In chemistry, periodic trends are specific patterns present in the periodic table that illustrate different aspects of certain elements when grouped by period and/or group. They were discovered by the Russian chemist Dmitri Mendeleev in 1863. Major periodic trends include atomic radius, ionization energy, electron affinity, electronegativity, nucleophilicity, electrophilicity, valency, nuclear charge, and metallic character. Mendeleev built the foundation of the periodic table. Mendeleev organized the elements based on atomic weight, leaving empty spaces where he believed undiscovered elements would take their places. Mendeleev's discovery of this trend allowed him to predict the existence and properties of three unknown elements, which were later discovered by other chemists and named gallium, scandium, and germanium. English physicist Henry Moseley discovered that organizing the elements by atomic number instead of atomic weight would naturally group elements with similar properties.
Summary of trends
Atomic radius
The atomic radius is the distance from the atomic nucleus to the outermost electron orbital in an atom. In general, the atomic radius decreases as we move from left-to-right in a period, and it increases when we go down a group. This is because in periods, the valence electrons are in the same outermost shell. The atomic number increases within the same period while moving from left to right, which in turn increases the effective nuclear charge. The increase in attractive forces reduces the atomic radius of elements. When we move down the group, the atomic radius increases due to the addition of a new shell.
Nuclear charge and effective nuclear charge
Nuclear charge is defined as the number of protons in the nucleus of an element. Thus, from left-to-right of a period and top-to-bottom of a group, as the number of protons in the nucleus increases, the nuclear charge will also increase. However, electrons of multi-electron atoms do not experience the entire nuclear charge due to shielding effects from the other electrons. In this case, the nuclear charge of atoms that experience this shielding is referred to as effective nuclear charge. Shielding increases as the number of an atom's inner shells increases. So from left-to-right of a period, the effective nuclear charge will still increase. But, from top-to-bottom of a group, as the number of shells increases, the effective nuclear charge will decrease.
Ionization energy
The ionization energy is the minimum amount of energy that an electron in a gaseous atom or ion has to absorb to come out of the influence of the attracting force of the nucleus. It is also referred to as ionization potential. The first ionization energy is the amount of energy that is required to remove the first electron from a neutral atom. The energy needed to remove the second electron from the neutral atom is called the second ionization energy and so on. As one moves from left-to-right across a period in the modern periodic table, the ionization energy increases as the nuclear charge increases and the atomic size decreases. The decrease in the atomic size results in a more potent force of attraction between the electrons and the nucleus. However, suppose one moves down in a group. In that case, the ionization energy decreases as atomic size increases due to adding a valence shell, thereby diminishing the nucleus's attraction to electrons.
Electron affinity
The energy released when an electron is added to a neutral gaseous atom to form an anion is known as electron affinity. Trend-wise, as one progresses from left to right across a period, the electron affinity will increase as the nuclear charge increases and the atomic size decreases resulting in a more potent force of attraction of the nucleus and the added electron. However, as one moves down in a group, electron affinity decreases because atomic size increases due to the addition of a valence shell, thereby weakening the nucleus's attraction to electrons. Although it may seem that fluorine should have the greatest electron affinity, its small size generates enough repulsion among the electrons, resulting in chlorine having the highest electron affinity in the halogen family.
Electronegativity
The tendency of an atom in a molecule to attract the shared pair of electrons towards itself is known as electronegativity. It is a dimensionless quantity because it is only a tendency. The most commonly used scale to measure electronegativity was designed by Linus Pauling. The scale has been named the Pauling scale in his honour. According to this scale, fluorine is the most electronegative element, while cesium is the least electronegative element. Trend-wise, as one moves from left to right across a period in the modern periodic table, the electronegativity increases as the nuclear charge increases and the atomic size decreases. However, if one moves down in a group, the electronegativity decreases as atomic size increases due to the addition of a valence shell, thereby decreasing the atom's attraction to electrons. However, in group XIII (boron family), the electronegativity first decreases from boron to aluminium and then increases down the group. It is due to the fact that the atomic size increases as we move down the group, but at the same time the effective nuclear charge increases due to poor shielding of the inner d and f electrons. As a result, the force of attraction of the nucleus for the electrons increases and hence the electronegativity increases from aluminium to thallium.
Valency
… excerpt ends here. Continue reading the full article.
