Explaining Periodic Trends! Many observable trends in the chemical and physical properties of elements are observable in the periodic table.! On trends you may be familiar with is reactivity, which is high in Group 1 elements, lower in the middle of the table, and high again in group 17. Noble gases are the least reactive. Within a group, reactivity increases as you move down the group for metals, and increases as you move up a group for non-metals.! The location of the elements in the periodic table can be used to predict their relative atomic radii, ionization energy, electron affinity and electronegativity, which in turn determine the chemical properties of the elements. A) Atomic Radius Atomic radius (r):the distance from the centre of an atom to the boundary in which the electrons spend 90% of their time. The size of an atom decreases going from left to right across a period. This is because as you move down the period, the number of positive protons in the nucleus increases. This increases attractive forces, and the attractive force on the outer electron shell gets stronger.
The size of an atom increases as you go down a group. This is because as you move down the groups, the number of occupied electron shells increases. Effective nuclear charge: Apparent nuclear charge as experienced by outermost electrons. As you go down group, the filled inner shells of an atom shield the outer shells from the positive charge of the nucleus. Because of this shielding, the effective nuclear charge for the outer electrons is smaller and radius increases. B) Ionization Energy The amount of energy required to remove an electron from an atom.
The first ionization energy refers to the amount of energy required to remove the outermost electron. The second ionization energy refers to the amount of energy required to remove the second outermost electron and so on. The outermost electron is the easiest electron to remove. Removal of subsequent electrons requires more energy. In terms of increasing energy: 1st Ionization Energy < 2nd Ionization Energy < 3rd Ionization Energy. Ionization Energy increases as you move left to right across a period. This is because as the atom gets smaller, the valence electrons become closer to the nucleus. This means the attractive force holding the electron is stronger and it takes more energy to pull the electron off. Ionization Energy decreases and you move down a group. This is because of shielding. The inner shells that are filled shield the outer shells from the positive charge of the nucleus, making outer electrons easier to remove.
C) Electron Affinity The amount of energy released when an electron is added to an element. The higher an elementʼs electron affinity the greater the attraction for an electron. Electron Affinity increase as you move left to right across a period. This is because as you move across a period, the attraction for electrons increases. Electron Affinity decreases as you move down a group. This is because as you move down a group, the attraction for electrons decreases.
D) Electronegativity The ability of an atom to attract an electron away from another atom. Elements with high electronegativity have a strong tendency to gain an electron or electrons from other atoms, whereas elements with low electronegativity have a strong tendency to lose an electron or electrons to other atoms. Electronegativity increases going from left to right across a period. This is because of the the relationship between electronegativity and radius. The farther an electron is from the nucleus, the easier it is to remove. Electronegativity decreases going down a group. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius.