The Periodic Table

 There are over 100 chemical elements which have been isolated and
identified
o Each element has one proton more than the element preceding it

o This is done so that elements end up in columns with other elements

which have similar properties
o Elements are arranged on the periodic table in order of increasing
atomic number
o The table is arranged in vertical columns called groups and in rows
called periods
 Period: These are the horizontal rows that show the number of shells of
electrons an atom has and are numbered from 1 - 7
o E.g. elements in period 2 have two electron shells, elements in period
3 have three electron shells
 Group: These are the vertical columns that show how many outer electrons
(also known as valency electrons) each atom has and are numbered from I –
VII, with a final group called Group 0 (instead of Group VIII)
o E.g. Group IV elements have atoms with 4 electrons in the outermost
shell, Group VI elements have atoms with 6 electrons in the outermost
shell and so on
 The group number can help determine the charge that metal and non-metal
ions form
 For metals, the group number corresponds to the number of electrons it will
lose to achieve a full outer shell and the charge of the metal ion
o E.g. sodium is in Group I, it will lose 1 electron and form an ion with
a 1+ charge
o Magnesium is in Group II, it will lose 2 electrons and form an ion with
a 2+ charge
 For non-metals in Group VII and VI, they will gain 1 and 2 electrons
respectively to gain a full outer shell
o E.g. non-metals in Group VII gain 1 electron to form ions with a 1-
charge
o Non-metals in Group VI gain 2 electrons to form ions with a 2- charge
All elements are arranged in the order of increasing atomic number from
left to right
Valency
 Valency (or combining power) tells you how many bonds an atom can make
with another atom or how many electrons its atoms lose, gain or share, to
form a compound
o E.g. carbon has a valancy of 4 as it is in Group IV so a single carbon
atom can share 4 electrons to make 4 single bonds or 2 double bonds
 The following valencies apply to elements in each group:
Exam Tip
An easier way of remembering which number is the mass number and which is the
atomic is:
Mass Number = The massive number i.e the larger of the two numbers.
The atomic number must be the smaller number.
The Metallic Character of Elements
 The metallic character of the elements decreases as you move across a
Period on the Periodic Table, from left to right, and it increases as you
move down a Group
 This trend occurs due to atoms more readily accepting electrons to fill
their valence shells rather than losing them to have the previous, already
full, electron shell as their outer shell
 Metals occur on the left-hand side of the Periodic Table and non-metals on
the right-hand side
 Between the metals and the non-metals lie the elements which display some
properties of both
 These elements are referred to as metalloids or semi-metals
Properties of metals and non-metals
A zig-zag line in this diagram separates the metals on the left, from the
non-metals on the right
Periodic Trends & Electronic Configuration
 The electronic configuration is the arrangement of electrons into shells for
an atom (e.g: the electronic configuration of carbon is 2,4)
 There is a link between the electronic configuration of the elements and their
position on the Periodic Table
 The number of notations in the electronic configuration will show the number
of occupied shells of electrons the atom has, showing the period
 The last notation shows the number of outer electrons the atom has, showing
the group number
Example: Electronic configuration of chlorine:
The electronic configuration of chlorine as it should be written

Period: The red numbers at the bottom show the number of notations which is 3,
showing that a chlorine atom has 3 shells of electrons.
Group: The final notation, which is 7 in the example, shows that a chlorine atom
has 7 outer electrons and is in Group VII

The position of chlorine on the Periodic Table

 Elements in the same group in the Periodic Table have similar chemical
properties
  When atoms collide and react, it is the outermost electrons that interact
 The similarity in their chemical properties stems from having the same
number of electrons in their outer shell
 For example, both lithium and sodium are in Group I and can react with
elements in Group VII to form an ionic compound (charges of Group I ions are
1+, charges of Group VII ions are 1-) by reacting in a similar manner and each
donating one electron to the Group VII element
 As you look down a group, a full shell of electrons is added to each
subsequent element
o Lithium's electronic configuration: 2,1

o Sodium's electronic configuration: 2,8,1

o Potassium's electronic configuration: 2,8,8,1

Exam Tip
Electronic configurations can be shown with the numbers separated by commas or
by full stops. In this course commas are used, but you will often see full stops used
elsewhere. Both are accepted.
Predicting Properties
 Because there are patterns in the way the elements are arranged on the
Periodic Table, there are also patterns and trends in the chemical behaviour
of the elements and their physical properties
 These trends in properties occur down groups and across the periods of the
Periodic Table
 As a result, we can use the Periodic Table to predict properties such as:

o boiling point

o melting point

o density

o reactivity

 Some common properties / trends in properties include:

o Group I elements react very quickly with water

o Noble gases are unreactive

o Transition elements are denser than Group I elements

o Reactivity decreases going down Group VII

 o Melting point decreases going down Group I

 In this way the Periodic Table can be used to predict how a particular
element will behave
Identifying Trends
 Using given information about elements, we can identify trends in properties
 An example of when this might be used is to determine the trend in reactivity
of Group I metals
 The table below shows the reactions of the first three elements in Group I
with water
Observations of Lithium, Sodium, and Potassium with Water

 The observations show that reactivity of the Group I metals increases as you
go down the group
  Using this information we can predict the trend going further down Group I for
the elements rubidium, caesium and francium
 As the reactivity of alkali metals increases down the group, rubidium,
caesium and francium will react more vigorously with air and water than
lithium, sodium and potassium
 Lithium will be the least reactive metal in the group at the top, and francium
will be the most reactive at the bottom
 Francium is rare and radioactive so is difficult to confirm predictions
Table to Show the Predicted Reaction of other Group I Elements with
Water

Exam Tip
You may be asked to identify other trends in chemical or physical properties of
Group I metals, given appropriate data.
Firstly, ensure that the metals and associated data are written in either descending
or ascending order according the their position in the Group. Then look for general
patterns in the data.
Group I Properties & Trends: Basics
The Group I metals
 The Group I metals are also called the alkali metals as they form alkaline
solutions with high pH values when reacted with water
 Group I metals are lithium, sodium, potassium, rubidium, caesium and
francium
 They all contain just one electron in their outer shell
Physical properties of the Group I metals
 The Group I metals:
o Are soft and easy to cut, getting even softer and denser as you move
down the Group (sodium and potassium do not follow the trend in
density)
o Have shiny silvery surfaces when freshly cut

o Conduct heat and electricity

o They all have low melting points and low densities compared to other
metals, and the melting point decreases as you move down the
Group; some would melt on a hot day

The alkali metals lie on the far left-hand side of the Periodic Table
Chemical properties of the Group I metals
 They react readily with oxygen and water vapour in air so they are stored
under oil to stop them from reacting
 Group I metals will react similarly with water, reacting vigorously to produce
an alkaline metal hydroxide solution and hydrogen gas
 The Group I metals get more reactive as you look down the group, so only the
first three metals are allowed in schools for demonstrations
Reactions of the Group I metals and water

Predicting the Properties of Group I Elements

 Knowing the reactions of elements at the top of the group allows you to
predict the properties of other elements further down Group I
Properties of other Alkali Metals (Rubidium, Caesium and Francium)
  As the reactivity of alkali metals increases down the group, rubidium,
caesium and francium will react more vigorously with air and water than
lithium, sodium and potassium
 Lithium will be the least reactive metal in the group at the top, and francium
will be the most reactive at the bottom
 Francium is rare and radioactive so is difficult to confirm predictions
 For example the reactions with water can be predicted:
Predicting the Reaction with Water

 You can also look at other properties such as boiling point, melting point and
density of Group I elements and use them to predict whether the other
properties are likely to be larger or smaller going down the group
Group VII Properties & Trends
The halogens
 These are the Group VII non-metals that are poisonous and include fluorine,
chlorine, bromine, iodine and astatine
 Halogens are diatomic, meaning they form molecules of two atoms
o The formulae of the halogens are F2, Cl2, Br2, I2 and At2

 All halogens have seven electrons in their outer shell

 They form halide ions by gaining one more electron to complete their outer
shells
 Fluorine is not allowed in schools so observations and experiments tend to
only involve chlorine, bromine and iodine
Properties of the halogens
 At room temperature (20 °C), the physical state of the halogens changes as
you go down the group
 Chlorine is a pale yellow-green gas, bromine is a red-brown liquid and
iodine is a grey-black solid
 This demonstrates that the density of the halogens increases as you
go down the group:

The physical state of the halogens at room temperature

 Reactivity of Group VII non-metals increases as you go up the group (this
is the opposite trend to that of Group I)
 Each outer shell contains seven electrons and when the halogen reacts, it will
need to gain one outer electron to get a full outer shell of electrons
  As you go up Group VII, the number of shells of electrons decreases (period
number decreases moving up the Periodic Table)
 This means that the outer electrons are closer to the nucleus so there
are stronger electrostatic forces of attraction, which help to attract the extra
electron needed
 This allows an electron to be attracted more readily, so the higher up the
element is in Group VII then the more reactive it is

Diagram showing the electronic configuration of the first three elements

in Group VII
Exam Tip
Solid iodine, iodine in solution and iodine vapour are different colours. Solid iodine is
dark grey-black, iodine vapour is purple and aqueous iodine is brown.
Predicting Group VII Properties
 You may be given information about some elements and asked to predict the
properties of other elements in the group
 The information you might be given could be in relation to melting/boiling
point or physical state/density so it is useful to know the trends in properties
going down the group
Melting and boiling point
 The melting and boiling point of the halogens increases as you go down the
group
 Fluorine is at the top of Group VII so will have the lowest melting and boiling
point
  Astatine is at the bottom of Group VII so will have the highest melting and
boiling point
Physical states
 The halogens become denser as you go down the group
 Fluorine is at the top of Group VII so will be a gas
 Astatine is at the bottom of Group VII so will be a solid
Colour
 The colour of the halogens becomes darker as you go down the group
 Fluorine is at the top of Group VII so the colour will be lighter, so fluorine
is yellow
 Astatine is at the bottom of Group VII so the colour will be darker, so astatine
is black
Exam Tip
You can be asked to identify trends in chemical or physical properties of the Group
VII elements, given appropriate data.
Firstly, make sure that you have placed the elements and associated data in either
ascending or descending order according to their position in Group VII. Then look for
any general patterns in the data.

Group VII Displacement Reactions

 A halogen displacement reaction occurs when a more reactive halogen
displaces a less reactive halogen from an aqueous solution of its halide
 The reactivity of Group VII non-metals increases as you move up the group
 Out of the three commonly used halogens, chlorine, bromine and iodine,
chlorine is the most reactive and iodine is the least reactive
Colour of Halogens in Aqueous Solutions
Halogen displacement reactions
Chlorine and bromine
 If you add chlorine solution to colourless potassium bromide solution, the
solution becomes orange as bromine is formed
 Chlorine is above bromine in Group VII so is more reactive
 Chlorine will therefore displace bromine from an aqueous solution of the
metal bromide
 The least reactive halogen always ends up in the elemental form
potassium bromide + chlorine → potassium chloride + bromine
2KBr (aq) + Cl2 (aq) → 2KCl (aq) + Br2 (aq)
Bromine and iodine
 Bromine is above iodine in Group VII so is more reactive
 Bromine will therefore displace iodine from an aqueous solution of metal
iodide
 The solution will turn brown as iodine is formed
magnesium iodide + bromine → magnesium bromide + iodine
MgI2 (aq) + Br2 (aq) → MgBr2 (aq) + I2 (aq)
Summary table of displacement reactions

Exam Tip
Iodine solid, solution and vapour are different colours. Solid iodine is dark grey-
black, iodine vapour is purple and aqueous iodine is brown.

Transition Elements
General properties of the transition elements
 They are very hard and strong metals and are good conductors
of heat and electricity
 They have very high melting points and are highly dense metals
 For example, the melting point of titanium is 1,688ºC whereas potassium in
Group I melts at only 63.5ºC, slightly warmer than the average cup of hot
chocolate!
 The transition elements form coloured compounds and often have more
than one oxidation state, such as iron readily forming compounds of both
Fe2+ and Fe3+
 These coloured compounds are responsible for the pigments in many paints
and the colours of gemstones and rocks
 Transition elements, as elements or in compounds, are often used
as catalysts to improve the rate or reaction in industrial processes
o Transition element catalysts of platinum or rhodium are also used in
car exhausts in the 'catalytic convertor' to reduce the levels of nitrous
oxides and carbon monoxide produced

The transition elements on the Periodic Table

  The transition elements have more than one oxidation number, as they can
lose a different number of electrons, depending on the chemical environment
they are in
 For example. iron either:
o Lose two electrons to form Fe2+ so has an oxidation number of +2

o Loses three electrons to form Fe3+ so has an oxidation number of +3

 Compounds containing transition elements in different oxidation states will

have different properties and colours

Ions of the same element can have different oxidation numbers forming
different colours
Uses of the transition elements
 The transition elements are used extensively as catalysts due to their ability
to interchange between a range of oxidation states
 This allows them to form complexes with reagents which can
easily donate and accept electrons from other chemical species within a
reaction system
 They are used in medicine and surgical applications such as limb and joint
replacement (titanium is often used for this as it can bond with bones due to
its high biocompatibility)
 They are also used to form coloured compounds
in dyes and paints, stained glass jewellery
Exam Tip
Although scandium and zinc are in the transition element area of the Periodic Table,
they are not considered transition elements as they do not form coloured
compounds and have only one oxidation state.

Noble Gases Properties & Electronic Configuration

The Noble Gases
 The noble gases are in Group VIII (or Group 0); they are non-metals and have
very low melting and boiling points
 They are all monoatomic, colourless gases
 The Group 0 elements all have full outer shells
 This electronic configuration is extremely stable so these elements are
unreactive and are inert
 Electronic configurations of the noble gases:
o He: 2

o Ne: 2,8

o Ar: 2,8,8

o Kr: 2,8,18,8

o Xe: 2,8,18,18,8
Noble gases are inert (unreactive) as they have a full outer shell of
electrons so do not easily lose or gain electrons