Ionization Potential And Electron Affinity

Subject: ChemistryCourse code: Chem-101

Submitted By: Aqsa ManzoorDiscipline: Zoology

Semester: IRoll No: 2180

Submitted To: Mrs. Uzma Imtiaz

Periodic TrendsPeriodic Trends

Ionization Energy, Electron Affinity and Atomic Radii.

The valence electron structure of atoms can be used to explain various properties of atoms.

In general, properties correlate down a group of elements.

In this Assignment I will discuss the trends of Ionization potential and electron affinity in the periodic table.

Ionization EnergyIonization Energy

Ionization : The process of removing an electron from an isolated atom to form a positive ion.

First ionization energy (I1): amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a cation.

Second ionization energy (I2): amount of energy required to remove a second electron from the gaseous monopositive cation to form dipositive cation.

Ionization energies are usually expressed in electron volts (eV) per atom or in kilojoules per mol (kJ/mol) 1eV/atom=96.48 kJ/mol

Value of each ionization energy will increase with each removed electron, since the attractive influence of the nucleus increases and will and will require more energy for the removal of an electron from more positive charges.

Ionization energies measure hoe tightlyelectrons are bound to atoms. Low energies indicate ease of removal of electrons and vice versa.

Factors affecting the magnitude of Ionization Potential:

1. Effective nuclear charge2. Atomic size i.e. atomic radius 3. Principle quantum number 4. Shielding effect 5. Half filled and completely filled

orbitals6. Nature of orbitals7. The extent of penetration of valence

electrons

Effective Nuclear Charge: Greater the magnitude of effective nuclear charge, higher is the amount of energy needed to remove the outermost shell electron. Thus with the increase of the magnitude of effective nuclear charge, the magnitude of ionization potential also increases. The effective nuclear charge increases from left to right in a period.

Atomic size: Greater is the atomic size of an atom, more far is the outermost shell electron from the nucleus and hence lesser will be the force of attraction exerted by the nucleus on the outermost shell electron. Thus higher the value of atomic radius of an atom, lower will be the ionization energy.

Principal Quantum Number (n): Greater is the value of n for the valence shell electron of an atom, further away this electron will be from the nucleus and hence lesser will be the force of attraction exerted by the nucleus on it so lesser energy will be required to remove the valence shell electron. Thus with the increase ot the principal quantum number of the orbital from which the electron is to be removed, the magnitude of ionization potential decreases.

Shielding Affect: The magnitude of shielding effect determines the magnitude of the force of attraction between the nucleus and the valence-shell electron. Greater is the magnitude of shielding effect working on the valence shell electron. Thus with the increase of shielding effect ionization potential increases.Half-filled and completely-filled orbitals: According to Hund’s rule, half-filled (ns1, np3, nd5) or completely-filled (ns2. ns6, nd10) orbitals are comparatively more stable and hence more energy is needed to remove an electron from such orbitals. Thus the ionization potential of an atom having half-filled or completely-filled orbitals in its electronic configuration is relatively higher than that expected normally from its position in the periodic table .

Nature Of Orbitals: The nature of orbitals of the valence-shell from which the electron is to be removed also influences the magnitude of ionization potential. The relative order of energy of s, p, d and f orbitals of a given nth shell is as: ns < np < nd < nfThis order clearly shows that to remove an electron from f-orbital will be easiest while to remove the same from s-orbital will be the most difficult.

The extent of penetration of valence electrons: The degree of penetration of valence electrons in a given principal energy level decreases in the order s>p>d>f, since ns electron is more tightly bound than any np electron, which in turn is more tightly bound than any nd electron etc.

Trends in Ionization Potential:

Ionization energy generally increases from left to right in a period because of the increase in nuclear charge and decrease in atomic radius.

Ionization energies generally decrease down a group due to the shielding effect and increase in atomic size.

Departures from these trends can usually be traced to repulsion between electrons, particularly electrons occupying the same orbitals.

It requires more energy to remove each successive electron.

When all valence electrons have been removed, the ionization energy takes a quantum leap.

Trends in First Trends in First Ionization EnergiesIonization Energies

As one goes down a column, less energy is required to remove the first electron.

For atoms in the same group, Zeff is essentially the same, but

the valence electrons are farther from the nucleus.

Generally, as one goes across a row, it gets harder to remove an electron. As you go from left to right, Zeff increases. The second occurs between Groups VA and VIA.

Electron removed

comes from doubly

occupied orbital.

Repulsion from other electron in orbital helps in its removal.

Elements in the lower left of the periodic table tend to have lower ionization energies than those in the upper right.

These are the elements in the s block, d block, f block and the lower left of the p block - metallic solids

Electron AffinityElectron Affinity

Electron Affinity: energy that occurs when an electron is added to a neutral atom in the gaseous state to form a negative ion.

X(g) + e- X- (g)E = electron attachment energy

Electron affinity = - E (electron attachment)

Second Electron Affinity: Second electron affinity of n element M(g) is defined as the amount of energy required to add one more electron to its mononegative anion, to form dinegative anion.

The addition of second electron to uni-negative ion is an endothermic process.

Factors Affecting Electron Affinity:

Nuclear charge: More the nuclear charge of the atom more strongly will it attract additional electron. Therefore, electron affinity increases as the nuclear charge increases.

Atomic size: The smaller the size of atom smaller will be the distance between the extra electron and the nucleus. Therefore, electrostatic force of attraction will be more and the electron affinity will be higher.

Electronic Configuration: Atoms having stable electronic configuration (i.e. those having completely filled or half filled outer orbitals) do not show much tendency to add extra electron, so have either zero or very low electron affinities.

Periodic Trends:

Electron affinity values generally increase on moving left to right in a period. Electron affinities undergo a general decrease down a group.Electron affinity is usually exothermic.Electron affinities can be obtained by using Born-Haber cycle.

In general, electron affinity becomes more exothermic as you go from left to right across a row.

The first occurs between Groups IA and IIA.◦ Added

electron must go in p-orbital, not s-orbital.

◦ Electron is farther from nucleus and feels repulsion from s-electrons.

Main Group Main Group Elements Elements (s and p)(s and p) An s-block element has low

ionization energy; outer electrons can easily be lost.

Group 1 form +1 ions; Group 2 form +2 ions

An s block element is likely to be a reactive metal

p- blockElements on the left have low

enough ionization energies to be metallic, but higher than the s-block elements and so are less reactive

Elements on the right have high electron affinities (tend to gain electrons to form closed shell ions)