Periodic Relationships Among

the Elements

Chapter 8

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2

When the Elements Were Discovered

3

ns

1

ns

2

ns

2n

p1

ns

2n

p2

ns

2n

p3

ns

2n

p4

ns

2n

p5

ns

2n

p6

d1

d5

d1

0

4f

5f

Ground State Electron Configurations of the Elements

4

Classification of the Elements

5

Electron Configurations of Cations and Anions

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s23p1 Al3+ [Ne]

Atoms lose electrons so that

cation has a noble-gas outer

electron configuration.

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Atoms gain electrons

so that anion has a

noble-gas outer

electron configuration.

Of Representative Elements

6

+1

+2

+3 -1-2-3

Cations and Anions Of Representative Elements

7

Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne]

O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne]

Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

What neutral atom is isoelectronic with H- ?

H-: 1s2 same electron configuration as He

Isoelectronic: have the same number of electrons, and

hence the same ground-state electron configuration

8

Electron Configurations of Cations of Transition Metals

When a cation is formed from an atom of a transition metal,

electrons are always removed first from the ns orbital and

then from the (n – 1)d orbitals.

Fe: [Ar]4s23d6

Fe2+: [Ar]4s03d6 or [Ar]3d6

Fe3+: [Ar]4s03d5 or [Ar]3d5

Mn: [Ar]4s23d5

Mn2+: [Ar]4s03d5 or [Ar]3d5

9

EXAMPLE 8.1

An atom of a certain element has 15 electrons. Without consulting a periodic table,

answer the following questions: (a) What is the ground-state electron configuration of

the element? (b) How should the element be classified? (c) Is the element diamagnetic

or paramagnetic?

Strategy (a) We refer to the building-up principle discussed in Section 7.9 and start

writing the electron configuration with principal quantum number n 5 1 and continuing

upward until all the electrons are accounted for. (b) What are the electron configuration

characteristics of representative elements? transition elements? noble gases? (c) Examine

the pairing scheme of the electrons in the outermost shell. What determines whether an

element is diamagnetic or paramagnetic?

Solution (a) We know that for n 5 1 we have a 1 s orbital (2 electrons); for n 5 2 we

have a 2 s orbital (2 electrons) and three 2 p orbitals (6 electrons); for n 5 3 we have a

3 s orbital (2 electrons). The number of electrons left is 15 2 12 5 3 and these three

electrons are placed in the 3 p orbitals. The electron configuration is 1 s 2 2 s 2 2 p 6 3 s 2 3 p 3 .

(b) Because the 3 p subshell is not completely filled, this is a representative element.

Based on the information given, we cannot say whether it is a metal, a nonmetal, or

a metalloid.

(c) According to Hund’s rule, the three electrons in the 3 p orbitals have parallel spins

(three unpaired electrons). Therefore, the element is paramagnetic.

Check For (b), note that a transition metal possesses an incompletely filled d subshell

and a noble gas has a completely filled outer shell. For (c), recall that if the atoms of

an element contain an odd number of electrons, then the element must be paramagnetic.

Practice Exercise An atom of a certain element has 20 electrons. (a) Write the

ground-state electron configuration of the element, (b) classify the element, (c) determine

whether the element is diamagnetic or paramagnetic.

© 2009, Prentice-Hall, Inc.

The shielding Effect and Effective

Nuclear Chargeتأثير الحجب في الذرات متعددة اإللكترونات وشحنة النواة المؤثرة

• In a many-electron atom,

electrons are both

attracted to the nucleus

and repelled by other

electrons.

• The nuclear charge that

an electron experiences

depends on both factors.

© 2009, Prentice-Hall, Inc.

The shielding Effect and Effective

Nuclear Charge

The effective nuclear

charge, Zeff, is found this

way:

Zeff = Z − S

where Z is the atomic

number and S is a

screening constant الجزء من شحنة )

(النواة المحجوب عن جذب الكترونات التكافؤ ,

usually close to the

number of inner

electrons.

12

Effective nuclear charge (Zeff) is the “positive charge” felt

by an electron.

Na

Mg

Al

Si

11

12

13

14

10

10

10

10

1

2

3

4

186

160

143

132

ZeffCoreZ Radius (pm)

Zeff = Z - s 0 < s < Z (s = shielding constant)

Zeff Z – number of inner or core electrons

13

Effective Nuclear Charge (Zeff)

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

What Is the Size of an Atom?

The bonding atomic

radius is defined as

one-half of the

distance between

covalently bonded

nuclei.

15

Atomic Radii

metallic radius covalent radius

16

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Sizes of Atoms

Bonding atomic

radius tends to…

…decrease from left to

right across a row

(due to increasing Zeff).

…increase from top to

bottom of a column

(due to increasing value

of n).

18

19

Trends in Atomic Radii

20

Comparison of Atomic Radii with Ionic Radii

21

Cation is always smaller than atom from

which it is formed.

Anion is always larger than atom from

which it is formed.

22

The Radii (in pm) of Ions of Familiar Elements

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Sizes of Ions

• Ionic size depends

upon:

– The nuclear

charge.

– The number of

electrons.

– The orbitals in

which electrons

reside.

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Sizes of Ions

• Cations are

smaller than their

parent atoms.

– The outermost

electron is

removed and

repulsions

between electrons

are reduced.

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Sizes of Ions

• Anions are larger

than their parent

atoms.

– Electrons are

added and

repulsions

between electrons

are increased.

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Sizes of Ions

• Ions increase in size

as you go down a

column.

– This is due to

increasing value of n.

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Sizes of Ions

• In an isoelectronic series, ions have the same

number of electrons.

• Ionic size decreases with an increasing

nuclear charge.

Periodic

Properties

of the

Elements

EXAMPLE 8.2

Referring to a periodic table, arrange the following atoms in order of increasing atomic

radius: P, Si, N.

Strategy What are the trends in atomic radii in a periodic group and in a particular

period? Which of the preceding elements are in the same group? in the same period?

Solution From Figure 8.1 we see that N and P are in the same group (Group 5A).

Therefore, the radius of N is smaller than that of P (atomic radius increases as we

go down a group). Both Si and P are in the third period, and Si is to the left of P.

Therefore, the radius of P is smaller than that of Si (atomic radius decreases as we

move from left to right across a period). Thus, the order of increasing radius is

N , P , Si .

Practice Exercise Arrange the following atoms in order of decreasing radius: C,

Li, Be.

© 2009, Prentice-Hall, Inc.

29

Chemistry in Action: The 3rd Liquid Element?L

iqu

id?

117 elements, 2 are liquids at 250C – Br2 and Hg

223Fr, t1/2 = 21 minutes

30

Ionization energy is the minimum energy (kJ/mol) required

to remove an electron from a gaseous atom in its ground

state.

I1 + X (g) X+(g) + e-

I2 + X+(g) X2+

(g) + e-

I3 + X2+(g) X3+

(g) + e-

I1 first ionization energy

I2 second ionization energy

I3 third ionization energy

I1 < I2 < I3

31

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Trends in First Ionization Energies

• Generally, as one

goes across a row, it

gets harder to

remove an electron.

– As you go from left to

right, Zeff increases.

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Trends in First Ionization Energies

However, there are

two apparent

discontinuities in this

trend.

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Trends in First Ionization Energies

• The first occurs between Groups IIA and IIIA.

• In this case the electron is removed from a p-orbital rather than an s-orbital.– The electron removed

is farther from nucleus.

– There is also a small amount of repulsion by the s electrons.

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Trends in First Ionization Energies

• The second occurs

between Groups VA

and VIA.

– The electron removed

comes from doubly

occupied orbital.

– Repulsion from the

other electron in the

orbital aids in its

removal.

36

General Trends in First Ionization Energies

Increasing First Ionization Energy

Incre

asin

g F

irst Io

niz

atio

n E

ne

rgy

37

Electron affinity is the negative of the energy change that

occurs when an electron is accepted by an atom in the

gaseous state to form an anion.

X (g) + e- X-(g)

F (g) + e- F-(g)

O (g) + e- O-(g)

DH = -328 kJ/mol EA = +328 kJ/mol

DH = -141 kJ/mol EA = +141 kJ/mol

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Trends in Electron Affinity

In general, electron

affinity becomes

more exothermic as

you go from left to

right across a row.

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Trends in Electron Affinity

There are

again,

however, two

discontinuities

in this trend.

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Trends in Electron Affinity

• The first occurs

between Groups IA

and IIA.

– The added electron

must go in a p-orbital,

not an s-orbital.

– The electron is farther

from nucleus and

feels repulsion from

the s-electrons.

Periodic

Properties

of the

Elements

© 2009, Prentice-Hall, Inc.

Trends in Electron Affinity

• The second occurs

between Groups IVA

and VA.

– Group VA has no

empty orbitals.

– The extra electron

must go into an

already occupied

orbital, creating

repulsion.

42

43

Variation of Electron Affinity With Atomic Number (H – Ba)

44

45

Diagonal Relationships on the Periodic Table

46

Group 1A Elements (ns1, n 2)

M M+1 + 1e-

2M(s) + 2H2O(l) 2MOH(aq) + H2(g)

4M(s) + O2(g) 2M2O(s)

Incre

asin

g r

eactivity

47

Group 1A Elements (ns1, n 2)

48

Group 2A Elements (ns2, n 2)

M M+2 + 2e-

Be(s) + 2H2O(l) No ReactionIn

cre

asin

g r

eactivity

Mg(s) + 2H2O(g) Mg(OH)2(aq) + H2(g)

M(s) + 2H2O(l) M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba

49

Group 2A Elements (ns2, n 2)

50

Group 3A Elements (ns2np1, n 2)

4Al(s) + 3O2(g) 2Al2O3(s)

2Al(s) + 6H+(aq) 2Al3+

(aq) + 3H2(g)

51

Group 3A Elements (ns2np1, n 2)

52

Group 4A Elements (ns2np2, n 2)

Sn(s) + 2H+(aq) Sn2+

(aq) + H2 (g)

Pb(s) + 2H+(aq) Pb2+

(aq) + H2 (g)

53

Group 4A Elements (ns2np2, n 2)

54

Group 5A Elements (ns2np3, n 2)

N2O5(s) + H2O(l) 2HNO3(aq)

P4O10(s) + 6H2O(l) 4H3PO4(aq)

55

Group 5A Elements (ns2np3, n 2)

56

Group 6A Elements (ns2np4, n 2)

SO3(g) + H2O(l) H2SO4(aq)

57

Group 6A Elements (ns2np4, n 2)

58

Group 7A Elements (ns2np5, n 2)

X + 1e- X-1

X2(g) + H2(g) 2HX(g)

Incre

asin

g r

eactivity

59

Group 7A Elements (ns2np5, n 2)

60

Group 8A Elements (ns2np6, n 2)

Completely filled ns and np subshells.

Highest ionization energy of all elements.

No tendency to accept extra electrons.

61

Compounds of the Noble Gases

A number of xenon compounds XeF4, XeO3,

XeO4, XeOF4 exist.

A few krypton compounds (KrF2, for example)

have been prepared.

62

The metals in these two groups have similar outer

electron configurations, with one electron in the

outermost s orbital.

Chemical properties are quite different due to difference

in the ionization energy.

Comparison of Group 1A and 1B

Lower I1, more reactive

63

Properties of Oxides Across a Period

basic acidic

64

Chemistry in Action: Discovery of the Noble Gases

Sir William Ramsay