The periodic table not only organises elements based on their atomic structure but also highlights recurring patterns in their properties. These trends, observed across periods and groups, provide insights into the behaviour of elements and their interactions with others.
Some key trends emerge from the periodic table, including how atomic size (atomic radii) changes across periods and down groups, how much energy is required to remove the first electron from the valence shell (first ionisation energy), how readily an atom can accept an electron (electron affinity), and how strongly atoms attract electrons in chemical bonds (electronegativity).
Atomic radii
Atomic radii (or radius) increases down a group (from top to bottom) of the periodic table. As \(n\) increases when moving down a group, the number of electron shells become larger; the electron orbitals are therefore larger, leading to an increase in atomic radii.
When moving across a period (from left to right), atomic radii decreases. As you move across a period, each successive element has an additional proton in the nucleus and an additional electron in the same valence shell. The increase in positive charge in the nucleus attracts the electrons more strongly, pulling them closer to the nucleus. Although the number of electrons increases, these electrons are added to the same energy level and do not significantly increase electron–electron repulsion that would otherwise counteract the pull of the nucleus.
This stronger attraction between the increased positive charge of the nucleus and the electrons results in a smaller atomic radii as you move across a period.
An exception to the trend in atomic radii across a period is in the transition metals. As you move across the transition metals (\(d\)-block elements), the decrease in atomic radii is less pronounced. This means that elements in the same period within the \(d\)-block will have a similar atomic radii.
First ionisation energy
The first ionisation energy is the minimum energy required to remove the first electron from a neutral atom. The second and third ionisation energies are the energies necessary to remove the second and third electrons from the atom, respectively.
First ionisation energy increases when moving across a period as the attraction between the nucleus and the valence electrons increases. More energy is therefore required to overcome this strong attraction and remove the first valence electron.
When moving down a group, the first ionisation energy decreases. This is because a valence electron in a higher energy level (greater \(n\)) is easier to remove.
Electron affinity
Electron affinity gives an indication of how easy it is for a neutral atom to accept an electron and form a negatively-charged ion (an anion). Quantitatively, it is the amount of energy released when an electron is added to an atom in the gaseous phase to form an anion.
The electron affinity increases when moving across a period. The interaction between the nucleus and the electrons increases, and more energy is released when an electron is added. Non-metals, particularly the halogens, have high electron affinities because they are close to achieving a full valence shell, making it easier for them to gain electrons.
Electronegativity
Electronegativity is the power of an atom in a molecule to attract electrons from another atom to form chemical bonds. The larger the value, the larger the electron-attracting ability.
Electronegativity decreases from top to bottom and increases from left to right of the periodic table. Across a period, atoms have more protons and a higher effective nuclear charge as you move across a period, which increases the attraction for electrons. Down a group, the atomic radii increases and the addition of electron shells leads to a greater distance between the nucleus and the valence electrons, and a reduced attraction for electrons.
Noble gases are generally not considered in electronegativity trends because they have full valence electron shells and are typically inert. They do not tend to form chemical bonds, so electronegativity is not so relevant—although, some heavier noble gases like xenon can form chemical bonds under specific conditions. This means that fluorine is generally considered the most electronegative element.
Periodic table illustrating trends. Atomic radii decrease across periods (pink arrow) and increase down groups. First ionisation energy increases across periods (blue arrow) and decreases down groups. Electronegativity increases across periods (orange arrow) and decreases down groups.
Example – predicting periodic trends
Predict which atom is larger for the following pairs of atoms:
\(\ce{C}\) and \(\ce{F}\)
\(\ce{Na}\) and \(\ce{Cs}\)
\(\ce{Sc}\) and \(\ce{Ni}\)
Step 1: Recall the trend in atomic radii in the periodic table.
When moving across a period (left to right), atomic radii decreases. As you move across a period, the number of valence increases, making the electron cloud more negative. However, the number of protons (atomic number) also increases, making the nucleus more positive. This allows the positively-charged nucleus to pull the electrons in closer, decreasing the atomic radius.
When moving down a group (top to bottom), atomic radii increases. Going down a group, elements gain an electron shell. As there will be a greater distance between the valence electrons and the nucleus, the atomic radius increases.
Step 2: Locate the elements on the periodic table and compare them.
\(\ce{C}\) and \(\ce{F}\) are both located in period \(2\), but \(\ce{F}\) is located further right in group \(17\) compared to \(\ce{C}\) in group \(14\).
\(\ce{Na}\) and \(\ce{Cs}\) are both located in group \(1\), but \(\ce{Cs}\) is located further down in period \(6\) compared to \(\ce{Na}\) in period \(3\).
\(\ce{Sc}\) and \(\ce{Ni}\) are both located in period \(4\), but \(\ce{Ni}\) is located further right in group \(10\) compared to \(\ce{Sc}\) in group \(3\).
Step 3: Apply your understanding of the relative locations of the elements on the periodic table with the trend in atomic radii to predict which atom is larger.
\(\ce{C}\) would be expected to have a bigger atomic radius than \(\ce{F}\) and would therefore be larger.
\(\ce{Cs}\) would be expected to have a bigger atomic radius than \(\ce{Na}\) and would therefore be larger.
\(\ce{Sc}\) would be expected to have a bigger atomic radius than \(\ce{Ni}\) and would therefore be larger. However, across \(d\)-block and \(f\)-block, atomic radii change is not significant. Thus, in reality, \(\ce{Sc}\) and \(\ce{Ni}\) have approximately the same radius.
Your turn – trends in the periodic table
Test yourself on your understanding of trends in the periodic table.
