Periodic Trends

What are periodic trends?

The periodic table is a wonderful thing. The version that we use today was created in 1869 by the Russian chemist Dmitri Mendeleev, and built on the work of scientists such as John Newlands. Mendeleev ordered the known elements by atomic mass, but noticed that they showed certain properties which repeated every eight or so elements. He therefore also arranged the elements in rows and columns, so that elements with similar properties were above and below each other in the table. In some cases the properties didn't quite match up; here, Mendeleev left gaps in the table, proposing that they would be filled by then-undiscovered elements. He was proven right when these elements were later discovered.

Fig. 1 - The periodic table

All in all, the periodic table is a regular arrangement of elements by atomic number. Rows are called periods, and columns are called groups. The elements are so arranged to show periodic trends.

Another word you might come across is periodicity. Periodicity implies repetition after a certain period or gap. In the case of the elements, periodicity means the repetition of properties of elements after a certain gap of atomic number. Mendeleev noticed the periodicity that occurred every eight or so elements, and arranged them accordingly. Let's now take a look at some of the trends he saw and explain why they occur.

Periodic table trends

As we saw above, rows in the periodic table are known as periods. Columns in the periodic table are known as groups. All in all, there are seven periods and eighteen groups in the periodic table. This might seem like a lot to learn about, but as you'll see, they show regular periodic trends when it comes to certain properties. The properties we'll look at today include:

• Electron configuration

• Atomic radius
• Electronegativity
• First ionisation energy
• Melting and boiling points
• Density
• Electrical conductivity

First up - electron configuration.

Periodic trends in electron configuration

What do an element's period and group tell you about its electron configuration? Luckily for you, knowing where it lies on the periodic table tells you exactly how it is structured. The period of an element tells you how many electron shells it has - for example, elements in period 2 all have two electron shells. The group of an element tells you how many electrons it has in its outer shell. For example, all the elements in group 3 have three electrons in their outer shell. Let's see how this varies across a period and down a group.

Electron configuration across a period

The number of electrons increases by one each time you move across a period. Each element has exactly one more electron than the element immediately to the left of it, and exactly one less than the element to the right. However, all elements in the same period have the same number of electron shells.

Take period 2, for example. All of the elements within it have just two electron shells. The first element in the period, lithium, has three electrons, whilst the next, beryllium, has four. This is shown in the periodic table by atomic number. An element's atomic number is its number of protons, but in a neutral atom, this is equal to its number of electrons. You can see in the image below that atomic number increases as you move along the period.

Fig. 2 - Period 2, shown with atomic number in the top left

Electron configuration down a group

The number of electron shells increases as you go down a group. However, elements in the same group have exactly the same number of electrons in their outer shell - the only difference is the number of inner shells.

For example, all the elements in group 1 have just one electron in their outer shell. But whilst hydrogen has only one electron shell in total, the next member lithium has two, and sodium has three.

Fig. 3 - Group 1 elements, shown with their electron configurations in brackets

Periodic table radius trend

Atomic radius depends on factors such as nuclear charge and number of electron shells, which we'll now explore.

Atomic radius across a period

Atomic radius decreases as you move along a period in the periodic table. But why is this the case? To explore this fully, we need to look at the atomic structure of elements in a period.

In the modern periodic table, elements are arranged by their atomic number. Atomic number is the number of protons an element has in its nucleus, found in the centre of the atom. Each element has the same number of protons as it does electrons, so atomic number also tells us the number of electrons in an element. These electrons are found orbiting the nucleus in shells. Although elements in the same period have different numbers of electrons, they have the same number of electron shells. Take a look at the example below. Carbon has six electrons, whereas oxygen has eight. However, both have just two electron shells.

Fig. 4 - The electron configurations of carbon and oxygen

Electron shells are attracted to the nucleus thanks to a strong electrostatic attraction between the positively charged nucleus and the negatively charged shells. This determines atomic radius.

As you go across a period, the atomic number increases - each element has one more proton and one more electron than the element before it. It means that the nucleus' charge increases. The outermost electron shell experiences a stronger attraction to the positively charged nucleus, so the negative electrons are pulled in closer to the nucleus in the centre of the atom. This decreases atomic radius.

Fig. 5 - Atomic radius of carbon and oxygen

Atomic radius down a group

Atomic radius increases as you go down a group. You might be wondering why. As you move down a group in the periodic table, atomic number increases - surely this means that nuclear charge increases as well? But as you go down a group, the number of electron shells increases, and so the outer electrons are found in shells further from the nucleus. This increases atomic radius.

Here's an example. Lithium and sodium are both found in group 1 in the periodic table. They both have one electron in their outer shell. However, sodium has more electron shells than lithium and so has a larger atomic radius.

Fig. 6 - The electron configuration of lithium and sodium, showing atomic radii

Periodic trends in Electronegativity

Now let's look at trends in electronegativity.

Shared pairs of electrons are always found in the outer shell of an atom. Electronegativity is all to do with the strength of the attraction between these bonded electrons and the atom's nucleus. Electronegativity depends on factors such as nuclear charge, number of electron shells, and shielding by inner electrons.

Electronegativity across a period

Electronegativity increases across a period. We have learnt that not only does nuclear charge increase as you move along a period, but also that atomic radius decreases. This brings the bonded pair of electrons closer to the nucleus. But we can't just look at the total nuclear charge - we need to look at the overall nuclear charge that the bonded electrons experience, thanks to shielding by inner electron shells.

We also know that elements in the same period have the same number of electron shells. This means that they have the same levels of shielding. If nuclear charge increases across a period, but shielding remains the same, then the overall nuclear charge felt by the bonded pair of electrons must increase. Thus, electronegativity increases.

Let's use carbon and oxygen as an example again. Carbon has six protons in its nucleus. However, the charge of two of these protons is shielded by two inner shell electrons. Oxygen has eight protons in its nucleus. Once again, the charge of two protons is shielded by two inner shell electrons. The outer shell electrons in carbon experience an overall effective nuclear charge of +4, whilst those in oxygen experience an overall effective nuclear charge of +6. Oxygen also has a smaller atomic radius than carbon. Therefore, it has a higher electronegativity.

Fig. 7 - How electron shielding affects the electronegativity of carbon and oxygen

Electronegativity down a group

Electronegativity decreases down a group. Although nuclear charge increases, the number of inner electron shells also increases, and these inner electron shells shield the charge of the additional protons in the nucleus. As a result, the overall nuclear charge felt by the bonded pair of electrons remains the same. But we also know that atomic radius increases as you move down a group, meaning the bonded pair of electrons are further away from the nucleus. Thus, electronegativity decreases.

For example, earlier, we looked at lithium and sodium. Lithium has three protons in its nucleus, but the charge of two of these protons is shielded by two inner shell electrons. Sodium has eleven protons in its nucleus, but the charge of ten of these protons is shielded by ten inner shell electrons. Outer shell electrons in both elements experience an overall nuclear charge of +1. However, sodium has a larger atomic radius than lithium. Therefore, it has a lower electronegativity.

Fig. 8 - Electronegativity trends in the periodic table

Periodic trends in first ionisation energy

Like electronegativity, ionisation energy is all about the attraction between electrons and the nucleus.

It depends on factors such as nuclear charge, atomic radius, and shielding by inner electron shells.

First ionisation energy across a period

Ionisation energy increases across a period. This is because nuclear charge increases and atomic radius decreases, but the number of inner shielding electron shells remains the same. Overall, there is a stronger attraction between the outermost electron and the nucleus, making it harder to remove it.

First ionisation energy down a group

Ionisation energy decreases down a group. Although nuclear charge increases, the number of inner electron shells also increases. These electrons shield the charge of the additional protons in the nucleus. In addition, atomic radius increases, meaning the outermost electron is further from the nucleus, and is, therefore, easier to remove.

Fig. 9 - The first ionisation energies of period 2

Periodic trends in melting and boiling points

All the trends that we've looked at so far have shown clear patterns and trends, either increasing or decreasing as you move along a period or down a group. We can say that they show periodicity. However, melting and boiling points don't show such clear trends.

Melting and boiling points across a period

Melting and boiling points vary across a period and depend on the structure and bonding of the particular element. Let's take period 3 as an example.

Fig. 10 - Melting points of period 3

Can you see that there is no pattern? Melting points rise and fall across the period. Let's look at why this is so in more detail.

Firstly, sodium, magnesium, and aluminium have medium melting points. This is because they bond using metallic bonding, forming giant metallic lattices held together by electrostatic attraction. To melt the metals, you need to overcome this metallic bonding. Aluminium has the highest melting point of the three as it has the strongest metallic bonding. This is because the ions within the structure have both a higher charge and a smaller atomic radius than the ions within sodium and magnesium: aluminium forms 3+ ions whereas sodium and magnesium form 1+ ions and 2+ ions respectively.

Secondly, silicon has a high melting point. This is because it is a giant covalent macromolecule. All of its atoms are held together by strong covalent bonds, which require a lot of energy to be broken.

Thirdly, phosphorous, sulphur, and chlorine have low melting points. This is because they are simple covalent molecules. Although there are strong covalent bonds within the molecules, the only forces between molecules are weak intermolecular forces which don't require much energy to break. Sulphur has a higher boiling point than phosphorous and chlorine because it forms larger molecules. This increases the strength of the intermolecular forces found between its molecules.

Finally, argon has a very low melting point. This is because it is a monatomic gas. It has extremely weak intermolecular forces between atoms which require hardly any energy to overcome.

Melting and boiling points down a group

Melting and boiling points also vary as you move down a group. For some groups, such as the halogens (group VII), melting points increase as you move down the group. But for other groups, such as the alkali metals (group I), melting points decrease as you move down the group. There's no one clear trend that can be applied to all groups - if only life was that simple!

Summary of periodic trends

That's a lot of different trends you need to know about! To help consolidate your knowledge, we've put together a handy table that summarises the periodic trends we explored above.