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Module 3 summary notes (A Level Chemistry OCR A)
This document contains summary notes for Module 3 Periodic table and energy, taken from the A Level Chemistry for OCR A OXFORD textbook.
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MODULE 3PERIODIC TABLE AND ENERGY
CHAPTER 7 – PERIODICITY
7.1 – THE PERIODIC TABLE
Periodic table was discovered by Mendeleev.
NOW – arranged in increasing atomic number, you gain 1 proton as you go along.
GROUP – each group has atoms with the same number of outer-shell electrons and similar
properties.
PERIOD – the period number indicates the number of electron shells in an atom.
PERIODICITY – across each period, there is a repeating trend in properties of the elements.
Topics in this chapter look at periodicity of several properties:
ELECTRON CONFIGURATION
IONISATION ENERGY
STRUCTURE
MELTING POINTS
ELECTRON CONFIGURATION
s, p, d, f
For each period the s and p sub-shells are filled in the same way – PERIODIC PATTERN
s block - Groups 1 and 2
d block - Transition group Each block corresponds to the electron’s highest energy sub-
p block - Groups 3,4,5,6,7 and 8 shell.
And f block
Group 1 – alkali metals.
Group 2 – alkaline earth metals.
Group 17 – halogens.
Group 18 – noble gases.
Groups 3-12 – transition metals.
7.2 - IONISATION ENERGY
IONISATION ENERGY - measures how easily an atom loses electrons to form positive ions.
The FIRST IONISATION ENERGY is the energy required to remove one electron from each
atom in one mole of GASEOUS atoms of an element to form one mole of gaseous 1 + ions.
E.g. The energy required to remove an electron from each atom of one mole of gaseous Na
to form one mole of gaseous ions is 496 KJ mol-1.
Na(g) Na+(g) + e- First ionisation energy = 496KJmol-1
+ 2+ -
Na (g) Na (g) + e Second ionisation energy
,The SECOND IONISATION ENERGY is the energy required to remove one electron from each
atom in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.
Third Na2+(g) Na3+(g) + e-
Fourth Na3+(g) Na4+(g) + e-
….
KEY POINT
An element has many ionisation energies as it does electrons.
FACTORS AFFECTING IONISATION ENERGIES
Electrons are held in their shells due to the attraction from the nucleus.
The first electron will experience the least attraction from the nucleus.
ATOMIC RADIUS easier to lose as distance increases IE
The bigger the atom, the less the nuclear attraction.
This is due to the greater distance between the nucleus and the outer electrons.
NUCLEAR CHARGE as number of protons increase, stronger attraction to outer shell. IE
The more protons in the nucleus, the greater the attraction between the nucleus and
the outer electrons.
SHIELDING repulsion theory. More shields, less attraction IE
Inner shell electrons repel outer shell electrons.
This repulsion is called the shielding affect.
It reduces the attraction between the nucleus and the outer electron.
EVIDENCE FOR ELECTRON SHELLS
Graphs can show number of electrons in each shell.
So, you can identify the element, elements group and its outer electrons.
7.3 - PERIODIC TRENDS IN BONDING AND STRUCTURE
BOILING POINTS ACROSS PERIODS 2 AND 3
1 – 4 groups HIGH boiling points (sharp increase)
5 – 0 groups LOW boiling points
4 – 5 groups (sharp decrease in boiling point)
This is because:
The forces between particles across a period.
Where there are strong forces, the melting and boiling points are high (metallic
bonding/giant covalent structure).
Where there are weak forces, the melting and boiling points will be low. (simple
molecular structures).
METALLIC BONDING
Bonding for metals.
It is the strong electrostatic attraction between cations (positive ions) and
delocalised electrons.
, The cations are fixed in position, maintaining the structure and shape of the metal.
The delocalised electrons are mobile and are able to move throughout the structure.
Only electrons move.
STRUCTURE, BONDING:
Sea of positive ions with delocalised electrons holding them together.
In a solid metal structure, each atom has donated its negative outer-shell electrons
to a shared pool of electrons, which are delocalised throughout the whole structure.
The positive ions (cations) left behind consist of the nucleus and the inner electron
shells of the metal atoms.
PROPERTIES:
CONDUCT ELECTRICITY metals conduct electricity in solid or liquid states.
When a voltage is applied across a metal, the delocalised electrons can move
through the structure carrying charge.
HIGH MELTING AND BOILING POINT Melting point depends on the strength of the
metallic bonds holding the atoms in the giant metallic lattice.
For most metals, high temperatures are necessary to provide the large amount if
energy needed to overcome the strong electrostatic attraction between the cations
and electrons.
This strong attraction results in most metals having high melting and boiling points.
SOLUBILITY Metals do not dissolve. Any interaction will lead to a reaction.
GIANT COVALENT BOND
STRUCTURE, BONDING:
Many atoms are held together by a network of strong covalent bonds to form a giant
covalent lattice.
Lattice structure
Tetrahedral structure (could be)
Strong covalent bonds
PROPERTIES:
HIGH MELTING AND BOILING POINTS covalent bonds are strong, high temps are
necessary to provide a large quantity of energy needed to break the strong bonds.
INSOLUBLE covalent bonds holding together the atoms in the lattice are too strong to be
broken by interactions with solvents.
DON’T CONDUCT ELECTRICITY, however graphene and graphite might as they are forms of
carbon.
In carbon (diamond) and silicon, all four outer-shell electrons are involved in
covalent bonding, so none are available for conducting electricity.
Carbon is special in forming several structures in which one of the electrons is
available for conductivity. Graphene and graphite are able to conduct electricity.
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