OBJECTIVES

• To list understand the factors that affect​ the columb force of attarction

• To understand how shielding and nuclear effective charge influence the nuclear force of attraction

• To differentiate between atomic radius and Van-dar Waal radius

• To compare the radius of different elements in the periodic table and explain the differences in size

• To identify and explain exceptions in atomic radius trend

• To compare the size of ions to their orginal atoms​

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Coulombic force

Shielding

​Shielding and effective nuclear charge

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• For multielectron atoms, any one electron experiences both the positive charge of the nucleus (which is attractive) and the negative charges of the other electrons (which are repulsive). Repulsion of one electron by other electrons is screening or shielding that shield electron from the full effects of the nuclear charge.

• The nuclear charge is the charge on the nucleus (number of protons) Effective nuclear charge (Z eff) of the electron is the total amount of attraction that an electron feels for the nucleus or a net positive charge that is attracting a particular electron

Effective nuclear charge

To determine the radius of an atom, the distance between two non-bonded atoms is measured. Half this distance is the radius and is called van der Waals radius. If atoms are chemically bonded, then the radius has a different value and is found to be different in non- metals and metals. In non-metals, it is half the distance between the nuclei of the two atoms bonded together. In metals, it is half the distance between the nuclei if two atoms in a crystal of the metal. When an atom bonds to another atom, the radius is always less than the van der Waals radium.

​Demo

​Demo

​Demo

Atomic Radius

• As you go across the period (row) Atomic Radius decreases.

• This is caused by an increasing positive charge in the nucleus.

• Electrons are on the same energy level going across the period, but more protons are pulling electrons in closer. This results in a smaller atomic radius.

Atomic Radius

• As you go down a group (column) Atomic Radius increases

• The number of electron shells increase as you move down the periodic table.

• This results in a larger atomic radius because the nucleus has less pull on the outer shell of electrons.

First Ionization Energy (IE1)

It is the minimum amount of energy required to remove one mole loosely bound electrons from one mole of neutral gaseous atom to form one mole of a single positively charged gaseous ion under standard conditions.

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IE1 of Ca : Ca (g) + IE1 → 𝐶𝑎¹⁺ + 1 e- IE1 of O (𝑔)

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O (g) + IE1 → 𝑂¹⁺ + 1 e-

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Second ionization energy​

It is the minimum amount of energy required to remove one mole of electrons from a single positively charged gaseous ion to form a 2+ charged gaseous ion.

IE2 of Ca : 𝐶𝑎¹⁺ (g) + IE2 → 𝐶𝑎²⁺ + 1 e-

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Sucessive ionization energies:

IE1< IE2<IE3<IE4-------

As more electrons are removed, more energy is required to kick electrons outside their atoms/ions​

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Trends in ionization energies

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Ionization Energy

Ionization Energy

• The energy required to remove one electron from the outer shell of a neutral atom.

• This is the energy needed to form an ion.

• When ionization energy is high it is harder to remove that electron.

Ionization Energy

• Moving across a period, ionization energy increases.

• Group one has a low ionization energy and has a high reactivity

• Group 18 has a high ionization energy and are unreactive.

• Caused by an increasing positive charge in the nucleus, attracting electrons more strongly.

Ionization Energy

• decreases as you move down the periodic table.

• Elements with larger atomic radius have less pull on their electrons in the outer shell and tend to give them up.

• Shielding effect also increases due to extra electron repulsion so electrons in the outermost principal level are increasingly farther away from the positively charged nucleus and are therefore held less tightly

​Demo

​Demo

​On the basis of periodic trends, determine which element in each pair has the higher first ionization

energy (if possible). Justify your answer.

a)  Al or S

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a)  Al or S

S has a higher ionization energy than Al As you go across a period from left to right ionization energy increases. For the same shielding effect, the effective nuclear charge increases, so electron are more strongly attracted to the nucleus and more energy is needed to remove the electron so IE increases.

​Demo

​As or Sb

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​Demo

​As or Sb

As has a higher ionization energy than Sb As you go down a group in the periodic table, the size of the orbital increases with increasing principle quantum number (n) so the outermost electron is highly shielded and easily removed. (Shielding effect increases and effective nuclear charge increases).

​Exception in IE1

​Exception in IE1

Successive ionization energies

• Generally, IE2 is always greater than IE1. Explain and justify your answer.

  IE1 involves the removal of an electron from a neutral atom while IE2 involves the removal of an electron from a positive ion where the same nuclear charge is attracting less number of electrons, thus the effective nuclear charge increases and electrons are more strongly attracted to the nucleus.​

• Successive ioniation energies are used to determine the identity of an element

• Large jumps or gaps in ionization energies indicate that electrons have been removed from another energy level

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Successive ionization energies

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Successive ionization energies

Successive ionization energies

​Ionization energy can be used to determine the group the element belongs to as well as the number of valence electrons. By examining the table below, we notice that the first large jump between two successive ionization energies can be used to determine the number of valence electrons.