Periodic Trends

Learning Targets/Content Objectives:

I can explain trends such as, Ionization energy, atomic and ionic radii, electron affinity and electronegativity in the periodic table using the position of the element in the periodic table, Coulomb's law, the shell model and the concept of shielding and the effective nuclear charge.

Summary of periodic table trends

Coulomb's law is used to explain trends in the periodic table including. atomic radius, ionization energy, electronegativity and electron affinity.

Coulomb's law describes the magnitude of the electrostatic force between two charged particles as follows. The force is proportional to the amount of charge that the two particles have and is inversely proportional to the distance that separates them. So, when we are talking about atoms we can consider Q1 as the effective nuclear charge of the atom and Q2 as the charge of the electron.

Q1 id considered as the effective nuclear charge represented by Z effective

Z effective= atomic # Z- Shielding effect S

S is estimated to be the same as the number of core electrons

What is the Effective Nuclear Charge?

The effective nuclear charge on an electron is just the amount of charge it experiences if we keep track of the repulsions of inner electrons. We will say that inner electrons shield or screen the electron from the charge of the nucleus, so that the charge that is effective at attracting the electron is the total charge minus the charge of the electrons that shield it. As an equation we can write

Z effective = Z - S

Below is an example of finding the effective nuclear charge of Lithium atom.

Z eff of Lithium

We can generalize by saying that as the effective nuclear charge increases, the force of attraction between the nucleus and the outermost shell increases (effective nuclear charge increase across the periodic table from left to right and therefore the force of attraction increases as we move from left to right across the periodic table.

Atomic Radius Trend

Atomic radius increases down a group because atoms have more energy levels/shells and therefore the distance between the nucleus and the electrons in the outermost shells increase and the attraction between the nucleus and the electrons in the outermost shell becomes smaller.

The stronger the Coulombic force is, the more tightly the electrons are attracted to the nucleus and the smaller the atomic radius will be. Therefore, we can say that the atomic radius decrease as we move from left to right in the periodic table because of the increased effective nuclear charge. Increased effective nuclear charge, increases the attractive force fo the nucleus and as a result it pulls the electron cloud closer resulting in a smaller atomic radius.

The diagram below shows how increased effective nuclear charge yields a stronger Coulombic force between the nucleus and the electrons and results in a smaller atomic radius. This explains how the atomic radius decrease as we move across the periodic table from left to right.

Ionization Energy

Ionization energy is the quantity of energy that an isolated, gaseous atom in the ground electronic state must absorb to discharge an electron, resulting in a cation.

H(g) →H+ (g) + e−

Ionization energy is usually expressed in kJ/mol, or the amount of energy it takes for all the atoms in a mole to lose one electron each.

Ionization Energy Trend

Ionization energy increases across the periodic table from left to right mainly because the effective nuclear charge of the atom gets bigger as we move from left to right. The bigger the effective nuclear charge of an atom, the stronger is the Coulombic force of attraction between the nucleus and the valence electrons and therefore more energy is needed to remove an electron from the atom. When the electron is strongly attracted by the nucleus it needs more energy to be taken away from the atom.

Ionization energy decreases down a group because the number of energy levels (n) increases and as a result the distance between the nucleus and the electrons increases. The increase in the distance causes less Coulombic (electrostatic) attraction between the electrons and the protons in the nucleus making the electron less attracted to the nucleus and easier to be removed.

Electron Affinity

Is the energy released when an electron is added to a gaseous atom to form a negatively charged ion as shown in the reaction below

For example, when a fluorine atom in the gaseous state gains an electron to form F⁻(g), the associated energy change is -328 kJ/mol. Because this value is negative (energy is released), we say that the electron affinity of fluorine is favorable.

Watch this video for a better understanding of the trends in the periodic table.

Hints when explaining trends in the periodic table

How-to-explain-Periodic-Trends.pdf

Check Your Understanding

Arrange the following elements in the order of increasing atomic radius (from smallest to largest). Justify your answer

Mg, Be, Ba, Sr and Ca

Be, Mg, Ca, Sr and then Ba and that is because down a group in the periodic table the principal quantum number increases which means that more shells are added to the atom and therefore the distance between the nucleus and the outermost shell increases resulting in bigger atomic radius.

2. Which atoms has higher ionization energy O or F ? Justify your answer

F has a higher Ionization energy than O and that is because F has a bigger effective nuclear charge, and because F has bigger effective nuclear charge, the valence electrons in F are held to the nucleus by stronger Coulombic forces than the forces in O and that makes it harder to lose an electron from F.

3. what evidence do you use to explain periodic trends across the periodic table from left to right?

Effective nuclear charge and its correlation to Coulomb's law

4. how would you justify periodic trends down a group in the periodic table?

Use increased number of shells and increased distance from the nucleus and its correlation to Coulomb's law.

Report abuse