1

Electron configurations follow the order of sublevels on

the periodic table.

The periodic table consists of sublevel blocks

arranged in order of increasing energy.

• Groups 1A(1)-2A(2) = s level

• Groups 3A(13)-8A(18) = p level

• Groups 3B(3) to 2B(12) = d level

• Lanthanides/Actinides = f level

2

To write an electron

configuration using

Sublevel blocks,

• locate the element on the periodic table

• starting with H in 1s,write

each sublevel block in order

going from left to right

across each period

• write the number of electrons

in each block

3

Solution

Period 1 1s block 1s2

Period 2 2s → 2p blocks 2s2 2p6

Period 3 3s → 3p blocks 3s23p2 (Si)

Writing all the sublevel blocks in order gives

1s22s22p63s23p2

4

A. 1s22p5

B. 1s22s22p6

C. 1s22s22p3

D. 1s22s22p4

5

1s22p5

1s22s2

2p6

1s22s2

2p3

1s22s2

2p4

25% 25%25%25%

A. 1s22s22p63s23p64s2

B. 1s22s22p63s23p64s1

C. 1s22s22p63s23p63d2

D. 1s22s22p63s23p54s23d1

6

1s22s2

2p63s23p64s2

1s22s2

2p63s23p64s1

1s22s2

2p63s23p63d2

1s22s2

2p63s23p54s2

3d1

25% 25%25%25%

• For 3d transition metals, 3d has lower energy after the

electrons is filled. 3d and 4s are close in energy

1s 2s 2p 3s 3p 4s 3d

Ar 1s2 2s2 2p6 3s2 3p6

K 1s2 2s2 2p6 3s2 3p6 4s1

Ca 1s2 2s2 2p6 3s2 3p6 4s2

Sc 1s2 2s2 2p6 3s2 3p6 3d14s2

Ti 1s2 2s2 2p6 3s2 3p6 3d24s2

7

8

9

A. 1s22s2 2p63s2 3p6 4s2

B. 1s22s2 2p63s2 3p6 4s2 3d5

C. 1s22s2 2p63s2 3p6 4s2 3d3

D. 1s22s2 2p63s2 3p6 4s2 3d2

10

1s22s2

2p63s2

3p6 4

s2

1s22s2

2p63s2

3p6 4

s2 3

d5

1s22s2

2p63s2

3p6 4

s2 3

d3

1s22s2

2p63s2

3p6 4

s2 3

d2

25% 25%25%25%

Solution

Period 1 1s block 1s2

Period 2 2s → 2p blocks 2s2 2p6

Period 3 3s → 3p blocks 3s2 3p6

Period 4 4s → 3d blocks 4s2 3d2 (at Ti)

Writing all the sublevel blocks in order gives

1s2 2s2 2p63s2 3p6 4s2 3d2

11

A. Na

B. S

C. K

D. Ca

12

Na S K Ca

0% 0%0%0%

A. 3p64s2

B. 4s24d7

C. 4s23d7

13

3p64s2

4s24d7

4s23d7

100%

0%0%

The last filled principal energy level is called the

valence level, or valence shell which is the

outermost sublevels that are highest in energy.

The valence electrons

• occupy orbitals in the valence level All the other

electrons are called core electrons, or inner electrons.

• determine the chemical properties of an element

• are related to the group number of the element

Example: Phosphorus has 5 valence electrons

5 valence electrons

P Group 5A(15) 1s22s22p6 3s23p3

14

All the elements in a group have the same number

of valence electrons.

Example:

Be 1s2 2s2

Mg 1s2 2s2 2p6 3s2

Ca 1s2 2s2 2p6 3s2 3p6 4s2

Sr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2

Can you see a pattern?

Elements in Group 2A (2) have two (2) valence electrons.

15

7-16

State the number of valence electrons for each of following:

A. oxygen

1) 4 2) 6 3) 8

B. aluminum

1) 13 2) 3 3) 1

C. chlorine

1) 2 2) 5 3) 7

Solution:

A. 2) 6 B. 2) 3 C. 3) 7

17

State the number of valence electrons for each.

A. 1s22s22p63s23p3

B. 1s22s22p63s23p64s23d104p4

C. 1s22s22p5

Solution:

A. 5 B. 6 C. 7

18

7-19

Noble gas-core abbreviated electron configurations are

often used for elements with many electrons.

Notice that iron’s electron configuration starts out with

argon’s electron configuration, but ends differently:

Fe 1s22s22p63s23p64s23d6

Ar 1s22s22p63s23p6

We use the symbol [Ar] to represent argon’s electron

configuration:

Fe [Ar] 4s23d6

argon core + valence electrons

7-20

In atoms, the number of electrons is equal to the

number of protons, which is the atomic number.

In ions, the number of electrons does not equal the

atomic number. We must add or subtract electrons,

depending on whether the ion is an anion or cation.

7-21

An example problem:

Write the electron configuration for Na+:

Na+ has a positive charge of 1; therefore,

we need to subtract 1 electron from the

total number of electrons, 11. Na+ has 10

electrons and is isoelectronic with Ne.

1s22s22p6

7-22

Write the electron configuration in spdf notation and noble gas-core abbreviated electronic configuration for the following ions.

1. Br-

2. N3-

3. K+

4. Sr2+

5. S2-

6. Ni2+

7-23

Write the electron configuration in spdf noation and

noble –gas core abbreviated for the following ions.

1. Br- 1s22s22p63s23p64s23d104p6 [Kr]

2. N3- 1s22s22p6 [Ne]

3. K+ 1s22s22p63s23p6 [Ar]

4. Sr2+ 1s22s22p63s23p64s23d104p6 [Kr]

5. S2- 1s22s22p63s23p6 [Ar]

6. Ni2+ 1s22s22p63s23p6 3d8 [Ar] 3d8

7-24

Electron configurations are related to the

following properties:

Relative reactivity

Atomic radii (atomic size)

Ionization Energy (tendency to lose

electrons)

7-25

From the activity series we see that the most active metals are those in Groups IA and IIA.

The more active the metal, the more easily it loses electrons to form ions.

Figure 7.24

7-26

The alkali metals form +1 ions since they all have 1 valence electron.

Oxides of the alkali metals:

Li2O, Na2O, K2O, Rb2O, Cs2O

The alkaline earth metals form +2 ions since they all have 2 valence electrons.

Oxides of the alkaline earth metals:

BeO, MgO, CaO, SrO, BaO

Ionization energy

• is the energy it takes to remove a electron from an atom

or ion in the gas state.

27

the ionization energy

decreases in going

down a group.

28

In general, for Main group elements:

the ionization energy increase in going left to right across

a in period. Except, IE(group 3A) < IE(group 2A) and

IE(group 6A) < IE(group 5A)

Select the element in each pair with the higher

ionization energy.

A. Li or K

B. K or Br

C. P or Cl

Solution

A. Li

B. Br

C. Cl

29

7-30

What do these values tell us about the stability of

core (inner) electrons?

31

Atomic size is usually reported atomic radius

covalent radius: half the distance between

identical nuclei of the neighboring atoms of the

same element in a molecules

Metallic radius: is based

on the distance of

closest approach of

adjacent atoms in a

solid metal.

The atomic radius in

general increases

• going down each group

as the number of energy

levels increases

• Except : Al > Ga

32

The atomic radius in general decreases

• going from left to right across a period

• because electrons are held closer to the nucleus by the

increasingly greater charge of the nucleus.

33

Select the element in each pair with the larger

atomic radius.

A. Li or K

B. K or Br

C. P or Cl

34

Select the element in each pair with the larger

atomic radius.

A. K is larger than Li

B. K is larger than Br

C. P is larger than Cl

35

A positive ion

• has lost its valence

electrons

• is smaller than the

corresponding neutral atom

• Example: Group 1A and

metals ions about half the

size of its corresponding

neutral atoms

36

The sodium ion Na+

• forms when the Na atom loses one electron from

the third energy level

• is smaller than a Na atom

37

A negative ion

• increases the number of

valence electrons

• is larger than the

corresponding neutral atom

• Example: Group 7A and

nonmetal ions about twice the

size of its corresponding

neutral atoms

38

The fluoride ion F-

• forms when a valence electron is added

• has increased repulsions due to the added valence

electron

• is larger than a F atom

39

7-40

For any isoelectronic (same number of electrona) series,

as the number of protons increases, the ion size

decreases.

Figure from p. 286

7-41

The general trend for ionic size (or radius) is for ionic size to increase from top to bottom in the periodic table.

For cations, as the charge increases, the ionic size decreases.

For anions, as the

charge increases

(becomes more

negative), the ionic

size increases.

1. Which is larger in each of the following?

A. K or K+

B. Al or Al3+

C. S2- or S

2. Which is smaller in each of the following?

A. N3- or N

B. Cl or Cl-

C. Sr2+ or Sr

42

1. Which is larger in each of the following?

A. K > K+

B. Al > Al3+

C. S2- > S

2. Which is smaller in each of the following?

A. N < N3-

B. Cl < Cl-

C. Sr2+ < Sr

43

Ability of an atom to attract bonding electrons

Proposed by Linus Pauling in the early 1930’s

A difference in electronegativity between the atoms in a covalent bond results in:

A polar covalent bond

Increased ionic character

The greater the difference in electronegativity, the greater the ionic character and the more polar the bond that joins the atoms.

Decreased bond length and increased bond strength

No difference in electronegativity between atoms in a covalent bond results in a nonpolar covalent bond.

8-44

8-45

Electronegativity Values

Figure 8.5

The difference in electronegativity between metals

and nonmetals is so large, that the electrons are

transferred, not shared.

The greater the electronegativity difference, the

more polar the bond.

Si-F > N-F> O-F >F-F

What partial charges go

on each atom?

8-46

Figure 8.6