When moving across the periodic table from Li to Be to B, the first ionization energy increases from Li to Be, then drops for B. The first ionization energy of B is greater than that of Li. Explain why.

• When moving across a period on the periodic table, the value of the effective nuclear charge increases with atomic number. This causes a general increase from Li to Be to B. 
• The even higher value of Be (greater than B) is due to the increased stability of the electron configuration of Be. 
• Beryllium has a filled s-subshell. Filled subshells have an increased stability, and additional energy is required to pull an electron away. 

Define the first and second ionization energy. State why the second ionization energy of elements is higher than first ionization energy?

• The first ionization energy is the energy it takes to remove an electron from a neutral atom.  
• The second ionization energy is the energy it takes to remove an electron from a 1+ ion.   
• After the removal of first electron, force of attraction between the nucleus and remaining outer most electrons increase. So, to remove 2nd electrons more amount of energy is required. Hence, the second ionization energy is higher than first ionization energy.  

How are the trends of atomic radius and ionization energy related? Explain in both groups and periods.

These trends are inverses of each other. As the size of the atom increases, the ionization energy (energy required to remove the electron) decreases because the electron being removed is farther from the nucleus in a larger atom.Less energy is required to remove an electron further from the nucleus. [AR increases down a group while IE decreases; AR decreases across a period while IE increases.]

What’s the difference between electronegativity and ionization energy?

Electronegativity is related to the pull of a nucleus on an electron being shared in a bond. Ionization energy is the energy required to remove an electron from an unbonded atom. Electronegativity is a ranking and has no units. Ionization energy is an amount of energy and has units such as kilojoules.

SPM Form 4: Periodic Table of elements (Checklist)

• Periodic Table: the table showing the elements in order of increasing proton number; similar elements are arranged in columns called groups.
• Group: A vertical column of elements in the Periodic table.
• Period: A horizontal row of the Periodic Table; its number tells you how many electron shells there are.
• Alkali metals: the Group I elements of the Periodic Table, which include lithium, sodium, potassium, rubidium, caesium and francium.
• Alkaline earth metals: the Group II elements of the Periodic Table, which include beryllium, magnesium, calcium, strontium, barium, and radium.
• Halogens: the Group VII elements of the Periodic Table, which include fluorine, chlorine, bromine, iodine and astatine.
• Noble gases: the Group 18 elements of the Periodic Table; they are called ‘noble’ or inert gases because they are so unreactive, which include helium, neon, argon, krypton, xenon and radon.
• Transition elements: the elements in the wide middle block of the Periodic Table (elements in group 3 to group 12).
• Metal: an element that shows metallic properties (for example conducts electricity, and forms positive ions)
• Non-metal - an element that does not show metallic properties: the non-metals lie to the right of the zig-zag line in the Periodic Table.
• Amphoteric oxide: An oxide that exhibits both acidic and basic properties.
• The atomic radius is a term used to describe the size of the atom.
• The ionization energy is the energy required to completely remove an electron from a gaseous atom or ion.
• Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom.
• Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. The higher the electronegativity of an atom, the greater its attraction for bonding electrons.

Factors that affect the size of the first ionization energy

The first ionization energy is the energy required to remove one electron from the parent atom. There are four factors that affect the size of the first ionization energy:

• The charge on the nucleus
• The distance of the electron from the nucleus
• The number of electrons between the outer electrons and the nucleus
• Whether the electron is alone or paired in its orbital

Factors Affecting the Ionization Energy

• Atomic radius: The valence electrons of an atom with a larger radius experience a less attraction towards nucleus, hence possesses low ionization energy.
• Effective nuclear charge:  The higher the nuclear charge the stronger the attraction between the nucleus and electrons. This causes the ionization energy to increase.
• Shielding effect (screening effect): The shielding effect of the electrons of the inner orbitals causes the outer electrons to be less attracted to the nucleus and thus decrease the magnitude of ionization energy.

Ionization Energy on Periodic Table

The ionization energy is the energy involved in removing one mole of electrons from one mole of atoms in the gaseous state.

Across a period from left to right, the ionization energy increases. This is due to the increase in atomic charge having a greater pull on the electrons and therefore more energy is required to remove electrons.

Going down a group, the ionization energy decreases. This is due to the outer electrons being further away from the nucleus and so the attraction is weaker and they are more easily removed.

Electronegativity and Ionization Energy

Electronegativity is related to ionization energy. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus.