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Periodic classification and periodicity: Class 12 Chemistry
The modern periodic table and its blocks, and the trends of atomic radius, ionisation energy, electron affinity, electronegativity and metallic character across periods and down groups.
Key points
In the modern periodic table the elements are arranged in order of increasing atomic number, which is the Modern Periodic Law: the properties of elements are a periodic function of their atomic numbers. There are 7 periods and 18 groups. Elements in a group have the same outer electron configuration, so they show similar chemical behaviour.
The table is divided into blocks by the subshell that is being filled:
- s block: Groups 1 and 2, with helium.
- p block: Groups 13 to 18, apart from helium.
- d block: the transition elements, Groups 3 to 12.
- f block: the lanthanides and actinides, placed below the main table.
Atomic radius is half the distance between the nuclei of two bonded atoms of the same element. Across a period from left to right the radius decreases, because the nuclear charge increases while the electrons are added to the same shell, so the pull on the outer electrons is stronger. Down a group the radius increases, because each element has an extra shell, and the shielding effect of the inner electrons outweighs the rise in nuclear charge.
The ionisation energy is the energy needed to remove the most loosely held electron from a gaseous atom. It generally rises across a period and falls down a group. There are exceptions across a period: the first ionisation energy of beryllium is higher than that of boron, and that of nitrogen is higher than that of oxygen. This is because of the extra stability of a full s subshell and of a half filled p subshell.
Successive ionisation energies of an element always rise, because each electron is removed from a more positive ion. A very large jump between two successive values shows that the next electron comes from an inner shell, which tells you the group of the element.
The electron affinity is the energy change when a gaseous atom gains an electron. It is most negative (most energy released) for the halogens. The electronegativity is the tendency of an atom in a bond to attract the shared electrons. Fluorine is the most electronegative element. Electronegativity rises across a period and falls down a group.
Metallic character falls across a period and rises down a group. Metals lose electrons to form positive ions, while non-metals gain electrons or share them. The period 3 oxides change from basic (Na₂O, MgO) through amphoteric (Al₂O₃) to acidic (SiO₂, P₄O₁₀, SO₃).
Shielding is the reduction in the pull of the nucleus on the outer electrons caused by the electrons in inner shells. The effective nuclear charge is the nuclear charge less the shielding. It rises across a period and stays about the same down a group, which explains most trends.
The size of ions differs from that of their atoms. A cation is smaller than its atom, because it has lost its outer shell or electrons and has a greater effective nuclear charge. An anion is larger, because the added electrons repel each other. In an isoelectronic series, such as O²⁻, F⁻, Na⁺ and Mg²⁺, the ion with the higher nuclear charge is the smallest.
Study tip: For any trend question, answer in three steps: name the trend, say what changes (nuclear charge, shells or shielding), and link it to the strength of the pull on the outer electron.
Definitions and flashcards
Say the meaning first, then open the card to check.
Periodic law
The properties of elements are a periodic function of their atomic numbers.
Atomic radius
Half the distance between the nuclei of two bonded atoms of the same element.
Ionisation energy
The energy needed to remove the outermost electron from a gaseous atom.
Electron affinity
The energy change when a gaseous atom gains one electron.
Electronegativity
The tendency of an atom to attract the shared pair of electrons in a bond.
Shielding effect
The reduction in the nuclear pull on outer electrons by inner electrons.
Effective nuclear charge
The nuclear charge felt by an outer electron after allowing for shielding.
Isoelectronic species
Atoms and ions that have the same number of electrons.
Short questions with model answers
These are the kind of questions asked for 2, 3 and 5 marks. Try each one before opening the answer.
Why does atomic size decrease across a period? (3 marks)
Across a period the nuclear charge increases by one for each element, while the added electrons enter the same shell. The shielding stays almost the same, so the effective nuclear charge rises and pulls the outer electrons closer.
Why does the first ionisation energy decrease down a group? (3 marks)
Down a group the outer electron is in a shell farther from the nucleus and is more shielded by inner shells, so it is held less strongly and needs less energy to remove.
Why is the first ionisation energy of oxygen lower than that of nitrogen? (3 marks)
Nitrogen has a half filled 2p subshell, which is especially stable. In oxygen two electrons share one 2p orbital and repel each other, so one of them is removed more easily.
Which is the most electronegative element and why? (2 marks)
Fluorine. It is small and has a high effective nuclear charge, so it attracts bonding electrons strongly.
Arrange O²⁻, F⁻, Na⁺ and Mg²⁺ in order of increasing radius and explain. (3 marks)
Mg²⁺ < Na⁺ < F⁻ < O²⁻. All have 10 electrons, so the ion with the greater nuclear charge holds them more tightly and is smaller.
Common mistakes
- Saying that atomic size increases across a period. It decreases.
- Forgetting the exceptions in the ionisation energy trend at Be and B, and at N and O.
- Comparing the sizes of ions without checking whether they have the same number of electrons.
- Giving the reason for the trend down a group as increased nuclear charge. The reason is the extra shell and shielding.
- Mixing up electron affinity, which is an energy change, with electronegativity, which is a tendency in a bond.