Periodic Relationships Among the Elements

Chapter 8

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8.2

ns1

ns2

ns2

np1

ns2

np2

ns2

np3

ns2

np4

ns2

np5

ns2

np6

d1

d5 d10

4f

5f

Ground State Electron Configurations of the Elements

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

8.2

+1

+2

+3 -1-2-3

Cations and Anions Of Representative Elements

8.2

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

8.2

Electron Configurations of Cations of Transition Metals

8.2

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

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 - 0 < < Z ( = shielding constant)

Zeff Z – number of inner or core electrons

8.3

Effective Nuclear Charge (Zeff)

8.3

increasing Zeff

incr

easi

ng Z

eff

8.3

8.3

Atomic Radii

8.3

8.3

Comparison of Atomic Radii with Ionic Radii

Cation is always smaller than atom from which it is formed.Anion is always larger than atom from which it is formed.

8.3

8.3

The Radii (in pm) of Ions of Familiar Elements

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

8.4

I1 < I2 < I3

8.4

Filled n=1 shell

Filled n=2 shell

Filled n=3 shell

Filled n=4 shellFilled n=5 shell

8.4

Variation of the First Ionization Energy with Atomic Number

General Trend in First Ionization Energies

8.4

Increasing First Ionization Energy

Incr

ea

sing

Firs

t Io

niz

atio

n E

ner

gy

Ionization Energy Trends DOWN a Group

Ionization Energy decreases as you move down a group because the electrons that are being lost are in successively higher energy levels – further from the nucleus – requiring less energy to remove them.

(Note – although the nuclear charge is greater (more protons), there is more shielding because of the additional electrons… so the overall effective nuclear charge doesn’t really change)

Ionization Energy Trends DOWN a Group

Which has a higher ionization energy, Na or Rb? Explain why.

Which has a higher ionization energy, Br or Cl? Explain why.

Ionization Energy Trends ACROSS a Period

Ionization Energy increases as you move across a period because there is an increase in the charge on the nucleus, which increases the effective nuclear charge – requiring more energy to remove electrons from the same energy level.

Note: The increase is not totally smooth. Between group 2 and 13 there is a ‘dip’ (decrease) in ionization energy because the electron removed from the element in group 13 is in a p-subshell which is slightly higher in energy than the s-subshells of group 2 and 1. This electron is slightly higher in energy already, so it takes less energy to remove it than it does to remove an electron from the s-subshell.

A similar glitch occurs between groups 15 and 16. In group 15 the element has only 1 electron in each p-subshell. In group 16, one of the p-subshells has a pair of electrons. The resulting greater electron-electron repulsion counteracts the effect of the increased nuclear charge... which lowers the activation energy.

Ionization Energy Trends ACROSS a Period

Which has a lower ionization energy, Na or Cl? Explain why.

Which has a higher ionization energy, B or Be? Explain why.

Which has a higher ionization energy, O or N? Explain why.

Electronegativity Trends – Why?

Down a Group – EN decreases• Valence electrons further from nucleus• Effective nuclear charge isn’t changing• Atom’s ability to attract an electron further away from the

nucleus is diminishingAcross a Period – EN increases• Charge on nucleus increases (more protons)• Electrons in valence shell are more attracted to nucleus –

closer • Atom’s ability to attract an electron is increased due to

closer distance to nucleus and increase effective nuclear charge.

Electronegativity Trends – Why?

Which has a higher electronegativity, Na or Rb? Explain why.

Which has a higher electronegativity, Br or Cl? Explain why.

Melting Point Trends – Why?Down a Metal Group – MPt decreases slightly• Valence electrons further from nucleus• Effective nuclear charge isn’t changing• Atom’s ability to attract an electron further away from the nucleus is

diminishing

Across a Period – MPt increases… then decreases (MPt reflects strength of forces between particles in solid vs liquid states.

• Gps 1&2 - Strength of metallic bond increases with inc nuclear charge, inc # mobile valence electrons, atomic radius decreases

• Gps 14 – giant covalent structures… v. strong covalent bonding in all directions

• Gps 15-16-17 – molecular structures… nonpolar… only weak van der Waals forces… mpt depends on mass and size of molecules

• Gp 18 – single atoms, v wk vdWaals… very low mpt

Melting Point Trends – Why?

Which has a higher melting point, Na or Cl2? Explain why.

Which has a higher melting point, Na or Rb? Explain why.

Which has a higher melting point, C or O2? Explain why.

Periodic TrendsIonization Energy, Electronegativity, Atomic Radius, Ionic Radius, Melting Point