2012
Definition: periodicity is the repeating pattern of properties of elements across different periods in the Periodic Table.
The modern periodic table encloses the elements in boxes and displays the boxes in groups (vertical columns) and periods (horizontal rows).
· There are 8 main groups.
· There are four periods of transition metal elements which slot into the
table between main group 2 and 3.
· The elements are arranged in order of atomic number.
Example
Please refer to the OCR text book pages 78-81 for a history of the Periodic Table - Development of the Periodic Table from Döbereiner, Newlands, Mendeleev, Moseley, Seaborg, et al. and try the questions to test your understanding.
The elements in period 2 resemble closely the elements in period 3, in terms of their repeating chemical behaviour and physical properties. For example the first element in period 2, lithium has similar chemistry to the first element in period 3, sodium. These elements are both metals, which float on water.
Electronic configurations
All the elements in the same period have a similar electron core.
Student activity 1
(a) Complete the electronic configurations of the period 3 elements sodium to chlorine (by adding the superscripts).
Na 1s 2s 2p 3s
Mg 1s 2s 2p 3s
Al 1s 2s 2p 3s 3p
Si 1s 2s 2p 3s 3p
P 1s 2s 2p 3s 3p
S 1s 2s 2p 3s 3p
Cl 1s 2s 2p 3s 3p
(b) Write the electronic configuration for the elements of group IV in
the space below.
Period Element Atomic Number Configuration
2 C 6 …………………………….
3 Si 14 …………………………….
4 Ge 32 …………………………….
Elements in a group resemble each other in their chemical reactivity and physical properties. The electronic configuration of the elements in group four all end with s2, p2. This means that the early elements in group 4 tend to get involved in covalent bonding and the later elements are metals and tend to lose 4 electrons. Oxidation states are sometimes +4 but not always.
All groups are vertical columns of elements with a characteristic electronic configuration for the higher energy sub-shells.
Group 1 elements are known as ……………………………
Group 2 elements are known as …………………………….
Group 7 elements are known as …………………………….
Group 8 elements are known as …………………………….
Ionisation energy
Student activity 2
1. Define the term first ionisation energy
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2. State the three factors that affect ionisation energy
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3. Write an equation to represent the first ionisation energy of oxygen
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Student activity 3
Complete the explanation of comparative ionisation energy for each of the pairs below, by crossing out the incorrect words.
a) Silicon and phosphorus
Silicon / phosphorus has the highest ionisation energy.
Because:
Silicon / phosphorus has more protons in its nucleus.
Silicon / phosphorus has the smallest atomic radius.
They both have the same number of shells/electrons
Silicon / phosphorus has the strongest attraction between the nucleus and the outer electrons.
b) Potassium and rubidium
Potassium/ rubidium has the highest ionisation energy.
Because:
Rubidium has fewer/more protons in its nucleus.
Potassium/ rubidium has the largest atomic radius.
Potassium/ rubidium has the most shells.
Potassium/ rubidium has the most shielding.
Potassium/ rubidium has the strongest attraction between the nucleus and the outer electrons.
c) Sodium and Magnesium
Sodium/magnesium has the highest ionisation energy.
Because:
Sodium/magnesium has more protons in its nucleus.
Sodium/magnesium has the smallest atomic radius.
They both have the same number of shells/electrons
Sodium/magnesium has the strongest attraction between the nucleus and the outer electrons.
Student activity 4
State and explain the general trend in ionisation energy. Complete the statements.
i) Across the period
Ionisation energy ………………………….
Because: atomic radii ……………………, nuclear charge……………….., shielding………………………………………….,
and nuclear attraction…………………………………
Therefore, it is ………………………… to remove an electron as atomic number increases.
ii) Down a group
Ionisation energy ………………………….
Because: atomic radii ……………………, nuclear charge……………….., shielding………………………………………….,
and nuclear attraction…………………………………
Therefore, it is ………………………… to remove an electron as atomic number increases.
Comparison of first and second ionisation energies
Explain why the first ionisation energy is always less than the second ionisation energy for all elements.
The proton number remains the same but there is one less electron .
The proton/electron ratio is greater; therefore nuclear attraction for the remaining electrons is greater. It follows; more energy is needed to remove a second electron. Also, this can be described as a reduction of electron-electron repulsion.
Atomic radii
Below you are given values for the size of the atoms in period three and in group 2. Study the data carefully. If the actual values are required in the exam they will be given on the paper.
Period 3
Atom / Na / Mg / Al / Si / P / S / ClRadius/nm / 0.186 / 0.160 / 0.143 / 0.117 / 0.110 / 0.104 / 0.099
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This general trend is the same for all periods in the periodic table.
Group 2
Atom / Radius/nmBe / 0.112
Mg / 0.160
Ca / 0.197
Sr / 0.197
Ba / 0.217
Ra / 0.220
This general trend is the same for all groups in the periodic table.
Student activity 4
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1. State what happens to the size of the atoms moving across period three?
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2. State what happens to the size of the atoms on descending group 2?
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3. Explain the following trends:
a) The change in size across a period
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b) The change in size down a group.
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Melting points and boiling points
Melting point and boiling point data for the elements in period three is shown below.
The melting point / boiling point of a particular element is dependant on the bonding present in the structure of the element. Metals have high melting points because they contain metallic bonding. Non-metals with giant covalent structures like silicon have high melting points because the covalent bonds are strong and difficult to break down. The simple covalent molecules like phosphorus and chlorine have low melting and boiling points because they are held together by intermolecular forces, which are weak, and easily broken down.
Atom / Na / Mg / Al / Si / P / S / Cl / ArMelting Point / OC / 98 / 649 / 660 / 1410 / 44 / 119 / -101 / -189
Boiling Point / OC / 890 / 1117 / 2447 / 2677 / 281 / 445 / -36 / -186
Student activity 5
1)The data above needs further explanation.
Atom / Na / Mg / Al / Si / P / S / Cl / ArBonding
Structure
Complete the bonding row in the table above using;
· M for metallic bonding
· C for covalent bonding
Complete the structure row using
· S for a simple molecular structure
· G for a giant structure
2) Why are the boiling points and melting points of the metals increasing from sodium to aluminium?
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3) Why does the group four element have the highest boiling and melting point?
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4) Why is there a sharp decrease in boiling point between group 4 and group 5?
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5) Explain the trend in melting and boiling points of phosphorus (P4), sulphur (S8) and chlorine (Cl2)
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Electrical Conductivity
Metallic elements are good conductors of electricity. This is because their structures contain delocalised electrons, which are free to move through the structure. Metallic character and therefore conductivity increases to a maximum in group three. Metallic character also increases down the group. Elements beyond silicon show no conduction as all electrons are involved in bonding and no free charge carriers are present. See notes on metallic bonding.
Describe the trend in the metallic character of the elements: -
Across a period:
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Down a group:
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Explain why aluminium is a better conductor of electricity than magnesium and why magnesium is better than sodium.
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