Electron Arrangement - Part 2

Chapter 9Some images Copyright © The McGraw-Hill Companies, Inc.

Brad Collins

Review Energy Levels

1s

2s

2p 2p 2p

3s

3p 3p 3p

3d4s

4p 4p 4p3d 3d 3d 3d

4d 4d 4d 4d 4d

n = 1

n = 2

n = 3

n = 4

Multi-electron

n=1, l = 0

n=2, l = 0

n=3, l = 0

n=2, l = 1

n=3, l = 1

n=3, l = 2

9.1

“Fill up” electrons in lowest energy orbitals (Aufbau principle)

H 1 electronH 1s1

He 2 electronsHe 1s2

Li 3 electronsLi 1s22s1

Be 4 electronsBe 1s22s2

B 5 electronsB 1s22s22p1

C 6 electrons

9.1

C 1s22s22p2N 1s22s22p3O 1s22s22p4F 1s22s22p5Ne 1s22s22p6

N 7 electronsO 8 electronsF 9 electronsNe 10 electrons

Electron ‘Filling’ Order• Electrons fill from lowest to highest energy

• Filling order by shell and subshell:

• 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

• No more than 2 electrons can occupy an orbital (Pauli exclusion principal)

• Within a subshell, one electron fills each orbital before the electrons pair up (Hund’s rule).

9.1

Order of orbitals (filling) in multi-electron atom

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s9.1

Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom.

1s1

principal quantum number n

angular momentum quantum number l

number of electrons in the orbital or subshell

Electron configuration of carbon: 1s2, 2s2 2p2

9.1

Electron Configuration

9.1

General Rules for Assigning Electrons to Atomic Orbitals

1. Each shell or principal level of quantum number n contains n subshells. If n = 2, there are two subshells (two values of l ) with angular momentum quantum numbers 0 and 1.

2. Each subshell of quantum number l contains 2l + 1

orbitals. For example, if l = 1, there are three p-orbitals.

3. No more than two electrons can be placed in each orbital.

4. The maximum number of electrons that an atom can have in a principal level n is equal to 2n2. 9.1

What is the electron configuration of Mg?Mg 12 electrons

1s < 2s < 2p < 3s < 3p < 4s

1s2 2s22p6 3s2 2 + 2 + 6 + 2 = 12 electrons

Abbreviated as [Ne]3s2 [Ne] =1s2 2s22p6

What are the possible quantum numbers for the last (outermost) electron in Cl?

Cl 17 electrons 1s < 2s < 2p < 3s < 3p < 4s

1s2 2s22p6 3s23p5 2 + 2 + 6 + 2 + 5 = 17 electronsLast electron added to 3p orbital

n = 3 l = 1 ml = -1, 0, or +1 ms = ½ or -½9.1

Valence Electrons• Valence electrons are electrons used by atoms to make chemical bonds.

• For a particular atom, the valence electrons are electrons with the highest n-value

• H has 1 electron in shell n = 1, so one valence electron

• 1s1!

• Li has 1 electron in shell n = 2, so one valence electron

• 1s2 2s1

• Cl has 7 electrons in shell n = 3, so seven valence electrons

• 1s2 2s22p6 3s2 3p5

9.1

Periodic Table and Electron Configuration

• Groups (columns) of representative elements in the periodic table have the same number of valence electrons

• Transition metals are predicted to have 2 valence electrons, but often do not.

• Filling d-orbitals

• Varying valence related to d-electron configurations

• Iron (Fe) configuration: 1s2, 2s2 2p6, 3s2 3p6, 4s2 3d6

• Fe2+ configuration: 1s2, 2s2 2p6, 3s2 3p6, 4s0 3d6

• Fe3+ configuration: 1s2, 2s2 2p6, 3s2 3p6, 4s0 3d5

• Note: The 5 d-electrons in Fe2+ all have parallel spins (Hund’s rule)

9.2

Electron Configurations and the Periodic Table

9.2

H HeLi Be B C N O F Ne

9.3

Electron Configurations - Alternate Approaches

Noble (rare) Gas abbreviation

Ne = 1s2 2s2 2p6

Mg = 1s2 2s2 2p6 3s1 or [Ne] 3s1

Orbital Diagram

• Uses boxes and arrows to represent orbitals and electrons

H

1s1

He

1s2

Anomalous Electron Configurations• Some elements have anomalous electron configurations (do not conform to Aufbau

principle)

• Transition metals e.g., Cr, Cu

• Predicted for Cr = [Ar] 4s2 3d

4

• Actual for Cr = [Ar] 4s1 3d

5

• Reason: Half-full d-orbital more stable, parallel spins (Hund’s rule)

• Predicted for Cu = [Ar] 4s2 3d

9

• Actual for Cu = [Ar] 4s1 3d

10

• Reason: Full d-orbital more stable

• More common for Inner-Transition metals e.g., La

• Predicted for [Xe] 6s2 4f1 Actual for [Xe] 6s2 5d1

9.3

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

Transition Metal Cations: Electron Configurations

9.3

Electron Configurations of Ions• Ions are created when neutral elements gain or lose electrons

• Electrons are gained or lost from the valence (outermost) shell

• Sodium: [Ne] 3s1, Sodium ion: [Ne], Na+

• Chlorine: [Ne] 3s23p5, Chlorine ion: [Ar], Cl–

• The tendency of ions to attain noble gas electron configurations is called the Octet Rule!

• Some elements cannot achieve a Noble Gas configuration

• Gallium: [Ar] 4s2 3d10 4s1, Gallium ion: [Ar] 3d10, Ga3+

• Called a psuedo-noble gas configuration!

• Consists of a noble gas abbreviation plus d and f subshells9.4

Practice ProblemHow many valence electrons does sulfur have?!

Sulfur is in Group 6, [Ne] 3s2 3p4

Sulfur has 6 valence electronsHow many valence electrons does nitride ion have?!

Nitrogen is in Group 6, [Ne] 3s2 3p4

Nitride ion electron configuration is [Ne] 3s2 3p6

Nitride ion has 8 valence electrons

Na+: [Ne] Al3+: [Ne] F–: 1s2 2s22p6 or [Ne]

O2–: 1s2 2s22p6 or [Ne] N3–: 1s2 2s22p6 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

Electron Configurations of Ions

9.4

Periodic Properties• Properties of elements depend on their position on the periodic

table

• Group 1 metals form 2:1 di-metal oxides, M2O

• Group 2 metals form 1:1 metal oxides, MO

• Atomic radius (CH 9.5)

• Ionization energy (CH 9.6)

• Electron affinity (CH 9.7)

• Electronegativity (CH 10.3)

Atomic Radius• Radius increases

down a group

• n increases with each period

• Trends: radius increases down and to the left

Measured from nucleus to nucleus

9.5

Ionic Radii

• Cation radius is always smaller than neutral atom

• Anion radius is always bigger than neutral atom

9.5

Ionic RadiiNa+ and Mg2+ both have 10 electrons, but Mg2+ is smaller than Na+

This is due to extra proton in Mg2+ that pulls the electrons closer to the nucleus than Na+

9.5

Ionization Energy• Ionization energy (IE) is the energy required to remove an

electron from an atom in its gaseous state.

• Ionization energy tends to increase as more electrons are removed from an atom (Table).

• Large jumps in IE mark transitions to stable electron configurations (Octet rule, Hund’s rule).

9.6

First Ionization Energy Trends

• First ionization energy is the energy to remove the first electron from an atom.

• Tends to increase across a period and up a group

• Exceptions due to more stable electron configurations, e.g., carbon (half-full p-subshell, Hund’s rule)

9.6

General Trends in First Ionization Energies(Mostly) Increasing First Ionization Energy

Incr

easi

ng F

irst I

oniz

atio

n E

nerg

y

9.6

Electron Affinity• Electron affinity (EA) is the enthalpy change that occurs when an

electron is added to an atom in the gaseous state.

• The greater the EA, the more energy an atom is willing to ‘spend’ to obtain an electron.

• Fluorine will release 328 kJ/mol to obtain an electron.

• Electron affinities are expressed as –∆H

• F(g) + 1 e– —> F–(g) ∆H = –328 kJ/mol, EA = 328 kJ/mol

• O(g) + 1 e– —> O–(g) ∆H = –141 kJ/mol, EA = 141 kj/mol

9.7

Electron Affinity Trends• First EA is the energy to add the first

electron to an atom.

• First EA ends to increase across a period and up a group

• Exceptions due to more stable electron configurations, e.g., carbon (half-full p-subshell, Hund’s rule)

9.7

General Trends in Electron Affinity(Mostly) Increasing Electron Affinity

Incr

easi

ng E

lect

ron

Affi

nity

9.7