Topic 5_Atomic structure

Figure 8.28 Ionic radius.

Figure 8.29 Ionic vs. atomic radii.

Sample Problem 8.8 Ranking Ions by Size

PROBLEM: Rank each set of ions in order of decreasing size, and
explain your ranking:
(a) Ca2+, Sr2+, Mg2+ (b) K+, S2−, Cl− (c) Au+, Au3+

PLAN: Find the position of each element on the periodic table and
apply the trends for ionic size.

SOLUTION:
(a) Sr2+ > Ca2+ > Mg2+
All these ions are from Group 2, so size increases down the group.

Sample Problem 8.8

SOLUTION:
(b) S2− > Cl− > K+
These ions are isoelectronic, so size decreases as nuclear charge
increases.

(c) Au+ > Au3+
Cation size decreases as charge increases.

Electron configurations of Monatomic Ions

Elements at either end of a period gain or lose electrons
to attain a filled outer level. The resulting ion will have a
noble gas electron configuration and is said to be
isoelectronic with that noble gas.

Na(1s22s22p63s1) → e– + Na+([He]2s22p6)
[isoelectronic with Ne]

Br([Ar]4s23d104p5) + e– → Br- ([Ar]4s23d104p6)
[isoelectronic with Kr]

Figure 8.25 Main-group elements whose ions have noble gas
electron configurations.

Sample Problem 8.6 Writing Electron Configurations of Main-Group
Ions

PROBLEM: Using condensed electron configurations, write reactions for
the formation of the common ions of the following elements:

(a) Iodine (Z = 53) (b) Potassium (Z = 19) (c) Indium (Z = 49)

PLAN: Identify the position of each element on the periodic table
and recall that:

• Ions of elements in Groups 1A(1), 2A(2), 6A(16), and 7A(17)
are usually isoelectronic with the nearest noble gas.

• Metals in Groups 3A(13) to 5A(15) can lose the ns and np
electrons or just the np electrons.

5.4 Ionization Energies

Ionization energy (IE) is the energy required for the
complete removal of 1 mol of electrons from 1 mol of
gaseous atoms or ions. Unit kJ/mol

Pulling an electron away from a nucleus requires energy to
overcome electrostatic attraction between negatively
charged electron and positively charged protons in the
nucleus. There is a first IE, second and so forth.

Atoms with a low IE tend to form cations during reactions.
Atoms with a high IE tend to form anions (except the noble
gases).

Ionization energy tends to decrease down a group and
increase across a period.

5.4 First Ionization Energies, IE1

The first Ionization energy (IE1) is the energy required
to remove 1 mole of electrons from 1 mole of gaseous
atoms.

The process can be represented by an equation:

M(g) → e– + M+(g) ∆H IE1 = +ve kJ/mol

The process is endothermic as energy must be supplied to
remove the electrons which are attracted by the positively
charged nucleus.
The magnitude of the ionization energy indicates the
strength of attractive force of nucleus acting on the valence
electrons of an atom.

Figure 8.15 Periodicity of first ionization energy (IE1).

Figure 8.16 First ionization energies of the main-group elements.

IE1 across a period.

Generally, when across a period IE1 increases.

The atomic size decreases. The outer electrons are more closely
attracted to the nucleus due to increasing in Zeff and thus difficult to be
removed and required more energy.

However, there are 2 exceptions in period 2 & 3.
• There are dips in IE1 at group 13 & dips at group 16.

Example: 4Be : 1s2 2s2 Period 2

5B : 1s2 2s2 2p1

• When Be gives up electron, it must come from completely filled s
orbital, which is more stable than electron from partially p orbital in B
and harder to be removed. Therefore IE1 of Be is greater than B.

* Similar reason is used to explain the irregularities of Mg & Al (Period 3)

Example: 7N : 1s2 2s2 2p3
8O : 1s2 2s2 2p4

• When N gives up electron, it must come from partially-filled p orbital,
which is more stable than electron in O. As for O, the presence of
paired electrons in p orbital experience electron-electron repulsion,
thus easier to be removed and less energy required.

* Similar reason is used to explain the irregularities of P & S (Period 3)

Sample Problem 8.4 Ranking Elements by First Ionization
Energy

PROBLEM: Using the periodic table only, rank the elements in each of
the following sets in order of decreasing IE1:

(a) Kr, He, Ar (b) Sb, Te, Sn
(c) K, Ca, Rb (d) I, Xe, Cs

PLAN: Find each element on the periodic table. IE1 generally
decreases down a group and increases across a period.

SOLUTION:

(a) He > Ar > Kr
Kr, He, and Ar are in Group 8A. IE1 decreases down the group.

Sample Problem 8.4

SOLUTION:

(b) Te > Sb > Sn
Sb, Te, and Sn are in Period 5. IE1 increases across a period.

(c) Ca > K > Rb
K has a higher IE1 than Rb because K is higher up in Group 1A. Ca
has a higher IE1 than K because Ca is farther to the right in Period 4.

(d) Xe > I > Cs
Xe has a higher IE1 than I because Xe is farther to the right in the
same period. Cs has a lower IE1 than I because it is farther to the left
in a higher period.

Figure 8.17 The first three ionization energies of beryllium.

This big jump appears after
the outer (valence) e- have
been removed and thus,
reflecting much more energy
needed to remove the (inner)
core electron.

Outer/valence
electrons

Deducing the position of element in Periodic Table based on successive IE:
• The drastic increase of IE reflects the difficulty of removing electrons from

valence to inner/core shell.
• Much higher energy is required to remove the first core/inner electron which

comes from the inner/core shell.
• The number of electrons lost before encounter the first inner electron is the total

of valence electrons (indicates the group in Periodic Table)

Table 8.5 Successive Ionization Energies of the Elements Lithium
Through Sodium

Determination the electron configuration of an atom using successive
ionization energies
Five successive ionization energies (kJmol-1) for atom M is shown below:

IE 1 2 3 4 5
(kJmol-1) 800 1580 3230 4360 16000

Determine:
i. Electron configuration of the valence electron for M.
ii. Group number of M in the periodic table.

Solution: The large increase in value of IE5 reflects that the fifth electron
is difficult to be removed as it comes from the core. There are 4 electrons
in the valence shell/orbital and the element has a configuration of ns2 np2
and belong to group 14

Sample Problem 8.5 Identifying an Element from Its Ionization
Energies

PROBLEM: Name the Period 3 element with the following ionization
energies (in kJ/mol) and write its electron configuration:

IE1 IE2 IE3 IE4 IE5 IE6
1012 1903 2910 4956 6278 22,230

PLAN: Look for a large increase in IE, which occurs after all valence
electrons have been removed.

SOLUTION:

The largest increase occurs after IE5, that is, after the 5th
valence electron has been removed. The Period 3 element
with 5 valence electrons is phosphorus (P; Z = 15).

The complete electron configuration is 1s22s22p63s23p3.

5.5 Electron Affinity

Electron Affinity (EA) is the energy change (released or absorbed) that

occurs when 1 mole of electrons is added to 1 mole of gaseous atoms or
ions. EA measures the ability of an atom or ion to receive /gain ē.

X (g) + e- X-(g) ∆H EA1 = -ve
X- (g) + e- X2-(g) ∆H EA2 = +ve

EA2 is ALWAYS positive. Why?
Because the energy must be absorbed to overcome electrostatic
repulsions to add another e- to a negative ion

Atoms with a low EA tend to form cations.
Atoms with a high EA tend to form anions.

The trends in electron affinity are not as regular as those for atomic size
or IE.

Figure 8.18 Electron affinities of the main-group elements (in kJ/mol).