Trends in the Periodic Table (VCE SSCE Chemistry): Revision Notes

Trends in the Periodic Table

The periodic table is more than just a list of elements organised by their electronic configuration. It's a powerful tool that helps us understand and predict the properties of elements. By examining the position of an element in the periodic table, you can determine key characteristics such as atomic radius, electronegativity, and first ionisation energy.

Understanding the periodic table structure

The periodic table organises elements in a way that reveals patterns in their properties. The group number tells you how many valence electrons an atom of that element has, while the period number indicates how many electron shells are occupied in an atom.

Trends in electronic configuration

To understand why element properties follow patterns, let's examine the electronic configurations of the alkali metals in group 1. These elements (lithium, sodium, potassium, rubidium and caesium) are all relatively soft metals that react vigorously with water and oxygen.

Their electronic configurations are:

• Li: $1s^2 2s^1$
• Na: $1s^2 2s^2 2p^6 3s^1$
• K: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1$
• Rb: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6 5s^1$
• Cs: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6 4d^{10} 5s^2 5p^6 6s^1$

As you move down a group, the number of electron shells increases. This means valence electrons occupy higher energy subshells and experience a weaker attraction to the nucleus. This changing attractive force between valence electrons and the nucleus as you move down a group creates observable trends in element properties.

Effective nuclear charge

Effective nuclear charge (sometimes called core charge) measures the attractive force that valence shell electrons experience towards the nucleus. This concept is crucial for predicting element properties and explaining trends across periods.

Consider a lithium atom with an atomic number of 3. It has three protons in its nucleus, two electrons in the first shell, and one electron in the second shell. The valence electron is attracted to the three positive charges in the nucleus, but it's also repelled by the two electrons in the inner shell. These inner shell electrons shield the valence electron from the full nuclear attraction. The valence electron effectively experiences the nucleus as if there were only a +1 charge. Therefore, lithium has an effective nuclear charge of +1.

Worked example: calculating effective nuclear charge for aluminium

Effective nuclear charge by group

The effective nuclear charge for main group elements corresponds to their group number:

Group	Effective nuclear charge
1	+1
2	+2
13	+3
14	+4
15	+5
16	+6
17	+7
18*	+8

*Note: Helium has an effective nuclear charge of +2.

How effective nuclear charge changes in the periodic table

The attraction between the nucleus and valence electrons changes predictably across the periodic table:

Direction	Trend in effective nuclear charge	Trend in attraction between nucleus and valence electrons
Down a group	Remains constant	Effective nuclear charge stays constant down a group, but valence electrons are held less strongly because they are further from the nucleus due to additional shells in the atom.
Left to right across a period	Increases	The valence electrons are more attracted to the nucleus as the effective nuclear charge increases.

Electronegativity

Electronegativity is the ability of an atom to attract electrons towards itself when forming a chemical bond. It measures how strongly an atom pulls on the electrons of nearby atoms. Since the positive pull comes from the nucleus, atoms with greater effective nuclear charge have higher electronegativity.

The diagram below shows electronegativity values for many main group elements. Note that noble gases (group 18) are not listed because they have a stable outer shell and don't readily form bonds with other atoms.

Electronegativity trends explained

Direction	Trend in electronegativity	Explanation
Down a group	Decreases	The effective nuclear charge stays constant and the number of shells increases down a group. Therefore, valence electrons are less strongly attracted to the nucleus as they are further from it. As a result, electronegativity decreases.
Left to right across a period	Increases	The number of occupied shells in the atoms remains constant but the effective nuclear charge increases across a period. Therefore, valence electrons become more strongly attracted to the nucleus. As a result, electronegativity increases.

Atomic radius

Atomic radius is a measurement of the size of atoms. It represents the distance from the nucleus to the valence shell electrons. Scientists usually measure atomic radius by halving the distance between the nuclei of two atoms of the same element that are bonded together.

How atomic radius is measured

Atoms don't have sharply defined boundaries, so we cannot measure their radii directly. Instead, we measure the distance between nuclei of atoms of the same element in molecules. For example, in a hydrogen molecule ($\text{H}_2$), the two nuclei are 74 picometres (pm) apart. The radius of each hydrogen atom is assumed to be half that distance, i.e. 37 pm.

As you move across a period, atomic radius decreases. This occurs because as effective nuclear charge increases, valence shell electrons are pulled in more tightly towards the nucleus.

Atomic radius trends explained

Direction	Trend in atomic radius	Explanation
Down a group	Increases	Effective nuclear charge stays constant and the number of shells increases as you move down a group. As a result, atomic radius increases.
Left to right across a period	Decreases	As you move across a period, the number of occupied shells in the atoms remains constant but the effective nuclear charge increases. The valence electrons become more strongly attracted to the nucleus, so atomic radius decreases across a period.

First ionisation energy

When an element is heated, its electrons can move to higher energy shells. If an atom receives sufficient energy, an electron can be completely removed from the atom. When this happens, the atom has one fewer electron than protons in the nucleus and becomes a positively charged ion.

First ionisation energy is the energy required to remove one electron from an atom of an element in the gas phase. For example, the first ionisation energy of sodium is 494 kJ per mole of sodium atoms. (A mole is a way of counting atoms—you'll learn more about this in later chapters.)

First ionisation energy trends explained

Direction	Trend in first ionisation energy	Explanation
Down a group	Decreases	Effective nuclear charge stays constant and the number of shells increases down a group. Therefore, valence electrons are less attracted to the nucleus as they are further from it. As a result, the energy required to overcome the attraction between the nucleus and the valence electron is less, and first ionisation energy decreases down a group.
Left to right across a period	Increases	Effective nuclear charge increases and the number of occupied shells remains constant across a period. As a result, valence electrons become more strongly attracted to the nucleus, and more energy is required to remove an electron. Therefore, first ionisation energy increases across a period.

The s- and p-block elements follow these patterns consistently. As effective nuclear charge increases across a period, so does ionisation energy. As atomic radius increases down a group (due to additional electron shells), electrons are further from the nucleus, so valence electrons in elements lower in a group can be removed more easily, meaning ionisation energy decreases.

Trends in metals and non-metals

One of the most useful distinctions in the periodic table is the classification of elements as metals, metalloids, or non-metals. Metalloids share properties of both metals and non-metals.

Understanding whether an element is a metal or non-metal helps you predict how it will bond with other elements. The reactivity of both metals and non-metals also follows predictable trends in the periodic table.

Metal reactivity trends

The properties of metals arise from their tendency to readily lose electrons. As you move down a group, valence shell electrons are further from the pull of the nucleus and are therefore lost more easily. This means that as you go down a group, metals become more reactive.

Across a period, the opposite occurs—as effective nuclear charge increases, metals become less reactive. This means the most reactive metals are in group 1.

Non-metal reactivity trends

In contrast, the properties of non-metals arise from their tendency to readily gain electrons. In general, non-metals become less reactive as you move down a group, because valence electrons are further from the nucleus. Non-metals become more reactive across a period, as effective nuclear charge increases. The noble gases are an exception to this rule, as they are inert (unreactive).

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