Trends of periodic atomic properties

Periodic TableTrends in periodic atomic properties

Sharifah Mona Abdul Aziz AbdullahAbdul Al-Hafiz Ismail

Centre for Pre-University StudiesUniversiti Malaysia Sarawak

This OpenCourseWare@UNIMAS and its related course materials are licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License.

Atomic radius

Atomic radius is defined as one-half of the average distance between the two covalently bonded atoms.

1.99 Å

Atomic radius of chlorine

Factor affecting the size of atoms:

a) Effective nuclear chargeThe charge felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus.

Atomic radius

Atomic radius

b) Shielding electron/effect

Shielding electrons are the electrons in the energy levels between the nucleus and the valence electrons

They are called "shielding" electrons because they "shield" the valence electrons from the force of attraction exerted by the positive charge in the nucleus.

Atomic radius

Trend in atomic radius across a period and down the group in periodic table:

Atomic radius decreases

Atom

ic ra

dius

incr

ease

s

Atomic radius

Decrease across a period• Proton number increases• Effective nuclear charge increases• Stronger attraction between nucleus and valence electron• Thus, atomic radius decreases

Increases down a group• Energy level increases• Shielding effect increases• Weaker attraction between nucleus and valence electrons• Thus, atomic radius increases.

Example: Arrange the following atoms in order of increasing atomic size.

Na, K, Cl, Ba, B, C, Ag

Atomic radius

Atomic radius

Solution:

Ba is at Period 6, the lowest position among all the elements, should be the largest atom.

Next should be Ag (Period 5) and then K (Period 4)

Both Na and Cl are at Period 3, but because of Na is at the left side of the table, Na is larger than Cl.

Same goes between B and C at Period 2.

Therefore,C < B < Cl < Na < K < Ag < Ba

Ionic Radius

Positive ion always has a smaller ionic radius than the original atom. Why?

The loss of electron(s) means that the remaining electrons each have a greater share of the positive charge of the nucleus so are more tightly bound. When the ion is formed a whole electron shell is usually lost.

Ionic Radius

Negative ion has a larger radius than that of the original atom. Why?

Even though the additional electrons are in the same shell as existing electrons, the addition of extra negative charge means that the electrons are less tightly bound to the nucleus and so the radius is larger.

Ionic Radius

Neutral atoms or ions that have same number of electrons and the same electronic configuration are said to be isoelectronic.

Examples:

Na atom (1s2 2s2 2p6 3s1) contain one extra shell compare to Na+ ions (1s2 2s2 2p6)

F- ion (1s2 2s2 2p6) contain an extra negative electron compare to F atom (1s2 2s2 2p5)

Therefore, Na+ and F- are isoelectronic.

Ionization Energy (IE)

Ionisation energy is the energy required to remove the valence electron from an atomic species.

First IE can be represented by an equation as:M(g) M+(g) + e-

Second IE can be represented by an equation as:M+ (g) M2+(g) + e-


Factor affecting the magnitude of ionisation energy:

a) Atomic radius• The valence electron of an atom with a larger

radius experience less attraction to the nucleus. Hence, the atom have lower IE.

b) Effective nuclear charge• Atom with higher effective nuclear charge

holds its electron closer to the nucleus than the atom of a lower effective nuclear charge.

• Hence, the IE increases as the effective nuclear charge become larger.



c) Shielding effect• The shielding effect of inner electrons causes

the valence electrons to be less strongly attracted to the nucleus and thus results in lower IE.


General trend in first ionisation energy across a period and down the group:

Ionisation energy increases

Ioni

satio

n en

ergy

dec

reas

es

Increases across a period• Proton number increases• Effective nuclear charge increases• Stronger attraction between nucleus and valence

electron• Atomic radius decreases• More difficult to remove valence electron• Thus, IE increases.


Decreases down a group• Energy level increases• Shielding effect increases• Weaker attraction between nucleus and valence

electron• Atomic radius increases• Easier to remove valence electron• Thus, IE decreases


Electronegativity

• Electronegativity is a measure of an atom’s attraction for another atom’s electrons.

Electronegativity increases

Elec

tron

egati

vity

incr

ease

s

Trend in electronegativity across a period and down the group:

Increases across a period• Proton number increases• Effective nuclear charge increases• Stronger attraction between nucleus and valence

electron• Atomic radius decreases• Greater ability to attract the bonding electrons to itself• Thus, electronegativity increases

Electronegativity

Electronegativity

Decreases down a group• Energy level increases• Shielding effect increases• Weaker attraction between nucleus and valence

electron• Atomic radius increases• Weaker ability to attract the bonding electrons to

itself.• Thus, the electronegativity decreases

Science

Trends of periodic atomic properties