous anion.
1. AgNO3(aq)
+
NaBr(aq)
AgBr(s)
+
NaNO3(aq)
2. CuSO4(aq)
+
Fe(s)
Cu(s)
+
FeSO4(aq)
3. Na2CO3(aq)
+
MgCl2(aq)
MgCO3(s)
+
2NaCl(aq)
4. 2HCl(aq)
+
CaCO3(s)
CaCl2(aq)
+
H2O(l)
+
CO2(g)
Check your answers.
44
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chemical reactions
7E
optional online activity
Visit the Topic webpage to use interactive animations describing different types of reactions
at the micro and macro level. There are also some animations showing how ionic equations are
written.
key points
•• Chemical reactions can be described at a macroscopic and a submicroscopic level.
•• Word equations and observations such as what can be seen, heard and felt, are
examples of descriptions at a macroscopic level.
•• Explanations using models such as the atomic model, bonding models, symbol
equations and particles are examples of descriptions at a submicroscopic level.
•• A symbol equation gives the symbol or formula of each substance and the state of each
substance in a reaction.
•• A symbol equation shows what happens to the individual atoms of each element and/or
compound and the numbers in which they combine.
•• We assume that what happens to a few atoms will also happen to a large number of
atoms as well (substances on a macroscopic scale), and in the same proportions of
atoms.
•• Ionic equations can be written that leave out the unreacting particles, called the
spectator ions.
© te ah o o t e k ur a p o un a m u
CH3001
45
balancing symbol equations
learning intention
In this lesson you will learn to:
•• balance symbol equations for chemical reactions.
rearranging atoms
don laing
When a chemical reaction happens, we explain the reaction on
the submicroscopic level as the rearrangement of the atoms in the
reactants to form the products.
For example, consider the reaction between magnesium and dilute
hydrochloric acid. On a macroscopic level we can observe bubbles of
hydrogen gas (which can be tested by the loud ‘pop’ sound it makes
when held near a burning splinter). When the reaction stops, a solution
of magnesium chloride is left behind.
Magnesium ribbon reacting
with hydrochloric acid,
producing hydrogen gas.
The chloride which was bonded to the hydrogen in hydrochloric acid is
now bonded to the magnesium.
don laing
8
The hydrogen which was bonded to the chloride in hydrochloric acid is
now bonded to itself.
A burning splint making a
‘pop’ sound is the test for
hydrogen gas.
The word equation is given below:
magnesium
(solid)
+ hydrochloric acid
(solution)
magnesium chloride +
(solution)
hydrogen
(gas)
The symbol equation can be written as:
Mg(s)
+
HCl(aq)
MgCl2(aq)
+
H2(g)
If you look carefully at the above equation, you will notice that there is one more chloride and
one more hydrogen atom that have been formed. One of the core concepts of chemistry is that
matter (or atoms) can’t be made or destroyed. So something is wrong!
We can overcome this problem by adding an extra HCl as a reactant. This is called balancing the
equation.
Mg(s)
+
HCl(aq)
HCl(aq)
MgCl2(aq)
+
H2(g)
+
H2(g)
Instead of writing HCl twice, we can write a 2 in front of the HCl.
Mg(s)
+
2HCl(aq)
MgCl2(aq)
There are now equal numbers of atoms for each element on both sides of the equation. No atoms
were made or destroyed during the reaction.
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balancing symbol equations
The kind of reactions that you normally carry out (what you might call ‘beaker’ or ‘test tube’ reactions)
cannot create or destroy atoms. Atoms can’t appear from nowhere, or disappear into nothing. What
your reactions can do is rearrange atoms, that is, change the way atoms are combined.
If an equation is to tell the whole story, it must trace where every atom in the reactants goes to,
and where every atom in the products comes from.
Note that you cannot change a formula in order to balance the equation.
For example, it would be wrong to alter the equation from the previous page like this:
Mg(s)
+
H2Cl2(aq)
MgCl2(aq)
+
H2(g)
There is no compound with a formula of H2Cl2. Experimental evidence has shown that only one
hydrogen atom combines with one chlorine atom to form HCl.
Here are some simple steps to balance an equation.
1. Check that all the formulae are correct. If they are, don’t change them.
2. Count the number of atoms of each element on either side of the equation.
3. Look at the elements where the numbers don’t match.
4. Balance the equation by putting a number in front of a formula. (You may need to try more
than one number – and your first number may not be correct.)
5. Remember, a number in front of a formula affects every element in the formula.
6. Keep the numbers as low as possible, and use only whole numbers (no fractions).
7. A final hint – balance single elements (for example H2, Mg, O2) last of all.
8A
online activity
Go to the Topic webpage to download the Balancing chemical equations simulation from PhET
website. This is a free download. (Or, just Google ‘PhET balancing chemical’.) If you cannot access
the Internet, contact your teacher to ask for the PhET CD.
Follow the instructions below before answering the
questions. Sufficient information is given in the diagrams to
enable you to answer these questions.
1. When you have downloaded the simulation, you should
see something like the diagram on the right. The
Introduction tab on the top left and Make Ammonia
should be highlighted. Click on Bar Charts.
© te ah o o t e k ur a p o un a m u
CH3001
47
balancing symbol equations
2. Change each of the values in the equation to 1. Look at
the reacting molecules and products as well as on the
bar chart showing the number of atoms of each element.
a. How many atoms of nitrogen and hydrogen are on
each side of the equation?
PhET
Questions
Information from step 2.
b. Have any atoms disappeared or created to form the products?
3. Change the value in front of NH3 from 1 to 2 to balance the
nitrogen atoms. Are the hydrogen atoms balanced? You
should notice that you have two hydrogen atoms in the
reactants but have six hydrogen atoms in the products.
Change the value in front of H2 until the bar chart shows that
the equation is balanced.
Question
a. Write the balanced equation for the reaction between
nitrogen gas and hydrogen gas to form ammonia.
Information from step 3.
4. Now click on Separate Water. Change the values in front of
H2O, H2 and O2 until the equation is balanced.
Question
a. Write the balanced equation for the decomposition of H2O
to form H2 and O2.
5. Now click on Combust Methane. Change the values in front
of CH4, O2, CO2 and H2O until the equation is balanced.
Question
Information from step 4.
Information from step 5.
a. Write the balanced equation for the combustion of CH4
with O2 to form CO2 and H2O.
6. Now click on Balancing Game. Choose your level and decide
whether you want the clock and sound on or off.
when you think that your equation is balanced.
7. Click on
You will be told if you are correct. If you are, click
.If you
are incorrect, you can choose to be shown why and/or just
.
Question
CH4
CH4
SO2
N2
CO2
48
+
+
+
+
+
CH3001
H2O
S
H2
H2O
H2O
H2
CS2
S
NH3
C2H2
+
+
+
+
+
CO
H2S
H2O
O2
O2
PhET
a. If you do not have access to a computer, balance these
equations below.
Information from steps 6 and 7.
© te ah o o te k u ra p ou n a mu
balancing symbol equations
Check your answers.
8B
check your understanding
Balance each of the following symbol equations, firstly as a full equation, and then as an ionic equation.
1.
2.
3.
4.
ZnCl2(aq)
Zn(aq)
Ca(OH)2(s)
Al2O3(s)
+ Na2CO3(aq)
+
HCl(aq)
+ HNO3(aq)
+
HCl(aq)
ZnCO3(s)
ZnCl2(aq)
Ca(NO3)2(aq)
AlCl3(aq)
+
+
+
+
NaCl(aq)
H2(g)
H2O(l)
H2O(l)
Check your answers.
8C
optional online activity
Visit the Topic webpage to use interactive animations to help you balance equations.
key points
•• An equation must account for every atom of every element. In a symbol/formula
equation, there must be the same number of atoms for every element in the reactants as
there are in the products.
•• Adjusting an equation to achieve this is called balancing the equation.
•• Equations are balanced by placing numbers in front of one or more formulae (not by
changing a formula).
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CH3001
49
9
the mole
learning intention
In this lesson you will learn to:
•• calculate relative, molecular and formula masses and link these to a mole of a substance.
relative masses
In the previous two lessons, you considered chemical reactions on both the macroscopic and
microscopic levels. You also learnt about the importance of balancing symbol equations, as
matter cannot be created or destroyed.
But how does a chemist know how much of each reactant to mix together? Atoms are too small to
count so a chemist can’t count out twice as many hydrogen atoms to oxygen atoms to make water!
Chemicals (even gases) can usually be weighed, so the mass of atoms of one chemical can be
compared with the mass of atoms of another chemical. One way of comparing atoms is to use a mass
spectrometer. You will learn more about analysing chemicals using a mass spectrometer in CH3021.
In a mass spectrometer, a gaseous sample of an element is passed through an electron beam producing
charged atoms (ions). These are accelerated through a magnetic field which deflects them, depending on
their mass and charge, onto a detecting screen.
Although the element hydrogen was originally used as the basis for
comparing the mass of atoms, carbon is now used as the standard. An
atom of carbon-12 is assigned a relative atomic mass (ram) value of
12.00. All other elements are compared to the mass of carbon-12.
Just as you might say you were twice as tall as your little brother,
relative atomic masses are comparisons or ratios. The terms
‘twice’, ‘three times’ and so on, have no units.
Relative atomic mass (ram) has no units.
50
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© te ah o o te k u ra p ou n a mu
the mole
Here are some common ram values in order of increasing mass. You do not have to learn ram
values, as they will always be given.
Element
Symbol
Relative atomic mass
(ram)
O
hydrogen
H
1.0
carbon
C
12.0
nitrogen
N
14.0
oxygen
O
16.0
magnesium
Mg
24.3
chlorine
Cl
35.5
16.0
N
14.0
O
16.0
C
12.0
N
14.0
O
16.0
C
N
Notice that the relative atomic masses don’t have to be whole numbers. This is because an
average is taken of all the naturally occurring different masses of the isotopes of an element. You
will investigate this in the next activity.
online activity
Go to the Topic webpage to download the Isotopes and atomic mass simulation. This is a free
download. (Or, just Google ‘PhET isotopes’.) If you cannot access the Internet, contact your
teacher to ask for the PhET CD. You previously used this interactive simulation in activity 2B.
Follow the instructions below before answering the questions. Sufficient information is given to
enable you to answer these questions.
1. When you have downloaded the simulation, click on the Mix isotopes tab on the top left.
Then click on Cl on the Periodic Table and Nature’s mix of isotopes. Make sure that Percent
Composition and Average Atomic Mass is showing. You should see something like the
diagram on the right.
2. Make a note of the percentage abundance of the 35Cl
isotope and the 37Cl isotope. The average atomic mass
is stated as 35.453 amu (atomic mass units). How is this
calculated? The following calculation uses a rounded off
75% and 25% for the abundance of each isotope.
75
25
( 100 x 35 ) + ( 100 x 37 )
=26.25+9.25
=35.5
So the relative atomic mass for chlorine is 35.5.
PhET
9A
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12.0
14.0
the mole
3. Now click on Li. Look at the percentage abundance of 6Li and
7
Li. Round these off to 7.6% and 92.4%. Make a note of the given
average atomic mass.
Question
a. Calculate the relative atomic mass for lithium. Show all your
working.
4. Click on F on the Periodic Table. Make a note of how many
isotopes it has and its average atomic mass.
Information from step 3.
5. Do the same for Na, Al and P. Is the average atomic mass exactly
the same as the mass number?
Question
PhET
a. Fluorine has only one isotope, 19F, yet does not have an
average atomic mass of exactly 19.000. What does this tell
you about the relative mass of protons compared to that of
neutrons?
Check your answers.
Information from step 4.
relative molecular mass (rmm)
The relative atomic mass scale can be used to find the relative masses of molecules. The relative
molecular mass is the sum of the ram values of all the atoms in a molecule.
Consider hydrogen gas and water shown below. What is the relative molecular mass of hydrogen
gas or water?
H H
52
H
O
H
Example 1:
Hydrogen gas is made up of molecules. Each
molecule contains two hydrogen atoms and is
represented by the formula H2.
Example 2:
Each water molecule is made up of two
hydrogen atoms and one oxygen atom,
represented by the formula H2O.
rmm for hydrogen
= 2 (ram for H)
= 2 (1.0)
= 2.0
rmm for water
= 2 (ram for H) + ram for O
= 2 (1.0) + 16.0
= 18.0
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Keep the number after
the decimal point.
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the mole
Example 3: What is the relative molecular mass of citric acid?
The formula for citric acid is given as H3C6H5O7. This means that each molecule has six carbon
atoms, seven oxygen atoms and eight hydrogen atoms (3 + 5 = 8).
rmm for citric acid = 6 (ram for C) + 7 (ram for O) + 8 (ram for H)
6 (12.0)+7 (16.0)+8 (1.0)
72.0+112.0+8.0
=192.0
There are no units, as
rmm values are relative.
relative formula mass (rfm)
You have seen that relative molecular masses can be calculated by adding the ram values of all
the atoms in the formula of a molecule. For this reason, relative molecular masses may also be
called relative formula masses.
Many compounds do not exist as molecules – these include most metal compounds, for example,
NaCl, KOH, CuSO4, Fe2O3 as well as sodium hydrogen carbonate.
For these compounds, you should not talk about a ‘relative molecular mass’, but you can work
out a relative formula mass in the same way as you did for H2O.
ram values
N
= 14.0
For example, given the ram values, calculate the relative formula mass for
O
= 16.0
Mg(NO3)2.
Mg = 24.3
The following working can be abbreviated, but initially it is a good idea to set it
out in full as shown.
Formula
Atoms in formula
Ram of atoms
Mg(NO3)2
1 Mg
24.3
1 × 24.3
=
24.3
2N
14.0
2 × 14.0
=
28.0
1O
16.0
6 × 16.0
=
96.0
Total
=
148.3
So the relative formula mass for Mg(NO3)2 = 148.3.
© te ah o o t e k ur a p o un a m u
Formula mass
There are no units, as
rmm values are relative.
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the mole
9B
check your understanding
Work out the relative formula masses for (1) to (6) below. Show your working.
1. oxygen, O2
2. zinc sulfide, ZnS
3. methane, CH4
4. ethanoic acid, CH3COOH
5. calcium carbonate, CaCO3
6. ammonium sulfate, (NH4)2SO4
ram values
H = 1.0
C = 12.0
N = 14 . 0
O = 16 . 0
S = 32 . 1
Ca = 40 . 1
Zn = 65 . 4
Check your answers.
the mole
Knowing the relative masses of elements is useful, but it still doesn’t
help the chemist know exactly how much of a chemical to weigh out
in grams needed for a reaction. An Italian chemist, Amedeo Avogadro,
worked out that 6.02 × 1023 atoms of an element are found in a mass
of the element equal to its relative atomic mass in grams.
Avogadro’s number (NA)
Or = 6.02 × 1023
= 60 200 000 000 000 000 000 000!
This number is also known as the mole.
This means that as neon has a relative atomic mass of 20.2, one mole of
neon gas (Ne) would weigh 20.2 grams and 20.2 g of neon would contain
6.02 × 1023 neon atoms. This is represented below.
If you had one mole of dollars
and divided it amongst all the
people on Earth (about
7 billion), each person would
be a multi trillionaire.
$6.02 x 1023
7 000 000 000 people
= $86 000 000 000 000 each
One mole of Ne
6.02 × 1023 atoms
of Ne
20.2 g of Ne
Meet the
chemistry mole.
Remember that a mole simply represents a number, just like ‘dozen’
represents the number 12, ‘pair’ represents two and ‘score’ represents 20.
The mole is a unit of measure, just as the kilogram and the metre
are units of measure. The mole is the chemist’s unit for measuring
amount.
It enables you to count atoms and molecules, because you know that
one mole of a substance contains the Avogadro number of particles.
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the mole
The term ‘particles’ can mean atoms, ions, molecules or formula units, depending on what is
being counted.
1. particles: atoms
For the following elements, ‘particles’ are single atoms:
•• all metals. For example: lithium, magnesium, iron, copper
•• the noble gases, Group 18 of the Periodic Table (e.g. He, Ne, Ar)
•• carbon, sulfur and phosphorus.
For example:
–– 1 mole of copper contains NA copper atoms or 6.02 × 1023 atoms.
–– ½ mole of copper contains ½NA copper atoms or 3.01 × 1023
atoms.
PhET
–– 3 moles of copper contain 3NA copper atoms or 1.81 × 1024 atoms.
copper ram = 63.5
When we say ‘one mole of carbon’, we know it means one mole of
1 mole copper
= 63.5 g
carbon atoms. However, what do we mean when we state ‘one mole of
hydrogen’? Do we mean one mole of hydrogen atoms (H) or one mole of hydrogen molecules
(H2)? This problem applies to any element that exists as diatomic molecules.
2. particles: diatomic molecules
This includes hydrogen (H2), nitrogen (N2), oxygen (O2), fluorine (F2), chlorine
(Cl2), bromine (Br2) and iodine (I2).
Note that this includes all the elements that are gases, apart from Group 18.
One mole of an element that consists of diatomic molecules will contain NA
number of particles. Are these particles atoms or molecules?
To avoid confusion the type of particle should always be specified, and so you
must say either:
1. a mole of oxygen atoms, or
2. a mole of oxygen molecules, or
3. give the formula as well as the name, so: a mole of oxygen, O, or a mole of
oxygen, O2.
oxygen ram = 16.0
oxygen rmm = 32.0
1 mole oxygen
molecules= 32.0 g
Given the ram for O = 16.0, the amounts (1) and (2) would have masses of 16.0 g and 32.0 g
respectively.
If the diatomic molecule is a gas, then the name followed by ‘gas’ implies the molecule. For
example, hydrogen gas refers to the hydrogen molecule H2.
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the mole
3. particles: compounds
Compounds can be either molecular or ionic. A molecular
compound is usually a group of non-metal atoms from two or more
different elements that are covalently bonded together. Examples of
molecular compounds are water and carbon dioxide. Molecular
compounds are usually gases or liquids, although large molecules
like sugar are solids.
O C O
The rmm of CO2 is 12 + 2(16) = 44.0.
So 1 mole of CO2 will weigh 44.0 g.
A mole of these compounds is the amount containing the Avogadro
number of molecules.
Ionic compounds include metal compounds such as metal oxides
and hydroxides (for example, MgO, Ca(OH)2, Fe2O3) and also salts (for
example, CuSO4, NaCl, PbCO3).
This lattice represents Fe2O3
with each oxygen (red) atom
surrounded by iron (dark grey)
atoms. The ratio is represented by
a formula unit Fe2O3.
These compounds are not made up of molecules. One mole of an ionic compound contains the
Avogadro number of its ‘formula unit’.
For example, 1 mole of iron(III) oxide, Fe2O3, contains NA Fe2O3 units made up from 2NA iron ions
and 3NA oxide ions.
9C
check your understanding
1. Use the ram values to give the mass of the following amounts. Show your
working.
a. 2 moles of carbon
b. half a mole of phosphorus
c. 1 mol chlorine atoms
d. 1 mol chlorine molecules, Cl2
e. 2 mol iodine molecules, I2
‘mol’ is the symbol
for the unit ‘mole’
ram values
C = 12.0
O = 16 . 0
P = 31 . 0
S = 32 . 1
Cl = 35 . 5
Cu = 63 . 5
I
= 127 . 0
f. 3 mol oxygen gas, O2
2. How many atoms of each of the following elements are there in 1 mole of the compound
Na2CO3?
a. sodium
b. carbon
c. oxygen
Check your answers.
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the mole
key points
•• The actual masses of atoms are inconveniently small to use, so chemists use a scale of
relative atomic mass (ram) values. An atom of carbon-12 is assigned a ram value of 12.00
and other atoms are compared with carbon-12 atoms to assign them values.
•• Ram values have no units; they are just numbers. You don’t have to learn ram values.
•• Compounds have a relative formula mass (rfm). The relative formula mass for a
compound is the sum of the ram values of all the atoms shown in the formula of the
compound. For compounds which are made up of molecules (for example, water)
relative formula masses can be called relative molecular masses.
•• A mole is the amount of a substance containing the Avogadro number (NA) of particles,
where:
–– ‘amount’ is a quantity chemists use for measuring substances
–– ‘particles’ can be atoms, molecules or formula units.
•• A mole of these elements contains NA atoms:
–– all metals
–– Group 18 (noble gases)
–– carbon, phosphorus, sulfur.
•• When using the term ‘mole’, care must be taken to specify whether you are referring
to atoms of the element or the molecule for elements which can exist as diatomic
molecules (e.g. H2, N2, O2, F2, Cl2, Br2, I2).
•• A mole of a molecular compound contains NA molecules. These include all compounds
which are liquids and gases, and some solids.
•• A mole of a metal compound contains NA formula units of the compound.
•• The symbol ‘mol’ is used for the unit ‘mole’.
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10
mole calculations
learning intention
In this lesson you will learn to:
•• use the mole and molar mass in calculations.
molar mass
You learnt in lesson 9 that you could find out the mass of one mole of any element by adding the
unit gram after the relative atomic mass for that element.
This value is known as the molar mass.
The photo above shows 1 mole of sugar (sucrose) which has been measured by weighing out the mass in grams numerically equal to
the relative molecular mass. Sucrose has the formula C12H22O11 and the relative molecular mass = 342. Its molar mass = 343 g.
For any substance the mass of 1 mole is called the molar mass. It has the symbol M and the unit
grams per mole (g mol–1).
So for sucrose, M = 342 g mol–1, or we can write M (sucrose) = 342 g mol–1.
Note these differences between the terms mole and molar mass:
1.
2.
Mole
Mole is a unit (for measuring the
amount of a substance).
A mole is a convenient way of counting
packages containing 6.02 × 1023
particles of a substance.
Molar mass
Molar mass is not a unit; rather, it is a
quantity which has units of grams per mole.
Molar mass is a measure of the mass of one
of these packages containing 6.02 × 1023
particles.
The units and symbols for molar mass and the amount of substance are listed in the table below.
Quantity
amount of a
substance
molar mass
58
CH3001
Symbol
n
Unit
mole
Symbol
mol
M
grams per mol
g mol–1
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mole calculations
using molar mass
You can use the molar mass to work out the mass of an amount of substance.
a worked example
What is the mass of 3.50 mole of aluminium, given the ram for Al = 27.0?
For aluminium, M (molar mass) = 27.0 g mol–1
so the mass of 3.50 mol = 3.5 mol × 27.0 g mol–1
= 94.5 g
This illustrates the following useful formula.
m
= n×M
mass
(g)
10A
amount of
substance
(mol)
molar mass
(g mol–1)
check your understanding
1. Rearrange the formula above to give two other equations, one to find n = , and the other to find M =.
2. Given the relative formula mass for Na2CO3 = 106, write down the molar mass for this
compound, giving its units.
3. Use your result from (2) to work out the masses of the following amounts of Na2CO3. Show
your working and give units.
a. 3.00 mol
b. 6.50 mol
c. 0.250 mol
d. 2.00 × 10–3 mol
Check your answers.
important formulae
You must remember the three formulae which express the relation between m, n and M. Perhaps
the easiest way to do this is to remember the unit for molar mass, M, is g mol–1.
This unit tells us that to find M we must divide
M (g mol–1)
so M
© te ah o o t e k ur a p o un a m u
=
=
grams
.
moles
grams
moles
m
(mass)
n
(amount of substance)
Be careful when
writing the letter ‘m’.
A lower case m means
mass. An upper case
M means molar mass.
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59
mole calculations
Now this formula can be rearranged to give:
m = n × M
n =
and
m
M
Alternatively, use the triangle method illustrated below.
m
m
× M
n
n
m=n×M
n=
× M
m
M
m
n
M=
× M
m
n
The symbol m must go at the top. Note how the three formulae above can be derived from the
arrangement of letters in the triangle.
worked examples
The following two examples and one given earlier this lesson will help you with the next activity.
1. Calculate the molar mass, M, of a compound, given that 0.150 mole of the compound has a
mass of 8.40 g.
m
n
8.40 g
=
0.150 mol
M (molar mass) =
= 56.0 g mol–1
2. Given the relative formula mass for oxalic acid = 126.0,
calculate the number of moles in 82.1 g of oxalic acid.
n(moles)
=
=
60
m
M
82.1 g
126.0
g mol–1
=
0.6515873 mol (by calculator)
=
0.652 mol (to three significant figures)
CH3001
ram for oxalic acid = 126.0
So M for oxalic acid
= 126.0 g mol–1
Would you expect the
answer to be bigger or
smaller than 1 mole?
© te ah o o te k u ra p ou n a mu
mole calculations
important steps
1. Show full working
Always show all your working for calculations. This is a good habit to develop because it
helps you work logically and make fewer mistakes. It also shows the teacher/examiner your
understanding of the topic. Always put in units with your working, and of course with your
answer.
2. Check your answer
It is always a good idea to check that an answer is sensible. For example, in the first worked
example you can do this by noting in the question that 0.150 mol = 8.40 g. So you would expect 1
mol of the compound to have a mass quite a lot greater than 8.40 g. If your answer was smaller
than 8.40 (for example, if in your working you had divided instead of multiplying) then this should
warn you that something is wrong.
In the second worked example you should check whether your answer is less than 1 mole. This is
because the given mass of 82.1 g is less than the mass of 1 mol (M for oxalic acid is 126 g mol–1).
Putting in units with the working also gives you a check. The units for the answer above are given
by dividing the units for m by the units for M.
m
M
=
g
g mol–1
=
g
g mol–1
=
1
mol–1
= mol
You can cancel the two g units, leaving the unit = mol. This gives you the correct units for n.
10B
online activity
Go to the Topic webpage to see how mass and moles are related. This is an interactive activity.
You will also complete an activity on the mole. Or you can complete the calculations below.
Show all your working for calculations.
1. Find the number of mol in each of the following amounts of substance.
a. 85.0 g of Mg
b. 8.50 g of SO2
c. 435 g of CaCO3
2. Calculate the mass, in grams, of each of these chemical amounts.
a. 0.255 mol of chlorine molecules, Cl2
b. 7.50 mol of NaOH
c. 4.50 × 10–3 mol of Na2CO3
M (g mol–1)
H = 1.0
C = 12.0
O = 16 . 0
Na = 23. 0
Mg = 24. 3
S = 32. 1
Cl = 35. 5
Ca = 40. 1
3. a.3.50 mol of a metallic element has a mass of 140 g. Calculate M (the
molar mass) and the ram for the element.
b.2.15 × 10–2 mol of a compound has a mass of 1.29 g. Calculate M and the rfm for the
compound.
Check your answers.
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mole calculations
concentration
It is also important to be able to work out the concentration of a substance dissolved in a
solution. It gives the ratio of amount of solute in a particular volume of solution. The diagrams
below all show the same concentration.
20 blobs in 2 litres
10 blobs in 1 litre
5 blobs in ½ a litre
The standard unit for measuring concentration is moles per litre (mol L–1). Concentration can be
calculated using the following formula:
c (mol L–1)
=
n (mol)
where
V (litre)
c = concentration
n = number of moles
V = volume in litres
Use capital V for volume.
Below is a sample calculation.
Question
A solution contains 0.125 mol of sodium hydroxide in 20.0 mL of solution. Find its concentration
in mol L–1.
Answer
First write down c, n and V and what you know about them from the question:
c=?
n = 0.125 mol
V = 20.0 mL
You don’t know the value of c
as you are asked to calculate it.
Next change the volume from mL to litres. The simplest way is to multiply the value in mL by 10–3.
V = 20.0 mL = 20.0 × 10–3 L
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mole calculations
Then substitute into the equation c = n/V
c (NaOH)
=
=
=
n
V
0.125 mol
0.0200 L
6.25 mol L–1
The concentration of NaOH = 6.25 mol L–1
10C
Check the significant figures.
Get into the habit of stating the answer
in full. If you only write a lot of numbers
it can be difficult following your working,
particularly if there is more than one step
involved.
check your understanding
Show all your working for calculations.
1. Calculate the concentration of a solution of sulfuric acid with 0.050 mol H2SO4 in 250 mL of
solution.
2. Calculate the amount of sodium carbonate (Na2CO3) in 1.00 × 102 mL of a 2.00 mol L–1
solution.
3. Calculate the concentration of 50.0 mL of HCl solution which contains 0.730 g HCl.
M(HCl) = 36.5 g mol–1.
Hint: You will also need to use the formula n = m/M in question 3 in order to first calculate the
amount in mol.
Check your answers.
significant figures
Throughout this course you will do many calculations. To achieve the necessary accuracy in
calculations you will need to use significant figures correctly.
Significant figures – These are figures all of whose digits (whole numbers) are known with
certainty. For example:
•• 261 has three significant figures
•• 79 has two significant figures
•• 1002 has four significant figures.
In general it is fairly easy to determine how many significant figures are present in a number by
using the following rules:
1. Any digit that is not zero is significant. Thus 845 cm has three significant figures, 1.234 kg has
four significant figures, and so on.
2. Zeros between non-zero digits are significant. Thus 606 m contains three significant figures,
40,501 J contains five significant figures, and so on.
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mole calculations
3. Zeros to the left of the non-zero digit are not significant. These zeros are used to indicate the
placement of the decimal point. Thus 0.08°C contains one significant figure, 0.000 0349 g
contains three significant figures, and so on.
4. If a number is greater than one, then all the zeros to the right of the decimal point count
as significant figures. Thus 2.0 mg has two significant figures, 40.062 mL has five significant
figures, 3.040 dm has four significant figures, and so on. If a number is less than one, then
only the zeros that are at the end of the number and the zeros that are between non-zero
digits are significant. Thus 0.090 kg has two significant figures, 0.3005 J has four significant
figures, 0.00420 min. has three significant figures, and so on.
5. For numbers that do not contain decimal points, the trailing zeros may or may not be
significant. Thus 400 cm may have one significant figure (the digit 4), two significant figures
(40), or three significant figures (400). You cannot know which of these is the case without
having more information.
By using exponential notation, however, you avoid this ambiguity. In this particular case, you
can express the number 400 as 4 × 102 for one significant figure, 4.0 × 102 for two significant
figures, and 4.00 × 102 for three significant figures.
10D
check your understanding
Give the number of significant figures for the following numbers:
1.3.06
2.0.0004
3.20.07
4.4.000
5.0.0708
6. 3.210 × 10–6
7.501.00
8.0.010
Check your answers.
significant figures in calculations
How old is
the dinosaur?
Wow! How can you
know so exactly?
Three million and
2 years old!
Easy! When I started working
here 2 years ago I was told it
was 3 million years old.
Knowing which numbers are significant can stop you from making your answers
more accurate than they really are!
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mole calculations
When measurements are added or subtracted, the answer can contain no more decimal places
than the least accurate measurement.
For example, you can’t say 3 000 000 years + 2 years = 3 000 002 years, as you are going from 1
significant figure to 7!
You can say 1.0 + 32.1 = 33.1 (you do not round off to 33, as both measurements are only uncertain
in the first decimal place.)
You can’t say 43.56 mL – 35.13 mL = 8.430 mL just because both the measurements have four
significant figures. The second decimal place is uncertain and so the answer cannot be given
more precisely than 8.43 mL.
When multiplying or dividing, you must round off your final answer to the same number of
significant figures as given in the measurement with the fewest significant figures used in the
calculation.
For example:
82.1
←
this value has three significant figures
126.0
←
this value has four significant figures
So you round off the answer to three significant figures.
key points
•• For any substance, the mass of one mol is called the molar mass, which has the symbol M.
•• Molar mass (M) has units of g mol–1, while amount (n) has units of mol.
•• m = n × M where
–– m = mass (g)
–– n = amount (mol)
–– M = molar mass (g mol–1)
•• You should be able to rearrange the above equation to find either n
or M. The triangle alongside provides one way of doing this.
•• The concentration of a solution can be calculated using the formula
c = n/V.
•• The unit for concentration is mol L–1.
m
n
× M
n
c
× V
•• To convert from mL to L, multiply by 10–3.
•• In any calculation, the answer must reflect the same degree of accuracy as the least
accurate data. This means that you need to consider the number of significant figures
you write in your answer.
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11
teacher-marked assignment
next steps
Before you finish off this topic you should have agreed your next steps with your teacher. If you do
not have your next set of study materials, contact your teacher immediately. If you are not sure
what to do next, ask your teacher for advice.
In this lesson, have a quick look back at all the lessons you have done in this topic. Think about
what you have learned. When you are ready, try the teacher-marked assignment CH3001A. If you
did not receive this with your booklet, contact your teacher.
When you have finished, complete the self-assessment section at the end of this booklet. Send
your answers to the activities in the booklet and the teacher-marked assignment to your teacher.
Make sure that you have written your name and ID number on the cover sheet of your booklet and
the teacher-market assignment. You can also use a label if you have one at hand.
You may send your work by post or electronically.
by post:
Put your answers with the cover sheet of the booklet and teacher-market assignment in the
plastic envelope provided. Make sure that the address card shows the address for Te Aho o Te
Kura Pounamu (The Correspondence School). Seal the envelope with tape before you post it. Do
not include booklets from other subjects with this Science work.
by email:
Scan your answers and the cover sheet and email to your teacher. The standard format for Te
Kura teacher email addresses is: [email protected]
If you aren’t sure who your teacher is, call 0800 65 99 88.
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12
glossary
acid
A substance that neutralises a base. A substance that releases H+
ions in solution.
anion
A negatively charged ion formed from an atom that has gained one
or more electrons.
atom
The smallest unit of an element that has the chemical
characteristics of that element.
atomic number
The number of protons in an atom.
Avogadro’s number
6.02 × 1023
cation
A positively charged ion formed from an atom that has lost one or
more electrons.
chemical change
A process where a new substance is formed.
coefficient
The number in front of a formula in a balanced equation.
compound
A substance made up of two or more elements chemically joined
together.
concentration
The amount of substance per volume of solution.
covalent bond
The attraction between the nuclei of the bonding atoms and the
shared pair of electrons. Covalent bonds form between non-metal
elements.
electron
A negatively charged particle found in the space around the
nucleus of an atom.
electron
configuration
The arrangement of electrons in the energy levels around the
nucleus.
element
A substance which cannot be broken down further by chemical
means. A substance that consists of atoms, all having the same
number of protons.
formula unit
The smallest repeating unit of a substance or the empirical formula
of an ionic or covalent network solid compound.
ion
A charged particle formed from an atom that has lost or gained an
electron.
ionic bond
The electrostatic attraction between the positive cation and the
negative anion.
isotope
An atom of the same element that has a different number of
neutrons.
mass number
The total number of protons and neutrons in the nucleus of the
atom.
metal
An element which is typically shiny, malleable, ductile with good
electrical and thermal conductivity.
metallic bond
The electrostatic attraction between the positive metal atom core
and the delocalised electrons. As the electrons are not fixed to a
particular atom, metallic bonds are non-directional.
© te ah o o t e k ur a p o un a m u
CH3001
67
glossary
68
metalloid
An element which exhibits some properties of metals and nonmetals.
molar mass
The mass of one mole of substance in grams.
mole
The amount that contains 6.02 × 1023 particles of that substance.
molecule
A discrete group of atoms bonded together, which retains the
chemical and physical properties of that substance.
neutron
An uncharged particle, similar in mass to a proton, found in the
nucleus of an atom.
non-metal
An element that lacks the physical and chemical properties of
metals. They are able to form anions and have acidic oxides.
nucleus
The central core of the atom containing the protons and neutrons.
polyatomic ion
A charged species composed of two or more covalently bonded
atoms.
product
A substance that is produced or formed by a chemical reaction.
proton
A positively charged particle, similar in mass to a neutron, found in
the nucleus of an atom.
reactant
A substance that takes part in and undergoes change during a
reaction.
relative atomic mass
The ratio of the average mass of an atom when compared to
carbon-12.
significant figures
The numbers that express the precision of a measurement instead
of its magnitude.
stoichiometry
The ratio of reactants and products for a reaction.
valence electrons
The electrons found in the outermost (or highest) energy level.
valency
A property of atoms or groups, equal to the number of atoms of
hydrogen that the atom or group could combine or displace in
forming compounds.
CH3001
© te ah o o te k u ra p ou n a mu
13
answer guide
1.
1A
elements and the periodic table
1. Sodium, magnesium and aluminium are all metals, as they are shiny and malleable (based on
the photos). You may have included silicon, as it is also shiny.
2. Phosphorus, sulfur, chlorine and neon are non-metals, as they do not exhibit the properties
of metals (e.g. they are not shiny).
3. Silicon is shiny like a metal, but is a poor conductor of electricity like a non-metal. (This is
one of the reasons why it is used in computer chips.)
1C
1. a.Na
e.Fe
i.Pb
b.Si
f.Mn
j.P
c.S
g.Mg
k.F
d.I
h.Li
l.Hg
d.copper
g.nitrogen
b.chlorine
e.calcium
h.sodium
c.carbon
f.cobalt
i.neon
2. a.chromium
1D
1. Si is the symbol for silicon. Ag = silver
2. Second letter should be lower case. Cl = chlorine (If you write a lower case ‘l’ with a flick,
make sure that it can’t be mistaken as an uppercase L.)
3. First letter should be upper case. Mg = magnesium
4. There should be no second letter. F = fluorine
5. N is the symbol for nitrogen. Na = sodium
6. First letter should be upper case. Ne = neon
7. Co is the symbol for cobalt. Cu = copper
8. Second letter should be lower case. Al = aluminium
2.
2A
structure of the atom
1. electrons and protons
2. neutrons and protons
3. protons
4. a. Number of protons and neutrons (or mass number)
b.Charge
c. Number of protons (or atomic number)
d. Symbol for the element
© te ah o o t e k ur a p o un a m u
CH3001
69
answer guide
2B
1. Beryllium (Be) and fluorine (F) have only one isotope.
2. 73Li and 63Li are isotopes of lithium.
3.
O : 99.757%
O : 0.038%
16
17
18
O : 0.205%
4. Hydrogen ( H) and carbon ( C).
3
14
5. Atoms tend to be unstable if too many neutrons are present.
(Extra information for interest: Smaller elements usually have close to a 1 : 1 ratio between
protons and neutrons. Larger elements tend to need more neutrons to hold a large number
of protons together. For example, gold-197 is stable, with 79 protons and 118 neutrons. There
are 36 unstable or radioactive isotopes of gold. The most stable elements usually have an odd
number of protons and an even number of neutrons.)
3.
3A
electrons and energy levels
1. ‘Billiard ball’ model, ‘plum pudding’ model and the classical solar system model.
2. Bohr model, de Broglie model and Schrödinger model.
3. The Schrödinger model most closely matched the lines in the atomic spectrum for hydrogen.
(The Bohr model gave too many lines in the infrared region and did not match the intensities
of the lines closely. The de Broglie model better matched the intensities but also had too
many lines in the infrared region.)
Snapshot 4: Experiment
92
UV
380
500
600
700
Snapshot 1: Bohr
780
7500
IR
Snapshot 2: de Broglie
92
70
UV
380
500
CH3001
600
92
UV
380
500
600
700
780
7500
IR
780
7500
IR
Snapshot 3: Schrödinger
700
780
7500
IR
92
UV
380
500
600
700
© te ah o o te k u ra p ou n a mu
answer guide
3C
3D
Element
Atomic number
Electron arrangement
H
1
1
He
2
2
Li
3
2, 1
Be
4
2, 2
B
5
2, 3
C
6
2, 4
N
7
2, 5
O
8
2, 6
F
9
2, 7
Ne
10
2, 8
Na
11
2, 8, 1
1. 1 valence electron
7. 3 valence electrons
2. 1 valence electron
8. 5 valence electrons
3. 1 valence electron
9. 5 valence electrons
4. 2 valence electrons
10. 6 valence electrons
5. 2 valence electrons
11. 6 valence electrons
6. 3 valence electrons
12. 7 valence electrons
4.bonding
4A
1. a.The attraction between the negative electrons of one atom and the positive protons of
the other atom reduces the potential energy of each atom. When the atoms get too close,
the electrons of the atoms repel each other strongly, resulting in an increase in potential
energy.
b.No. The forces of attraction between the two helium atoms are very weak.
2. a.There is a strong force of attraction between the hydrogen atoms. There are fewer
electrons to repel each other. It is also thought that the large drop in potential energy is
due to each hydrogen atom gaining stability from a full valence energy level.
4. a.The electrons get too close and repel each other. This repulsive force results in an
increase in potential energy.
5. a.The atoms are far away from each other so do not attract each other. Due to random
motion, the atoms may get closer together until they are close enough to feel an
electrostatic force of attraction.
© te ah o o t e k ur a p o un a m u
CH3001
71
answer guide
7. a.Hydrogen has fewer electrons to repel each other. The hydrogen atom has a much smaller
atomic diameter than a fluorine atom so there will be a stronger attraction between
protons and electrons.
b.Hydrogen forms a stronger bond than fluorine as more energy has been lost (or has been
stabilised to a greater extent).
8. a.From strongest to weakest: Oxygen–oxygen > hydrogen–hydrogen > fluorine–fluorine.
b. The more energy released to form the bond, the stronger the bond.
4B
1. Covalent bonding (as between two non-metal atoms)
2. Covalent bonding (as between two non-metal atoms)
3. Ionic bonding (as between a metal and a non-metal atom)
4. Ionic bonding (as between a metal and a non-metal atom)
5. Metallic bonding (as between two metal atoms)
6. Ionic bonding (as between a metal and a non-metal atom)
7. Ionic bonding (as between a metal and a non-metal atom)
8. Covalent bonding (as between two non-metal atoms)
5.
5A
5B
5C
valency and formulae
1.NaCl
6.CH4
2.CaO
7.H2S
3.MgBr2
8.Al2O3
4.PH3
9.SiO2
5.FeCl3
10.HCl
1.NaOH
6.Ca(HCO3)2
2.Mg(NO3)2
7.(NH4)2CO3
3.Fe(OH)2
8.H2SO4
4.Na2SO4
9.Al2(SO4)3
5.ZnCO3
10.Ca3(PO4)3
1. iron : chlorine = 1 : 3
2. aluminium : oxygen = 2 : 3
3. sodium : carbon : oxygen = 2 : 1 : 3
4. copper : oxygen : hydrogen = 1 : 2 : 2
5. potassium : oxygen = 2 : 1
6. calcium : phosphorus : oxygen = 3 : 2 : 8
72
CH3001
© te ah o o te k u ra p ou n a mu
answer guide
6.
6A
6B
naming compounds
1. potassium chloride
6. ammonium sulfate
2. magnesium iodide
7. sodium carbonate
3. iron(II) oxide
8. silver phosphate
4. iron(III) chloride
9. magnesium hydrogen carbonate
5. lead sulfide
10. calcium carbide
1. nitrogen dioxide
2. sulfur dioxide
3. sulfur trioxide
4. carbon disulfide
If the first element in a
formula only has one atom,
‘mon’ is usually ignored.
5. silicon tetrachloride
6. oxygen difluoride
7. dichlorine monoxide (or chlorine monoxide)
Chlorine monoxide is
acceptable as there is no
ClO compound.
8. chlorine dioxide
6C
1.
Cl–
chloride
2.
ClO–
hypochlorite
3.
ClO2–
chlorite
4.
ClO3–
chlorate
5.
ClO4–
perchlorate
6.
S2–
sulfide
7.
SO32–
sulfite
8.
SO42–
sulfate
9.
S2O32–
thiosulfate
© te ah o o t e k ur a p o un a m u
CH3001
73
answer guide
7.
chemical reactions
1. a.When colourless aqueous solutions of lead nitrate and sodium iodide were mixed
together, a bright yellow precipitate of lead iodide was formed.
7A
lead iodide(solid) + sodium
b.lead nitrate(aqueous) + sodium iodide(aqueous)
nitrate(aqueous)
2. a.When colourless aqueous solutions of lead nitrate and sodium chloride were mixed
together, a white precipitate of lead chloride was formed.
lead chloride(solid) +
b.lead nitrate(aqueous) + sodium chloride(aqueous)
sodium nitrate(aqueous)
3. a.When colourless aqueous solutions of lead nitrate and sodium bromide were mixed
together, a white precipitate of lead bromide was formed.
lead bromide(solid) +
b.lead nitrate(aqueous) + sodium bromide(aqueous)
sodium nitrate(aqueous)
1. The solid white marble chip of calcium carbonate became smaller as it reacted with the acid.
Bubbles of gas were seen.
7B
2. calcium carbonate(solid) + hydrochloric acid(aqueous)
+ water(liquid) + carbon dioxide(gas)
3. The solid calcium carbonate is a base and reacts with the
hydrogen ions from the hydrochloric acid. A hydrogen ion
combines with one of the oxygen atoms in the carbonate ion
to form water. The remaining carbon and two oxygen atoms
form carbon dioxide gas. The chloride and calcium ions
remain in solution.
4. CaCO3(s) + 2HCl(aq)
7C
calcium chloride(aqueous)
You should be able
to answer question 3
using NCEA Level 1 and 2
chemistry. Visit the Topic
webpage for more help if
you did not know how to
start explaining this.
CaCl2(aq) + H2O(l) + CO2(g)
1. The dark grey iron nail became coated with a
pink solid. The blue colour of the solution faded
to colourless.
2. iron(solid) + copper sulfate(aqueous)
Don’t worry if you did not have the
‘2’ in front of the HCl. You will learn
about balancing equations in the
next lesson.
copper(solid) + iron(II) sulfate(aqueous)
3. The iron displaces the copper ions in solution. It does this by giving its electrons to the
copper. This means that the iron atoms form Fe2+ ions and the Cu2+ ions form Cu atoms,
resulting in the pink deposit. The solution was blue due to the Cu2+ ions. The colour fades as
the concentration of Cu2+ decreases and colourless Fe2+ ions are formed.
7D
4. Fe(s) + CuSO4(aq)
Cu(s) + FeSO4(aq)
1. Ag+(aq) + Br–(aq)
AgBr(s)
2. Cu (aq) + Fe(s)
2+
Cu(s) + Fe2+(aq)
3. CO32–(aq) + Mg2+(aq)
MgCO3(s)
4. 2H+(aq) + CaCO3(s)
Ca2+(aq) + H2O(l) + CO2(g)
74
CH3001
HCl dissociates into H+
and Cl– ions in solution. The
Cl– ions do not react.
© te ah o o te k u ra p ou n a mu
answer guide
8.
balancing symbol equations
8A
2. a. There are 2 nitrogen atoms and 2 hydrogen atoms on the left side of the equation.
There is 1 nitrogen atom and 3 hydrogen atoms on the right side of the equation.
b.One nitrogen atom has disappeared in the products and one hydrogen atom has been
created.
2NH3
3.a.2N2 + 3H2
4.a.2H2O
O2 + 2H2
CO2 + 2H2O
5.a.CH4 + 2O2
3H2 + CO
7. a.CH4 + H2O
CS2 + 2H2S
CH4 + 3S
S + 2H2O
SO2 + 2H2
4NH3 + 3O2
2N2 + 6H2O
4CO2 + 2H2O
2C2H2 + 5O2
8B
1. Full:ZnCl2(aq) + Na2CO3(aq)
Ionic:Zn (aq) + CO (aq)
2. Full: Zn(s) + 2HCl(aq)
2–
3
2+
ZnCO3(s) + 2NaCl(aq)
ZnCO3(s)
ZnCl2(aq) + H2(g)
Ionic: Zn(s) + 2H+(aq)
Zn2+(aq) + H2(g)
Ca(NO3)2(aq) + 2H2O(l)
3. Full: Ca(OH)2(s) + 2HNO3(aq)
Ionic:Ca(OH)2(s) + 2H+(aq)
4. Full:Al2O3(s) + 6HCl(aq)
Ionic: Al2O3(s) + 6H+(aq)
9.
9A
Ca2+(aq) + 2H2O(l)
2AlCl3(aq) + 3H2O(l)
Reactants: Zinc metal
contains atoms, so must be
written as Zn (not Zn2+).
Products: Hydrogen gas
contains molecules so
must be written as H2 and
not 2H+.
2Al3+(aq) + 3H2O(l)
the mole
3.a.
7.6
92.4
( 100 x 6 ) + ( 100 x 7 )
= 0.456 + 6.468
= 6.924
5.a.Protons and neutrons do not have exactly the same mass. A neutron is slightly heavier
than a proton.
9B
1. rfm (O2) = 2(16.0) = 32.0
rfm = relative formula mass
2. rfm (ZnS) = 65.4 + 32.1 = 97.5
3. rfm (CH4) = 12.0 + 4(1.0) = 16.0
4. rfm (C2H4O2)
= 2(12.0) + 4(1.0) + 2(16.0)
= 24.0 + 4.0 + 32.0
= 60.0
© te ah o o t e k ur a p o un a m u
Ethanoic acid, CH3COOH, is written
as C2H4O2 for ease of counting the
atoms of each element.
CH3001
75
answer guide
5. rfm (CaCO3)
= 40.1 + 12.0 + 3(16.0)
= 40.1 + 12.0 + 48.0
= 100.1
6. rfm (NH4)2SO4 = 2(14.0 + 4.0) + 32.1 + 4(16.0)
= 36.0 + 32.1 + 64.0
= 132.1
*rfm = relative formula mass
9C
1. a. ram (C) = 12.0 so 2 mol of C atoms = 2(12.0) = 24.0 g
b. ram (P) = 31.0 so ½ mol of P atoms = ½(31.0) = 15.5 g
c. ram (Cl) = 35.5 so 1 mol of Cl atoms = 35.5 g
d. rfm (Cl2) = 2 × 35.5 = 71.0
So mass of 1 mole of Cl2 molecules = 71.0 g
e. rfm (I2) = 2 × 127 = 254
So mass of 2 moles of I2 molecules = 2 × 254 g
= 508 g
f. rfm (O2) = 2 × 16.0 = 32.0
So mass of 3 moles of I2 molecules = 3 × 32.0g
= 96.0 g
2. a. 2 mol of Na atoms
b. 1 mol of C atoms
c. 3 mol of O atoms
10. mole calculations
10A
1.
n =
m
M
m
n
M =
2. 106 g mol–1
10B
3. a. 3.00 × 106
= 318 g
= 689 g
= 26.5 g
= 0.212 g
b. 6.50 × 106
c. 0.250 × 106
d. 2.00 × 10–3 × 106
1. a. m = 85.0 g
so
b. m = 8.50 g
so
n
c. m = 435 g
so
76
n
n
CH3001
M(Mg) = of 24.3 g mol–1
=
m
M
=
85.0
= 3.50 g
24.3
M(SO2) = 64.1 g mol–1
m
8.50
=
=
= 0.133 g
M
64.1
0.132 is wrong as it has been
rounded off incorrectly (or you
did not keep all the figures in
your calculator).
M(CaCO3) = 100.1 g mol–1
m
435
=
=
= 4.35 g
M
100.1
© te ah o o te k u ra p ou n a mu
answer guide
2. a. n = 0.255 mol M(Cl2) = 71.0 g mol–1
so
m = n × M = 0.255 × 71.0 = 18.1 g
b. n = 7.50 mol M(NaOH) = 40.0 g mol–1
so m = n × M = 7.50 × 40.0 = 3.00 × 102 g
c. n = 4.50 × 10–3 mol M(Na2CO3) = 106.0 g mol–1
so m = n × M = 4.50 × 10–3 × 106.0 = 0.477 g
3. a. n = 3.50 mol
so
=
m = 140 g
140
=
= 40.0 g mol–1
3.50
and rfm = 40.0
b. n = 2.15 × 10–2 mol m = 1.29 g
m
1.29
so
M =
=
= 60.0 g mol–1
n
2.15 × 10–2
10C
M
m
n
Don’t forget to round off to
three significant figures.
and rfm = 60.0
1. c = ? mol L–1
n = 0.050 mol
V = 250 mL = 250 × 10–3 L
So
c
2. n = ?
=
n
V
=
0.050 mol
250 × 10–3 L
Answer is given to two
significant figures as each value
only had two significant figures.
= 0.20 mol L–1
V = 1.00 × 102 mL = 1.00 × 102 × 10–3 L
c = 2.00 mol L–1
n = c × V = 2.00 × 1.00 × 102 × 10–3 = 0.200 mol Na2CO3
3. c = ? mol L–1
You need to calculate n
n = ? mol
before you can calculate c.
m = 0.730 g
M = 36.5 g mol–1
V = 50.0 mL = 50.0 × 10–3 L
m
0.730 g
So
n =
=
= 0.0200 mol
M
36.5 g mol–1
So
10D
c
=
n
V
=
0.0200 mol
50.0 × 10–3 L
Answer is given to three
significant figures.
= 0.400 mol L–1
1.3
3.4
5.3
7.5
2.1
4.4
6.4
8.2
© te ah o o t e k ur a p o un a m u
CH3001
77
78
CH3001
39
Y
88.9
38
Rb Sr
87.6
37
85.5
261
262
226
262
263
106
223
105
104
Rf Db Sg Bh Hs Mt
103
88
Ir
264
107
265
108
268
109
192
87
190
Fr Ra Lr
186
184
181
179
W Re Os
77
103
175
76
101
45
58.9
27
137
75
98.9
44
55.8
26
133
74
95.9
43
54.9
25
Cs Ba Lu Hf Ta
73
92.9
42
52.0
24
9
11
12
195
197
79
108
47
63.6
29
201
80
112
48
65.4
30
204
81
115
49
69.7
31
Pt Au Hg Tl
78
106
46
58.7
28
10
15
207
Pb
82
119
209
Bi
83
122
51
74.9
33
31.0
I
210
210
85
127
53
79.9
35
35.5
17
222
86
131
Xe
54
83.8
36
Kr
40.0
18
20.2
10
Ne
4.0
He
Cl Ar
19.0
9
F
17
Po At Rn
84
128
52
79.0
34
32.1
16
S
16.0
8
O
16
Sn Sb Te
50
72.6
32
28.1
P
14
Si
14.0
7
N
15
12.0
6
C
14
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
8
56
72
91.2
41
50.9
23
V
7
13
Zr Nb Mo Te Ru Rh Pd Ag Cd In
40
47.9
22
Ti
6
55
71
45.0
40.1
21
39.1
20
Ca Sc
19
K
5
27.0
4
24.3
23.0
3
Al
12
10.8
11
1.0
Na Mg
Molar mass/g mol–1
5
B
9.0
1
H
6.9
Atomic Number
Be
4
3
13
Li
2
1
2
18
14
periodic table
© te ah o o te k u ra p ou n a mu
acknowledgements
Every effort has been made to acknowledge and contact copyright holders. Te Aho o Te Kura Pounamu apologises for any
omissions and welcomes more accurate information.
Extracts: PhET Interactive Simulations, University of Colorado, http://phet.colorado.edu. Extract only.
Photos
Cover: Red cabbage, © Don Laing, Wellington, NZ: Te Aho o Te Kura Pounamu. Used by permission.
Gold ring; Copper bowl; Aluminium foil, by Delene Holm, © Te Aho o Te Kura Pounamu, Wellington, NZ.
Magnesium (used twice in booklet), by Don Laing, © Te Aho o Te Kura Pounamu, Wellington, NZ.
Fluorine (close-up), © Don Laing, Wellington, NZ: Te Aho o Te Kura Pounamu. Used by permission.
Sulfur; Iodine, by Ben Mills (benjah-bmm27), Wikimedia Commons, http://commons.wikimedia.org/wiki/File:Sulfur-sample.jpg,
accessed 20 August 2012. Public domain.
Fluorine (close-up), © Don Laing, Wellington, NZ: Te Aho o Te Kura Pounamu. Used by permission.
Bromine, by Jurii, Wikimedia Commons, http://commons.wikimedia.org/wiki/File:Bromine-ampoule.jpg, accessed 20 August 2012.
Licensed under Creative Commons Attribution 3.0 Unported.
Sodium; Aluminium, by Delene Holm, © Te Aho o Te Kura Pounamu, Wellington, NZ.
Sulfur; Chlorine; Neon, by Don Laing, © Te Aho o Te Kura Pounamu, Wellington, NZ.
Si (silicon), user Ondrej Mangl, accessed 2011, Wikimedia Commons, http://commons.wikimedia.org/wiki/
File:k%c5%99em%C3%ADk.png. Public domain.
Phosphorus, by Dnn87, Wikimedia Commons, http://commons.wikimedia.org/wiki/File:Phosphor.JPG, accessed 20 August 2012.
Licensed under Creative Commons Attribution 3.0 Unported.
An alchemist carrying out experiments, from http://www.duesseldorfer-auktionshaus.de/, Wikimedia Commons, http://commons.
wikimedia.org/wiki/File:Joseph_Leopold_Ratinckx_Der_Alchemist.jpg, accessed 20 August 2012. Public domain.
Large Hadron Collider quadrupole magnets, by gamzis, Wikimedia Commons, http://commons.wikimedia.org/wiki/File:LHC_
quadrupole_magnets.jpg, accessed 20 August 2012. Licensed under Creative Commons Attribution 2.0 Generic.
Doughnut, by Delene Holm, © Te Aho o Te Kura Pounamu, Wellington, NZ.
Magnesium ribbon; Magnesium burning; White magnesium oxide, © Don Laing, Wellington, NZ: Te Aho o Te Kura Pounamu. Used by
permission.
Silver precipitates, © Don Laing, Wellington, NZ: Te Aho o Te Kura Pounamu. Used by permission.
Lead precipitates, © Don Laing, Wellington, NZ: Te Aho o Te Kura Pounamu. Used by permission.
Marble reacting with hydrochloric acid, © Don Laing, Wellington, NZ: Te Aho o Te Kura Pounamu. Used by permission.
Iron nail in blue copper sulfate solution (2 x photos), © Don Laing, Wellington, NZ: Te Aho o Te Kura Pounamu. Used by permission.
Magnesium ribbon reacting with hydrochloric acid; Burning splint, © Don Laing, Wellington, NZ: Te Aho o Te Kura Pounamu. Used by
permission.
Copper on scales, © Don Laing, Wellington, NZ: Te Aho o Te Kura Pounamu. Used by permission.
Sugar and matches, by Delene Holm, © Te Aho o Te Kura Pounamu, Wellington, NZ.
© te ah o o t e k ur a p o un a m u
CH3001
79
acknowledgements
Illustrations
Haematite, by Benjabmm-27, Wikimedia Commons, http://upload.wikimedia.org/wikipedia/commons/1/10/Haematite-unit-cell-3Dballs.png, accessed September 2012. Public domain.
All other illustrations copyright © Te Aho o Te Kura Pounamu, Wellington, NZ.
CH3001A
Forming barium sulfate and sodium chloride, © Te Aho o Te Kura Pounamu, Wellington, NZ.
80
CH3001
© te ah o o te k u ra p ou n a mu
self-assessmentch3001
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Describe the structure of the
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