4.6 Periodic Properties of Elements

The periodic properties of elements refer to certain characteristics that show a repeating pattern as you move across periods (rows) or groups (columns) in the periodic table. These properties include metallic and non-metallic properties, atomic radius, ionization energy, electronegativity, and electron affinity.

a. Metallic Property

• Definition: Metals are elements that are shiny, conduct heat and electricity, and produce a metallic sound when struck. Modern definition: metals are elements that can lose one or more electrons to form positive ions.

• Trend in the Periodic Table:

• As you move left to right across a period, elements tend to lose their metallic properties. This means that metallic property decreases as you move from left to right in a period.

• Example: Lithium (Li) is a metal because it can lose an electron to form a positive ion:

• Li → Li⁺ + e⁻

b. Non-metallic Property

• Definition: Non-metals are elements that are not shiny, do not produce metallic sounds, and are poor conductors of heat and electricity. They tend to gain electrons to form negative ions.

• Trend in the Periodic Table:

• As you move right to left across a period, non-metallic properties increase. This means that the further right you go, the more non-metallic the elements become.

• Example: Chlorine (Cl) is a non-metal because it accepts an electron to form a negative ion:

• Cl + e⁻ → Cl⁻

• Sub-metals: Some elements, like silicon (Si), can exhibit properties of both metals and non-metals, depending on the conditions. These elements are called sub-metals.

c. Atomic Radius/ Size of Atom

• Definition: Atomic radius refers to the size of an atom, or the distance from the nucleus to the outermost electrons.

• Trend in the Periodic Table:

• Across a period (left to right): The atomic radius decreases. As the atomic number increases, more protons and electrons are added, causing a stronger pull on the electrons, which brings them closer to the nucleus.

• Down a group (top to bottom): The atomic radius increases because a new electron shell is added for each period, making the atom larger despite the addition of more protons.

d. Ionization Energy

• Definition: Ionization energy is the amount of energy required to remove one mole of electrons from one mole of atoms in their gaseous state, forming positive ions.

• Trend in the Periodic Table:

• Across a period (left to right): Ionization energy increases because the atomic radius decreases, causing a stronger attraction between the nucleus and electrons, making it harder to remove an electron.

• Down a group (top to bottom): Ionization energy decreases because the atomic radius increases, causing the outer electrons to be farther from the nucleus, and easier to remove.

• Example:

• Silicon (Si) has a greater ionization energy than sodium (Na) because Si has a smaller atomic radius.

• Lithium (Li), with the smallest atomic radius among the alkali metals (Group 1), has the highest ionization energy.

e. Electron Affinity

• Definition: Electron affinity is the energy released when an electron is added to an atom in its gaseous state to form a negative ion.

• Trend in the Periodic Table:

• Across a period (left to right): Electron affinity increases because the atomic radius decreases, making it easier to attract an electron.

• Down a group (top to bottom): Electron affinity decreases because the atomic radius increases, making it harder to attract an electron.

• Example:

• Be (Beryllium) has a higher electron affinity than Ra (Radium) because Be has a smaller atomic radius.

• Among elements in Period 3 (Na, Mg, Al, Si), Si has the highest electron affinity because it has the smallest atomic radius, while Na has the least because it has the largest atomic radius.

f. Electronegativity

• Definition: Electronegativity is the ability of an atom in a molecule to attract electrons towards itself when covalently bonded.

• Trend in the Periodic Table:

• Across a period (left to right): Electronegativity increases because the atomic radius decreases, leading to a stronger attraction for electrons.

• Down a group (top to bottom): Electronegativity decreases because the atomic radius increases, weakening the attraction for electrons.

• Example: In Period 3, Cl (Chlorine) has the highest electronegativity, while Na (Sodium) has the lowest.