The Periodic Table
The periodic table
• is a table of the chemical elements in which the
elements are arranged by order of their atomic
• it is arranged so that the periodic properties
(repeating properties) of the elements are made
• The standard form of the table includes periods
(usually horizontal in the periodic table)
• and groups (usually vertical).
• Elements in groups have some similar properties
to each other.
• This has a quiz you might like to try
Describe the way the Periodic Table classifies
elements in order of proton number.
• This link gives you some history of the periodic
• In the periodic table the elements are aranged
in order of their proton number
Describe the change from metallic to
non-metallic character across a period
• Non-metallic properties increase from left to
• On the left elements have only a few electrons
in the outer shell
• And become stable by giving them away
• On the right they have more in the outer shell
and become stable by accepting electrons to
fill the shell
• There are three types of elements
• Here is the boundary between them
Periodic trends
Periodic properties
• Elements in the same group have similar
• In groups one and two the reactivity of the
elements increases down the group
• On the other side of the table in group 7 the
reactivity decreases down the group
• The noble gases (group 0) have a full outer
shell of electrons and won’t react with any
other element
• You should be able to draw out the electronic
structures for the first 20 elements if you know
the proton and nucleon number
Describe the relationship between Group number,
number of outer-shell (valency) electrons and
metallic/non-metallic character.
• The group number tells us how many
electrons are in the outer shell
• Metals donate electrons when they ract
• Non-metals accept electrons when they react
• The number of electrons that can be given
away or accepted is known as the valency
Use the Periodic Table to predict properties of
elements by means of groups and periods.
• In each group the elements have similar
• eg in group 1 the reactivity increases down
the group, but all show similar reactions
Group properties
Describe lithium, sodium and potassium in Group I as a
collection of relatively soft metals showing a trend in melting
point and reaction with water.
• This video shows you the reaction of the group
1 metals with water
Group 1 properties
• All soft metals that are easily cut with a scalpel or
• The freshly cut surface is a shiny, silver colour, it
quickly tarnishes to a dull grey as the metal reacts
with oxygen and water in the air.
• Pieces of such metals are stored in oil to prevent these
• The shiny surface of sodium tarnishes more quickly
than that of lithium.
• And potassium tarnishes more quickly than sodium.
• This shows the increasing reactivity of the metals as
we go down the group.
Reaction with cold water
• All the alkali metals react vigorously with cold
• In each reaction, hydrogen gas is given off
and the metal hydroxide is produced.
• The speed and violence of the reaction
increases as you go down the group.
• This shows that the reactivity of the alkali
metals increases as you go down Group 1.
• When lithium is added to water, lithium floats.
It fizzes steadily and becomes smaller, until it
eventually disappears.
• lithium + water → lithium hydroxide + hydrogen
• 2Li(s) + 2H2O(l) → 2LiOH(aq) + H2(g)
• When sodium is added to water, the sodium
melts to form a ball that moves around on the
• It fizzes rapidly, and the hydrogen produced may
burn with an orange flame before the sodium
• sodium + water → sodium hydroxide + hydrogen
• 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
• When potassium is added to water, the metal melts
and floats.
• It moves around very quickly on the surface of the
• The hydrogen ignites instantly.
• The metal is also set on fire, with sparks and a lilac
• There is sometimes a small explosion at the end of
the reaction.
• potassium + water → potassium hydroxide +
• 2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
Predict the properties of other elements in
Group I, given data where appropriate
• You may be given data and asked to predict
the properties of other members of the group
Group 7 – the halogens
• The elements in Group 7 of the Periodic Table
are called the halogens.
• They include chlorine, bromine and iodine.
The halogens are diatomic - this means they
exist as molecules, each with a pair of atoms.
Chlorine molecules have the formula Cl2,
bromine Br2 and iodine I2.
Describe the trends in properties of chlorine, bromine and iodine
in Group VII including colour, physical state and reactions with
other halide ions.
• The electronegativity of halogens decrease down
the group.
• Fluorine is the most reactive
• The halogens boiling points increase down group
• This is because the atoms get bigger and so the
Van der Waals forces get smaller.
• So fluorine is a gas, bromine is a liquid and
iodine is a solid
• The oxidising power of the halogens
decreases from Fluorine to Iodine.
• This property can be demonstrated by
displacement reactions.
• a halogen higher up the will displace a lower
• For example, when chlorine is added to Br-(aq),
the following reaction takes place.
• 2Br-(aq) + Cl2 ® 2Cl-(aq) + Br2
Predict the properties of other elements in
Group VII, given data where appropriate
• Reactivity here decreases down the group
• Remember the opposite is the case in group 1
Displacement reactions
• When chlorine (as a gas or dissolved in water) is
added to sodium bromide solution the chlorine takes
the place of the bromine.
• Because chlorine is more reactive than bromine, it
displaces bromine from sodium bromide.
• The solution turns brown.
• This brown colour is the displaced bromine.
• The chlorine has gone to form sodium chloride.
• If you look at the equation, you can see that the Cl
and Br have swapped places.
• chlorine + sodium bromide → sodium chloride +
• Cl2(aq) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
Videos of elements
• Here are some videos of elements in the
periodic table
• Have a look at some of them
• Form group 1 – Li, Na, K and Fr
• From groups 2 – Mg Ca
• From group 7 – F, Cl Br and I
Transition elements
Describe the transition elements:•
as a collection of metals having
high densities,
high melting points
forming coloured compounds,
and which, as elements and compounds, often
act as catalysts
The transition metals
• The elements in the centre of the periodic
table, between groups 2 and 3, are called the
transition metals.
• Most of the commonly used metals are there,
including iron, copper, silver and gold
Common properties
• The transition metals have the following properties in
• they form coloured compounds
• they are good conductors of heat and electricity
• they can be hammered or bent into shape easily
• they are less reactive than alkali metals such as
• they have high melting points - but mercury is a liquid
at room temperature
• they are usually hard and tough
• they have high densities
Noble gases
Group 0 – noble gases
• Argon makes up about 0.9 per cent of the air.
• It is one of a group of elements called the
noble gases.
• The noble gases are in Group 0 of the periodic
• The noble gases are all chemically unreactive
Describe the noble gases as being
• The noble gases are all chemically unreactive
• This is because they have a full outer shell
(i.e. 8 electrons)
Noble gases
Describe the uses of the noble gases
• in providing an inert atmosphere,
• i.e. argon in lamps,
• helium for filling balloons.
The main uses of the noble gases
• Used in balloons
and airships.
• It is much less
dense than air, so
balloons filled with
it float upwards.
• Used in advertising
signs, it glows when
electricity is passed
through it.
• Different coloured
neon lights can be
made by coating the
inside of the glass
tubing of the lights
with other
• Used in light bulbs.
• The very thin metal
filament inside the bulb
would react with oxygen
and burn away if the bulb
were filled with air instead
of argon.
• Argon stops the filament
burning away because it is
• Used in lasers.
Krypton lasers are
used by surgeons
to treat certain eye
• and to remove
Properties of metals
Distinguish between metals and non-metals by
their general physical and chemical properties.
• See next slides
State at room
Solid (except mercury,
which is a liquid)
About half are solids,
about half are gases, and
one (bromine) is a liquid
High (they feel heavy for
their size)
Low (they feel light for
their size)
Malleable or brittle
Malleable (they bend
without breaking)
Brittle (they break or
shatter when hammered)
Conduction of heat
Poor (they are insulators)
Conduction of electricity
Poor (they are insulators,
apart from graphite)
Magnetic material
Only iron, cobalt and nickel None
Sound when hit
They make a ringing sound
They make a dull sound
(they are sonorous)
solids at room temperature (except Hg)
metallic lustre - shiny
malleable and ductile
good conductors of heat and electricity
Metals can be easily oxidized or corroded
Metals can react with acid in a single displacement
reaction to make hydrogen gas and an aqueous
solution of a salt
• Some metals, like sodium or calcium will react with
water to make a base
• gases or solids at room temperature (except
• variety of colour and appearance
• brittle solids
• insulators (poor conductors)
Chemical properties
•Usually have 1-3 electrons in their outer
•Lose their valence electrons easily.
•Form oxides that are basic.
•Are good reducing agents.
•Have lower electronegativities.
•Usually have 4-8 electrons in their outer
•Gain or share valence electrons easily.
•Form oxides that are acidic.
•Are good oxidizing agents.
•Have higher electronegativities.
Physical Properties
•Good electrical conductors and heat
•Malleable - can be beaten into thin
•Ductile - can be stretched into wire.
•Possess metallic luster.
•Opaque as thin sheet.
•Solid at room temperature (except Hg).
•Poor conductors of heat and electricity.
•Brittle - if a solid.
•Do not possess metallic luster.
•Transparent as a thin sheet.
•Solids, liquids or gases at room
Reaction with oxygen
• Remember that metals react with oxygen in
the air to produce metal oxides, like
magnesium oxide.
• Non-metals react with oxygen in the air to
produce non-metal oxides.
Explain why metals are often used in
the form of alloys.
• An alloy is a mixture of two elements, one of
which is a metal.
• Alloys often have properties that are different
to the metals they contain.
• This makes them more useful than the pure
metals alone.
• For example, alloys are often harder than the
metal they contain.
Identify and interpret diagrams that
represent the structure of an alloy.
• Alloys contain atoms of different sizes, which
distorts the regular arrangements of atoms.
• This makes it more difficult for the layers to
slide over each other, so alloys are harder than
the pure metal.
Reactivity series
Place metals in order of reactivity:
 potassium,
 sodium,
 calcium,
 magnesium,
 zinc,
 iron,
 hydrogen
 copper,
• by reference to the reactions, if any, of the
elements with
water or steam,
dilute hydrochloric acid (except for alkali
Putting metals in order of reactivity
• The reactivity series for some common metals
• Other metals may be more reactive than
magnesium, or in between magnesium and
• If we put the metals in order of their reactivity,
from most reactive down to least reactive, we
get a list called the reactivity series.
• Here's a mnemonic to help learn the reactivity
• "Pond slime can make a zoo interesting - the
long crinkly sort goes purple."
Compare the reactivity series to the tendency
of a metal to form its positive ion,
• illustrated by its reaction, if any, with:
• the aqueous ions of other listed metals,
• the oxides of the other listed metals.
Displacement reactions
• A metal will displace (take the place of) a less
reactive metal in a metal salt solution.
For example,
iron + copper(II) sulfate iron sulfate + copper.
Fe(s) + CuSO4(aq)
FeSO4(aq) + Cu(s)
• Copper(II) sulphate is blue, iron sulphate is
• During the reaction the blue solution loses its colour
• and the iron metal is seen to turn pink-brown
• as the displaced copper becomes deposited on it.
A compound will always displace a less reactive metal
from solutions of its compounds.
• Another example of 'competition' between metals
to form compounds is observed in the reaction
between metals and metal oxides.
• When a metal is heated with the oxide of a less
reactive metal, it will displace the metal from it.
• Iron displaces the copper from the oxide - in
fact iron is behaving as a reducing agent, since
it is removing oxygen from the other metal.
• Observe this at the website below
Deduce an order of reactivity from a
given set of experimental results.
• You can use these displacement reactions to do this.
• Remember:
• The more reactive a metal is the more likely it is to form
a compound.
• The more reactive a metal, the more stable its
• More reactive metals displace less reactive metals
• Copper, silver and gold appear as elements in the earth due to their
unreactivity with their environment. They are easy to extract.
• Reactive metals are more difficult to extract. They are often found as
compounds or ores.

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