Chemistry AS Unit 1


  • Created by: sanam
  • Created on: 15-10-12 18:57

The S, P, D, F Blocks

4.1 The periodic table.

The s-, p,- d- and f- blocks of the periodic table:


  • All elements with their highest energy level in s-orbitals are in the s-block.
  • All elements with their highest energy level in p-orbitals are in the p-block.
  • All elements with their highest energy level in d-orbitals are in the d-block.
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  • A vertical column of elements.
  • Form a chemical family.
  • All have similar properties.
  • Same number of electrons in their outer main level.


  • In the s-block, elements get more reactive down the group.
  • To the right of the periodic table (non-metals), elements get more reactive up a group.
  • Transition metals are rather unreactive. Most useful metals are found here.
  • Lanthanides are metals.
  • Actinides are radioactive metals.
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  • Horizontal rows.
  • Trends and physical properties across a period.

Placing of hydrogen and helium:

  • Helium has noble gas properties so it is sometimes placed over the noble gases in group 0. However it is not a p-block element, so it is sometimes placed on its own.
  • Hydrogen is sometimes placed above group 1, however it is usually placed on its own.
  • Hydrogen is not similar to group 1 as it is a gas and they are all reactive metals.



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Trends in Period 3

4.2 Trends in the properties of elements of Period 3.

Periodicity and properties of elements in Period 3:

  • Periodicity is explained by electron arrangements of the elements.
  • Sodium, magnesium and aluminium are metals. They have giant structures. They lose their outer electrons to form ionic compounds.
  • Silicon has four outer electrons. It forms four covalent bonds. Has some metallic properties and is classed as a semi-metal.
  • Phosphorous, sulphur and chlorine are non-metals. Either accept electrons to form ionic compounds or covalently bond.
  • Argon has a full outer shell so it is unreactive.
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Trends in Melting/ Boiling Points

Trends in Melting and Boiling points:

  • Left = high melting points (sodium is the exception) / Right = low melting points.
  • These are due to their structure.
  • The m.p.t and the b.p.t increase from Na to Al because the metal-metal bonds get stronger. This is because metal ions have an increasing number of delocalised electrons and a decreasing atomic radius. This leads to a higher charge density, which attracts the ions together more strongly.
  • Silicon is macro molecular. It has a tetrahedral structure. Strong covalent bonds link all its atoms together. A lot of energy is needed to break these bonds. Silicon has high m.t.p and b.p.t
  • Phosphorous, sulphur and chlorine are all molecular substances. Their m.t.p and b.t.p depend upon the strength of the van der Waals forces between their molecules. Van der Waals are weak and easily overcome. The m.t.p and b.t.p of these elements are low.
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Trends in Melting/ Boiling Points (2)


  • Van der Waals becomes stronger with an increase of atoms in a molecule. Sulphur is the biggest out of the three non-metals and so it has higher melting and boiling points than P and Cl.
  • Argon has low m.t.p and b.t.p because it exists as single atoms. TF. very weak van der Waals forces.


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Atomic Radii

4.3 More trends in the properties of the elements of period 3.

Atomic radii:

  • This is the size of atoms
  • Depends on the type of bonding.
  • Decreases across each period and there is a jump between periods.
  • Atoms get larger down groups.

Why the radii of atoms decrease across a period:

  • From Na to Cl, protons are being added to the nucleus and electrons are being added to the outer main level.
  • The charge on the nucleus increases (atomic number)
  • The increased charge pulls the electrons closer to the nucleus.
  • There isn’t an increase in shielding as there are so additional electron shells to provide it.
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Atomic Radii (2)


Why the radii of atoms increase down a group:

  • Down a group the atoms of each element have one extra complete level than the one before it.
  • The outer main level becomes further away from the nucleus, so the atomic radii increases.
  • There are more inner levels of electrons to provide shielding of the nucleus so the the nuclear charge has less pull on the outer electrons.
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First Ionisation Energy

First Ionisation energy:

  • This is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state.
  • These have periodic patterns.
  • Generally increases across a period.
  • Decreases down a group.

Why first ionisation energy increases across a period:

  • From left to right the number of protons increases but the electrons enter the same main level.
  • There is an increased charge on the nucleus; the nuclear charge has a greater pull on the electrons in the outer main level, so it becomes more difficult to remove one.
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First Ionisation Energy (2)

Why the first ionisation energy decreases going down a group:

  • The number of filled inner levels increases down the group.
  • There is an increase in electron shielding of the nucleus.
  • The outer electrons are held less strongly.
  • Outer electrons are easier to remove.

Why is there a drop in ionisation energy from one period to the next?

  • At the start of each period there is a new main level.
  • This causes an increase in atomic radii due to electron shielding of the nuclear charge.
  • The outer electrons are further from the nucleus and less strongly attracted to it, so it easier to remove one of them.
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Drop in IE Group 2 and 3

4.4 A closer look at ionisation energies.

The drop in ionisation energy between Groups 2 and 3:

  • This is to with the sub-level from which the first electron is removed.
  • Magnesium loses a 3s and aluminium loses a 3p electron.
  • The p-shell electron is already in a higher energy level than the s-level electron.
  • So less energy is needed to remove it.
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Drop in IE Group 5 and 6

The drop in first ionisation energy between groups 5 and 6:

  • This is about pairing of electrons.
  • Phosphorous has no paired electrons as each electron is in a different p-orbital.
  • Sulphur has two of its p-electrons paired. One of these is easier to remove as there is repulsion between the pair.
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Successive Ionisation Energies

Successive Ionisation Energies:

  • Is electrons are removed one at a time, it gets harder each time to remove and electron.
  • When the main energy level changes there is a sharp increase in ionisation energy.
  • Throughout a main level there is a gradual increase.
  • The closer the main level is to the nucleus the harder it becomes to remove and electron.
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