Atomic Structure – Revision Pack (C4)
Atoms:
A nucleus is made up of neutral neutrons and positive protons. The nucleus is
surrounded by negatively charged electrons. The nucleus is positive but the atom
has no overall charge because the protons and electrons cancel each other out.
Subatomic
particle
Relative charge
Relative mass
Proton
+1
1
Neutron
0
1
Electron
-1
0.0005
Element – made up of one type of atom; cannot be chemically broken down
Compound – made of two or more elements that are chemically combined
Mass and Atomic Number:
19
9
F
The mass number is the larger number of the two. It represents how
many neutrons and protons there are in the atom.
The atomic number is the smaller number of the two. It represents
how many protons there are in the atom. The number of protons is
equal to the number of electrons.
An atom is neutral because it has the same number of protons as
neutrons; the positives cancel out the negatives.
The number of neutrons is equal to the mass number take away the atomic number.
Arrangement of electrons:
The elements of the periodic table are arranged in order of increasing atomic
number.
The amount of electrons is different for the shells of an atom:
Atomic Structure – Revision Pack (C4)
The maximum number of electrons for the first shell is 2.
The maximum number of electrons for all of the shells from then is 8.
Na (shown to the left) has an electronic structure. This is 2.8.1 – two
electrons in the 1st shell, 8 in the 2nd and 1 in the 3rd. The number of
electron is equal to the proton number; so 2 + 8 + 1 = 11 which is the
atomic (proton) number of sodium (Na).
Groups and Periods:
Periods =
rows =
number
of
occupied
shells
Groups = columns =
number of electrons on
outer shell
Isotopes:
Isotopes are elements that have the same atomic number, but differing mass
numbers; this means that they have different numbers of neutrons.
Isotope
The different number of neutrons
can be deduced, as with these
lithium isotopes.
Electrons
Protons
Neutrons
3
3
3
3
3
4
3
3
5
Atomic Structure – Revision Pack (C4)
Developing the Theory:
John Dalton created the original theory, which was provisional meaning it could be
changed later. The theory was later confirmed by better evidence.
J.J. Thompson, Rutherford and Bohr all found new evidence that changed the
explanation of the model of the atom. With later evidence their predictions were
confirmed.
Unexpected results also change ideas rapidly. For example, Geiger and Marsden
had some unexpected results which made significant contributions to the ideas of
Rutherford and Bohr.
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Why bonding?
To be stable, atoms must have a full outer shell. Atoms can be made stable by
transferring electrons and this is called ionic bonding.
Ionic bonding ONLY happens between the group 1 (alkali) metals and non-metals.
Ions are charged particles, they are either:
-
Positive – if they have lost an electron (e.g. Na+ or Cu2+)
Negative – if they have gained (or taken) an electron (e.g. O2- or Cl-)
Non-metals gain electrons to help them to get a stable electronic structure. If an
atom gains an electron it becomes negatively charged because there are more
negative electrons than there are positive protons.
Metals gain electrons to help them to get a stable electronic structure. If an atom
loses an electron it becomes positively charged because there are more positive
protons in the nucleus than there are negative electrons.
When ionic bonding happens, the metals become negative ions and the nonmetals become negative ions. The positive and negative ions are then attracted to
one another.
Dot and Cross diagrams are used to show ionic bonding:
This is the bonding of magnesium oxide (MgO).
Mg has the electronic structure 2.8.2
O has the electronic structure 2.6
For Mg to gain stability it has to lose two electrons.
For O to gain stability it has to gain two electrons.
When they reactant, there is transfer of electrons
and both Mg and O are now stable.
YOU must be able to draw these!
Ionic compounds:
Positive and negative have to cancel out. Some examples of this include:
Na+ and Cl- - cancel out to make NaCl
Cu2+ and Cl- - you need two Cl to cancel out the Cu2+ so makes CuCl2
Atomic Structure – Revision Pack (C4)
Mg2+ and O2- - cancel out to make MgO
Mg2+ and Cl- - you need two Cl to cancel out the Mg2+ so makes MgCl2
Na+ and O2- - you need two Na to cancel out the O2- so makes Na2O
There is a big difference between the atom of an
element and the ion of an element. For example,
for a sodium atom to become stable it has to lose
one electron from its outer shell. This forms a
sodium ion which is fully stable.
However, for negative ions there is a slight
difference in what the ion is called; if an atom has
to gain electrons to become stable its ion ends in
‘ide’. For example, a fluorine atom needs one
electron to become stable or have a full outer
shell. This forms an ion called a fluoride ion.
Giant Ionic Lattice:
The structure of compounds like NaCl and MgO is a giant ionic lattice. In this, positive
ions have strong electrostatic forces to negative ions. They ALWAYS exist as solids.
The strong electrostatic forces
mean that these compounds
have very high melting points.
They can conduct electricity
when molten or in solution
because the ions are free to
move.
Magnesium Oxide
High melting point (strong bonds) – used
in fire-resistant materials
Sodium Chloride
High melting point (strong bonds)
Insoluble in water
Dissolves in water
Conducts electricity when molten (ions
free to move)
Conducts electricity when molten and
in solution (ions free to move)
The melting point of different compounds is dependent on the amount of electrons
that are transferred. For example, in MgO there are Mg2+ ions as opposed to in NaCl
Atomic Structure – Revision Pack (C4)
where there is just Na+ ions. This means there is a stronger electrostatic attraction and
because two electrons have been donated there is a stronger bond.
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Metal Properties:
Metal properties can be both chemical and physical.
Chemical Properties
- Resistance to attack by oxygen
or other acids (shown by gold
and aluminium) – makes it useful
for car bodies, so they are not
oxidised by rain
Physical Properties
- High thermal conductivity (shown
by copper) – which is good for
pots and pans
- Lustrous and shiny (shown by
aluminium & most other metals) –
aesthetically pleasing
- Malleable – can be hammered
or pressed in to shape without
breaking or cracking (shown by
most metals) – good for making
things like coins
- Ductile – meaning it can be
deformed and maintain
Atomic Structure – Revision Pack (C4)
-
toughness (like steel) – used in
wires
Low Density (shown by
aluminium) – lightweight so is
often used in car bodies or
aircrafts
Metallic Bonding:
Metals will generally always have high melting and boiling points. This is because of
the strong metallic bonds that they possess. The bonding between atoms is very
strong and a lot of energy is required to break these bonds.
Metals like copper, gold and silver all conduct electricity. When they do this, the
electrons in the metal move generating a charge.
A metallic bond is a strong electromagnetic force
between positively charged and tightly-packed
metal ions and a ‘sea’ of delocalised electrons.
(NOTE - ELECTROMAGNETIC FORCE is generated
by differences in electrical charge – here the ions
are positive and the electrons are negative!)
The delocalised electrons can move through the
metal very easily – this is why metals can conduct
electricity.
Metals have high melting and boiling points
because a great deal of energy is required to
break and overcome the STRONG attraction
between delocalised electrons and close-packed
metal ions.
Superconductors:
A superconductor is a material that conducts electricity with little or no resistance.
Although copper, gold and silver are good conductors of electricity they DO NOT
become superconductors.
Mercury does become a superconductor, but at a very low temperature (-268.8oC).
At this temperature the electrical resistance falls to zero.
After a substance goes from its normal state
to an electromagnetic state, it no longer
has magnetic fields inside it. When you put
a small magnet next to a superconductor,
it is repelled. If you put a small permanent
magnet above a superconductor, it
levitates (as seen to the left).
Atomic Structure – Revision Pack (C4)
There are many potential benefits of superconductors, including:
-
Loss-free power transmission
Super-fast electronic circuits
Powerful electromagnets
However, there are also many drawbacks with superconductors as well, for
example:
-
They work at very low temperatures and this limits their uses
However, scientists are working on a superconductor that works at room
temperature (20 degrees Celsius).
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Purifying Water
River water is often cloudy and not suitable to drink. To make it in to clean and
drinkable water, it goes through a ‘water purification works’.
Some pollutants get into water before purification. For example, nitrates often enter
rivers from nitrogen in fertilisers and pesticides from crop spraying to kill insects often
gets into the water prior to purification.
Occasionally, some pollutants also get into the water AFTER purification. For
example, some older houses will have lead water pipes. Gradually, the lead
dissolved into the water.
Water that has not yet been purified may also contain: microbes, leaves, salts and
minerals. This is why purification must take place; this happens in three simple steps:
Sedimentation:
Chemicals are added
to make solid particles
and bacteria settle
out.
Filtration:
Fine particles are
filtered (and thus
removed). A layer of
sand on gravel filters
Chlorination:
Finally, chlorine is
added to kill any
remaining microbes
within the water. This
Atomic Structure – Revision Pack (C4)
These steps do cost money, so water must be conserved (not wastefully overused).
This can be used to provide clean water.
Water is RENEWABLE (i.e. not used up or depleted when used) but the supply is
certainly not endless. If it doesn’t rain enough in the winter then reservoirs don’t fill
up. Also, in the UK more and more homes are being built which increases the
demand for water. We also need to be careful as it takes energy to pump and purify
water and this contributes to global warming (an increase in the Earth’s
temperature).
On occasion, soluble substances are not removed by water purification because
they are dissolved in the water. Examples include nitrates from fertilisers and
pesticides from crop spraying. Extra processes are required to remove these
substances.
This cannot happen with water from the sea because there are too many
substances dissolved within it. To remove these, further distillation is required which
required huge amounts of energy meaning that the process is very expensive. This is
why sea water is only purified in rare cases, like when there is NO fresh water.
Testing Water:
Water can be tested using a precipitation reaction. Two aqueous solutions are used:
Barium Chloride (BaCl2) or Silver Nitrate (AgNO3).
In a precipitation reaction, two solutions react to form an insoluble chemical. This
chemical will turn from liquid to solid (this is known as the precipitate) and from the
colour you can identify what chemical was present.
If you are testing for SULPHATE ions:
STEP 1 – add the water sample to a test tube
STEP 2 – add some barium chloride (BaCl2)
STEP 3 - if any SULPHATE ions are present a white precipitate is
formed (as seen to the left).
The balanced symbol equation for this reaction would be:
BaCl2 + MgSO4  BaSO4 + MgCl2
Barium Chloride + Magnesium Sulphate  Barium Sulphate +
Magnesium Chloride
Precipitate
Atomic Structure – Revision Pack (C4)
If you were testing for CHLORIDE, BROMIDE or IODIDE ions:
STEP 1 – add the water sample to a test tube
STEP 2 – add some Silver Nitrate (AgNO3)
STEP 3 – If any:
-
CHLORIDE ions are present, a white precipitate is
formed (1st in picture to left)
BROMIDE ions are present, a cream precipitate is
formed (2nd in picture to left)
IODIDE ions are present, a yellow precipitate is
formed (3rd in picture to left)
The balanced symbol equations for the reaction with silver nitrate would be:
AgNO3 + NaCl  AgCl + NaNO3 (CHLORIDE ions – white precipitate)
AgNO3 + NaBr  AgBr + NaNO3 (BROMIDE ions – cream precipitate)
AgNO3 + NaI  AgI + NaNO3 (IODIDE ions – yellow precipitate)
NOTE – this is used to test water – do NOT forget that!
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6
(a) This question is about drinking water.
Washing-up liquid can pollute drinking water.
Another pollutant is nitrates.
Suggest how nitrates could get into drinking water.
................................................................................................................................. [1]
(b) Joe is testing some water samples.
He adds silver nitrate solution and barium chloride solution to different water
samples.
Look at the table. It shows his results.
Water sample
Result with silver nitrate solution
Result with barium chloride
solution
A
White Precipitate
Colourless Solution
B
Cream Precipitate
Colourless Solution
C
Colourless Solution
White Precipitate
......................................................
.......................................................
D
(i)
What water sample contains bromide ions?
...................................................................................................................... [1]
(ii)
What water sample contains chloride ions?
...................................................................................................................... [1]
(iii)
Sample D contains iodide and sulphate ions. Complete the table.
...................................................................................................................... [1]
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Question
Answer
Marks
Guidance
Atomic Structure – Revision Pack (C4)
9 (a)
(Run off) from
1
Allow sewage / NPK
fertilisers (1)
9 (b) (i)
B (1)
1
9 (b) (ii)
A (1)
1
9 (b) (iii)
Silver nitrate – yellow
1
Both required for the mark
precipitate and
barium chloride –
white precipitate (1)
PPQ(5):
Question
4 (a)
Answer
Insoluble material is
Marks
1
removed / named
Guidance
Allow twigs / leaves / sand / dirt removed
Ignore small particles removed
insoluble material
removed (1)
4 (b)
Soluble substances
1
Ignore lead particles present by allow lead
are not removed by
ions present (1)
water purification /
Allow there are still dissolved substances
AW (1)
present / soluble pollutants still present /
nitrates present / pesticides present (1)
4 (c) (i)
Distillation (1)
1
Allow evaporation and condensation (1)
Allow reverse osmosis (1)
4 (c) (ii)
Large energy
1
Reject uses fossil fuels (0) but needs a lot
requirement / need to
of fossil fuels for heating (1)
boil a lot of water /
Ignore references to other costs (0)
needs lots of heat /
needs lot of electricity
(1)
Properties of Alkali Metals:
The group 1 metals are also known as the ‘alkali metals’. They include caesium,
rubidium, lithium, sodium and potassium. They all have similar properties because
they come from group one and have one electron on their outer shell. These
properties include:
-
They react vigorously with water
Hydrogen is given out
Atomic Structure – Revision Pack (C4)
-
The metal reacts with water to form an alkali (the hydroxide of the metal)
You can predict the reactivity of certain
elements by looking at the pattern of reactivity
in other alkali metals.
The balanced symbol equation for Li and H2O
is:
2Li + 2H2O  2LiOH + H2
Reactivity of Group 1 Elements:
Atoms of Group 1 elements will have similar properties because they
have one electron in their outer shell.
When the atoms of alkali metals react, they lose one electron and
form:
-
A full outer shell, so have a stable electronic structure
A positive ion which has more positive charge in its nucleus
than negative charge from the surrounding shells
This can be represented through the equation:
Na – e-  Na+
The closer the outer shell’s electron is
to the nucleus, the more attractive
force there is, so the electron is less
easily lost. Na is more reactive than
Li.
The easier it is for an atom of an alkali
metal to lose its electron, the more
reactive it is.
Oxidation:
Oxidation Is Loss; for example:
K
-
e-

K+
An atom of potassium loses an electron to form a positive ion. This is an example of
oxidation.
Flame Tests:
Atomic Structure – Revision Pack (C4)
A flame test is used to find out if lithium, sodium and potassium are present in a
compound.
Alkali metal in the compound
Lithium
Sodium
Potassium
Colour of flame
Red
Yellow
Lilac
How to do it:
STEP 1 – Put on some safety goggles
STEP 2 – A flame-test wire is moistened with dilute HCl
STEP 3 – The flame-test wire is then dipped into a solid chemical
STEP 4 – The flame-test wire is then put in a blue Bunsen Burner flame
STEP 5 – You can then identify which alkali metal is present in the compound
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The Halogens:
The Group 7 Elements all have similar properties because they all have seven
electrons in their outer shells. They have similar properties because they are trying to
gain one electron when they react. In the reaction with Group 1 elements, the
halogens form a negative ion and gain a full electronic structure.
Group 7 Element
Chlorine
Bromine
Iodine
Colour and state at room temperature
Green gas
Orange liquid
Grey solid
It is possible to predict the properties of
group 7 elements if you know properties of
the other halogens because they follow a
particular trend.
Atomic Structure – Revision Pack (C4)
The nearer the outer shell is to the nucleus, the easier it is for an atom to gain one
electron. Therefore, the easier it is to gain the electron, the more reactive the
halogen. Fluorine is more reactive than Chlorine.
Reactions of Halogens:
When a halogen reacts with an alkali metal, a
metal halide is made. For example, when
potassium reacts with iodine, the metal halide
formed is potassium iodide. Word equation:
Potassium + iodine  Potassium Iodide
OR
2K + I2  2KI
The ‘ide’ is always the group 7 element.
Another equation is:
2Li
+
Br2
Lithium is the alkali metal.

2LiBr
Displacement Reaction:
Halogen
Gas (R)
Chlorine
Bromine
Chlorine
Bromine is the halogen.
Lithium Bromide is the metal halide.
Metal Halide (R)
Metal Halide (P)
Potassium
Bromide
Potassium
Iodide
Potassium
Iodide
Potassium
Chloride
Potassium
Bromide
Potassium
Chloride
Halogen
Gas (P)
Bromine
Colour
Iodine
Red/Brown
Solution
Red/Brown
Solution
Iodine
Orange Solution
REMEMBER:
Bromine can displace iodine.
Chlorine can displace bromine and iodine.
Iodine cannot displace either because it is the least reactive out of the three.
Atomic Structure – Revision Pack (C4)
BECAUSE displacement is dependent on reactivity, we can predict the reactions
between halogen and metal halides.
Reduction:
REMEMBER - Reduction Is Gain, as shown in this ionic equation:
Br2
+
-
2e-

2Br-
A molecule of bromine gains two electrons (one for each atom)
It becomes two bromide ions that are negative ions
This is reduction
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Covalent Bonding:
Non-metals can share electron pairs between atoms – this is known as covalent
bonding. The idea of covalent bonding is that when the elements react, they gain
electrons to fill their outer shell and become stable.
When the atoms join together, there is a strong bond that holds them together –
covalent bonds. The formation of simple molecules that contain single and double
bonds can be represented by ‘dot and cross’ diagrams which only show the outer
shells electrons. For example:
H2O has two single covalent bonds. Two
pairs of electrons are shared.
Oxygen atoms can each form two
covalent bonds. Two pairs of electrons
are shared in an O2 molecule – this is a
Atomic Structure – Revision Pack (C4)
Predicting chemical properties:
The attraction between molecules like carbon dioxide and water is called an
intermolecular force.
The dot and cross diagrams for water and carbon dioxide represent how ALL the
atoms bond to make a molecule – this means that you have to show ALL the
electrons in each of the atoms.
Carbon dioxide is made up of two oxygen and one
carbon atom. The carbon atom has four electrons in its
outer shell, so it needs four more. Oxygen atoms have
six electrons in its outer shell, so needs to more. Two
double covalent bonds are created – each oxygen
outer shell is shared with two of the electrons from the
carbon outer shell. This way all of the atoms have a full
outer shell.
The simple molecular structures, like CO2, have weak intermolecular forces so are
very easy to break apart so the substances have low melting points.
In covalent bonding there are no free electrons so these molecules do NOT conduct
electricity.
Groups and Periods:
The group number (columns) is the same as the number of electrons in the outer
shell.
The period number (rows) is the number of shells needed for all the electrons.
The development of the periodic table:
Scientist
Discovery
Newlands
1865 – Newlands puts 56 elements into groups and noticed
that every 8 elements behaved similarly. This was not
accepted for another 50 years until other scientists
discovered more evidence.
Mendeleev
1869 – Mendeleev put all the elements in order in a table. He
noted periodic changes in the elements and made the
prediction that new elements would be discovered.
1891 – Mendeleev did not include the noble gases in his
periodic table.
Atomic Structure – Revision Pack (C4)
After 1891 – Later investigations by other scientists confirmed
Mendeleev’s idea of periodicity. His prediction that other
elements would be found are correct.
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What are the Transition Metals?
The transition elements are a block of metallic
elements that are located between group 2
and group 3 of the periodic table. This block
contains elements such as iron (Fe), copper
(Cu), platinum (Pt) and zinc (Zn).
The transition elements have typical metallic
properties and their compounds are often
coloured, for example:
-
Copper compounds are blue
Iron (II) compounds are pale green
Iron (III) compounds are often orange or
brown
Atomic Structure – Revision Pack (C4)
Transition elements and their compounds are often used as catalysts, for example:
-
Iron (Fe) is used in the Haber process to make ammonia, which is used in
fertilisers
Nickel (Ni) is used in the manufacture of margarine to harden the oils
Thermal Decomposition of metal carbonates:
If a transition metal carbonate is heated it undergoes thermal decomposition to
form metal oxide and carbon dioxide; you can adapt this formula to any transition
metal:
(Transition) Metal Carbonate  Metal Oxide + Carbon Dioxide
Some examples of thermal decomposition include:
-
Iron carbonate  Iron Oxide + Carbon Dioxide
Nickel carbonate  Nickel Oxide + Carbon Dioxide
Zinc carbonate  Zinc Oxide + Carbon Dioxide
The metals will generally change colour during this reaction.
You should also be able to write balanced symbol equations for this reaction, for
example:
-
FeCO3  FeO + CO2
NiCO3  NiO + CO2
ZnCO3  ZnO + CO2
The CO3 part is the carbonate.
Metal
Metal
C
O
+
O
C
The O part is the oxide.
The CO2 part is the carbon dioxide.
O
O
O
O
Atomic Structure – Revision Pack (C4)
This is an experiment to
test whether thermally
decomposing a transition
element makes carbon
dioxide. During heating, a
transition element should
ALWAYS give off carbon
dioxide! The delivery tube
should send the gas
produced down into the
limewater, if it turns milky
then carbon dioxide was
made.
Precipitation Reactions with Sodium Hydroxide (Na+OH-) solution:
Sodium hydroxide is added to solutions to identify what transition elements are
present. The NaOH solution reacts with compounds of each transition metal to make
a solid of a certain colour, for example:
-
Cu2+ ions form a blue solid
Fe2+ ions form a grey / green solid
Fe3+ ions form a orange / brown solid
You will often be asked to write a balanced symbol equation of a precipitate
reaction, for example:
-
Cu2+ + 2OH-  Cu(OH)2
Fe2+ + 2OH-  Fe(OH)2
Fe3+ + 3OH-  Fe(OH)3
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PPQ(1):
PPQ(2):
11
Iron and copper are transition elements.
(a)
Brahim adds a small volume of sodium hydroxide solution to five different solutions.
An insoluble solid called a precipitate is made each time.
Look at the results table. It is not finished.
Solution
Copper chloride
Formula
Colour of precipitate made
CuCl2
Blue
Copper nitrate
Cu(NO3)2
.......................................................
Iron (II) chloride
FeCl2
Green
Iron (II) sulphate
FeSO4
Green
Iron (III) nitrate
Fe(NO3)3
......................................................
(i)
Finish the table.
(ii)
Look at the formulas in the table.
[2]
Which formula contains six oxygen atoms?
Choose from the table.
......................................................................................................... [1]
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Question
7 (c)
Answer
Fe(OH)2
Marks
Guidance
1
PPQ(2):
Question
11 (a) (i)
Answer
Copper nitrate – blue
Marks
Guidance
2
(1)
Iron (III) nitrate –
Allow foxy red / red-brown (1)
orange / rusty / brown
(1)
11 (a) (ii)
Ignore red or yellow
Copper nitrate /
1
Cu(NO3)2 (1)
Ignore CuNO3
PPQ(3):
Question
Answer
11 (a) (iii)
Cu2+ + 2OH- 
Marks
2
Guidance
Allow any correct multiple
Cu(OH)2
Allow 1 mark for correct balancing with
Correct formula (1)
minor errors of case or subscript
Correct balancing (1)
Allow = instead of arrow
Ignore state symbols
Not + for and
PPQ(4):
Question
8 (c)
Answer
Fe3+ +3OH- 
Fe(OH)3 (1)
Marks
1
Guidance
Look for balancing on the printed
examination as well as the answer line.