User:Chaccocat/Periodic trends

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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 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.

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Electrophilicity refers to the tendency of an electron-deficient species, an electrophile, to accept electrons. Nucleophilicity is defined as an electron-rich species, a nucleophile, and its affinity to donate electrons to another species. Trends in the periodic table can help predict an element's nucleophilicity and electrophilicity. Nucleophilicity generally increases as electronegativity increases, meaning the nucleophilicity for elements increases left-to-right on the periodic table. Electronegativity increases as electronegativity decreases, meaning electronegativity decreases left-to-right on the periodic table.

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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.

Trend-wise, 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.

Lead: Electron Affinity

Electron affinity between two electronegative atoms (i.e., Chlorine and Bromine) decreases as the space between the valence shell and nucleus increases.

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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. Similarly to ionization energy, this is caused by the increase in atomic size due to the addition of a valence shell, which weakens 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.

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Nuclear charge is defined as the number of protons in the nucleus of the relevant element. Electrons of multi-electron atoms do not experience the entire nuclear charge due to shielding effects from the electrons on other atoms. Effective nuclear charge refers to the nuclear charge of atoms that experience shielding.

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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.