GRADE 10 SCIENCE
THE PERIODIC TABLE
The Periodic Table is to arranged in a way that enables chemists to understand patterns in
the properties of elements.
The elements are laid out in order of Atomic (proton) Number.
Many of the similarities and differences in the properties of elements can be explained by
the electronic structure of the atoms - electron configuration
The main structural features of the periodic table are:
Columns of similar elements called Groups.
They are usually similar chemically and physically BUT there are often important
trends in physical properties and chemical reactivity up/down a group.
Complete horizontal rows are called Periods.
Usually consist of a range of elements of different character. There are important
trends from left to right across a period e.g. the most important overall change is from
metallic  non-metallic element character.
The metal/non-metal divide – Know it!
Elements on the left are METALS
Elements on the right are NON-METALS
Some elements along the divide are referred to as SEMI-METALS or METALLOIDS.
They have mixed metal/non-metal character.
PERIODICITY
Many physical and chemical properties of the elements follow a sequence and change
systematically as one moves across from one side of a period to the other. This is called
periodicity. We study a few of these properties through periods 2 and 3
Density
3,5
TRENDS:
- General increase in density
from Group 1 to the middle of
the periodic table
3
Density (g.cm-3)
2,5
- General decrease from the
middle of the table to Group 8
2
1,5
- Density in a group in Period 3
elements is higher than in
Period 2
1
0,5
Element at top of group
0
0
1
2
3
4
5
6
7
9 Elements
8
Group Number
at bottom of group
Melting Points
4000
TRENDS:
3500
- Melting points of metals
increase across a period
Melting Point (°C)
3000
2500
- Melting points of non-metals
decrease across a period
2000
- Melting points of metals
decrease down groups
1500
1000
- Melting points of non-metals
increase down groups
500
0
-500
0
1
2
3
4
5
Group Number
6
7
8
9
Boiling Points
TRENDS:
6000
- Boiling points of metals
increase across a period
Boiling Point (°C)
5000
- Boiling points of non-metals
decrease across a period
4000
3000
- Boiling points of metals
decrease down groups
2000
- Boiling points of non-metals
increase down groups
1000
0
-1000
0
1
2
3
4
5
6
7
8
ie SAME as MELTING POINT
9
Group Number
Atomic Radius
Definition: Distance between the nucleus and the outermost orbital of an atom


True atomic radius
Group 8 exist as free atoms
Covalent radius
Diatomic molecules
½d= radius
Metallic radius
Atoms in metallic lattice
TRENDS:
- increase down a group
(extra shell added in each
period)
- decrease across a period
(adding electrons to the
same shell but, increased
negative charge means
increased electrostatic
attraction and thus
‘contracts’ the atom
Formulae of Halides
The halogens (Group 7 – F2, Cl2, Br2, I2) are reactive and form halides (fluoride, chloride etc)
when reacting with other elements (except with unreactive Group 8). Chlorides are used as
the example below.
Group
1
2
3
4
5
6
7
Period 2 Halide
LiCl
BeCl2
BCl3
CCl4
NCl3
OCl2
FCl
Period 3 Halide
NaCl
MgCl2
AlCl3
SiCl4
PCl3
SCl2
Cl2
Ratio element:Cl
1
2
3
4
3
2
1
NOTE: Peak of combining power occurs at changeover from metals to non-metals
Formulae of Oxides
The same trend as for the halides is noted – this is true of many of the compound form –
combining power increases across the metals and decreases across the non-metals
Group
1
2
3
4
5
6
7
Period 2 Oxide
Li2O
BeO
B2O3
CO2
varies
O2
F 2O
Period 3 Oxide
Na2O
MgO
Al2O3
SiO2
P4O6
varies
Cl2O
Ratio element:O
0.5
1
1.5
2
1.5
-
0.5
Ionisation Energy
Definition: The ionization energy of an element is the amount of energy required to remove
an electron from a gaseous atom to form a cation.
The higher the ionisation energy, the more difficult it is to remove an electron from an atom
It is possible to remove successive electrons from an atom:
First eFirst ionisation energy
M  M+ + eSecond eSecond ionisation energy
M+  M2+ + eThird eThird ionisation energy
M2+  M3+ + eWe only look at the trends in first ionisation energy:
TRENDS:
- ionisation energies increase
across each period – as an
extra proton and electron are
added to an atom, electrostatic
attraction increases.
- large drop at the start of each
new period. Electrons are being
added to a new shell which is
further from the nucleus thus
experience weaker electrostatic
forces
Electron affinity
Definition: The energy change when an electron is added to a gaseous atom to form an anion.
(ie a measure of the ability of an atom to accept an electron)
TRENDS:
- electron affinity increases
across each period – those of
metals are generally lower than
those of non metals
- electron affinities of the
halogens are highest – by
accepting an electron, they
attain the noble gas
configuration of the noble gas
next to them
Electronegativity
Definition: The ability of an atom in a compound to attract the bonding electron pair(s)
TRENDS:
- increases across each period halogens are highest
– decreases down a group
CHEMICAL PROPERTIES among ELEMENTS
Reactivity determined by the no of valence electrons that need to be lost or gained to
achieve a full outer shell.
- lose electrons - no of protons > no of electrons
- positive ion (CATION) formed
- gain electrons - no of electrons > no of protons
- negative ion (ANION) formed
Elements in same group – same no of valence electrons, so they will lose/gain the same no of
electrons to achieve a full outer shell, so reactivity is similar.
GROUP 1 – alkali metals
- all have 1e in outer shell.
- lose 1 electron when they react – form cations with a 1+ charge
- react vigorously with water to give hydrogen gas and an alkaline solution of the metal
hydroxide:
eg.
2Li + 2H2O  2LiOH+ H2
- burn strongly in oxygen to form solid oxides:
eg.
4Li + O2  2Li2O
- ionisation energy is low in Group 1, so losing an electron to form +1 ion is easy, so elements in
Group 1 are very reactive.
- ionisation energy decreases as we go down the group (shell size increases), so reactivity
INCREASES down the group.
- physical properties – soft, low density (float on water), silvery colour, low melting point
GROUP 2 – alkali earth metals
- all have 2e in outer shell.
- lose 2 electrons when they react – form cations with a 2+ charge
- less reactive than Group 1 since more energy needed to remove second electron.
- reactivity with water varies. Be – not reactive
Mg – very slow with cold water, vigorous with steam
Mg + 2H2O(l)  Mg(OH)2+ H2
Mg + H2O(g)  MgO + H2
Ca, Sr, Ba react more vigorously with cold water
Ca + 2H2O  Mg(OH)2+ H2
Group 2 hydroxides are only very slightly soluble.
- burn in oxygen to form solid oxides:
eg.
2Mg + O2  2MgO
- ionisation energy decreases as we go down the group (shell size increases), so reactivity
INCREASES down the group.
- physical properties – silvery colour but harder, more dense, higher melting than group 1
GROUP 7 – halogens
- form diatomic molecules (F2, Cl2, Br2, I2) – poisonous!
- all have 7e in outer shell.
- gain 1 electron when they react – form anions with a 1- charge
- react with Group 1 metals to form salts.
2Li + Cl2  2LiCl
- Since an electron is being, reactivity depends on electron affinity.
- Electron affinity decreases down the group, so reactivity DECREASES down the group.
Fluorine is frequently explosive – we don’t use it.
- Physical properties: - low melting and boiling points (increase as you go down group)
At room temp, F2, Cl2 - gases, Br2 - liquid, I2 - solid,
- colour varies
F2 – colourless, Cl2 - green-yellow, Br2 - red-brown, I2 – purple-black
GROUP 8 – noble/inert gases
- have a full outer shell – no gain or loss of electrons, unreactive or INERT
- Physical properties – low boiling point, colourless, gases at room temperature