Chapter Seven: Atomic Structure and Periodicity. In a Hydrogen atom (1–electron) the orbitals of a subshell are equal in energy (degenerate) 1s 2s3s4s2p3p3d.

Chapter Seven:Atomic Structure and Periodicity

In a Hydrogen atom (1–electron) the orbitals of a subshell are equal in energy (degenerate)

1s

2s

3s

4s

2p

3p 3d

Ene

rgy

Single-Electron Atom Energy Levels

1s

2s

3s

4s

2p

3p

3d

E =

n +

l

Screening results in the 4s-orbital having a lower energy that that of the 3d-orbital.

4s: n + l = 4 + 0 = 4vs.

3d: n + l = 3 + 2 = 5

Multi-Electron Atom Energy Levels

• In the H-atom, all subshells of same n have same energy.

• In a multi-electron atom:1. subshells increase in energy as

value of n + l increases.2. for subshells of same n + I, the

subshell with lower n is lower in energy.

Assigning Electrons to Subshells

Aufbau Principle: Lower energy orbitals fill first.

Hund’s Rule: Degenerate orbitals (those of the same energy) are filled with electrons until all are half filled before pairing up of electrons can occur.

Pauli exclusion principle: Individual orbitals only hold two electrons, and each should have different spin.

“s” orbitals can hold 2 electrons

“p” orbitals hold up to 6 electrons

“d” orbitals can hold up to 10 electrons

Orbital Filling Rules:

Hund’s Rule. Degenerate orbitals are filled with electrons until all are half-filled before pairing up of electrons can occur.

Consider a set of 2p orbitals:

2p

Electrons fill in this manner

Orbital Filling: The “Hund’s Rule”

“Pauli exclusion principle”Individual orbitals only hold two electrons, and each should have opposite spin.

Consider a set of 2p orbitals:

2p

Electrons fill in this manner

= spin up = spin down*the convention

is to write up, then down.

Orbital Filling: The “Pauli Principle”

Electrons fill the orbitals from lowest to highest energy.

The electron configuration of an atom is the total sum of the electrons from lowest to highest shell.Example:

Nitrogen: N has an atomic number of 7, therefore 7 electrons

1s 2s 2p

1s2 2s2 2p3

Orbitals

Electron Configuration (spdf) notation:

Electron Configuration: Orbital Box Notation

Atomic Electron Configurations

Electron Configurations & the Periodic Table

Electron Filling Order

Group 4AAtomic number = 66 total electrons1s2 2s2 2p2

1s 2s 2p

Carbon

Orbital box notation:

1s 2s 2p 3s 3p

This corresponds to the energy level diagram:

1s

2s

2p

3s

3p

Electron Configuration in the 3rd Period

Orbital box notation:

1s 2s 2p 3s 3p

Aluminum: Al (13 electrons)

1s

2s

2p

3s

3p

1s2 2s2 2p6 3s2

spdf Electron Configuration

3p1

Electron Configuration in the 3rd Period

3s 3p

[Ne]

The electron configuration of an element can be represented as a function of the core electrons in terms of a noble gas and the valence electrons.

Orbital Box Notation

Noble gas Notation

[Ne] 3s23p2

Full electron configuration spdf notation

1s22s22p63s23p2

Noble Gas Notation

• The innermost electrons (core) can be represented by the full shell of noble gas electron configuration:1s22s2 = [He], 1s22s22p6 = [Ne], 1s22s22p63s23p6 = [Ar]...

• The outermost electrons are referred to as the “Valence” electrons”.

Element Full Electron Config. Noble Gas Notation

Mg 1s2 2s2 2p6 3s2 [Ne] 3s2

Core or Noble Gas Notation

All 4th period and beyond d-block elements have the electron configuration [Ar] nsx (n - 1)dy Where n is the period and x, y are particular to the element.

CopperIronChromium

Transition Metal

Electron Configurations are written by shell even though the electrons fill by the periodic table:

Ni:

1s2 2s2 2p6 3s2 3p6 4s2 3d8

last electron to fill: 3d8

electron configuration by filling:

1s2 2s2 2p6 3s2 3p6 4s23d8

electron configuration by shell: (write this way)

Transition Elements

Ions with UNPAIRED ELECTRONS are PARAMAGNETIC (attracted to a magnetic field).

Ions without UNPAIRED ELECTRONS are DIAMAGNETIC (not attracted to a magnetic field).

Fe3+ ions in Fe2O3 have 5 unpaired electrons.This makes the sample paramagnetic.

Electron Configurations of Ions

f-block elements: These elements have the configuration [core] nsx (n - 1)dy (n - 2)fz Where n is the period and x, y & z are particular to the element.

Cerium:[Xe] 6s2 5d1 4f1

Uranium:[Rn] 7s2 6d1 5f3

Lanthanides & Actinides

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Some Anomalies

Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.

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Some Anomalies

For instance, the electron configuration for copper is[Ar] 4s1 3d10

rather than the expected[Ar] 4s2 3d9.

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Some Anomalies

• This occurs because the 4s and 3d orbitals are very close in energy.

• These anomalies occur in f-block atoms, as well.

Valence Electrons:• Electrons at the highest energy level. • Electrons available to be gained/lost/shared in a chemical

reaction• Valence electron configuration: nsxpx (total of 8 valence electrons)

• Representation of valence electrons in Lewis dot notation:

Formation of ions & Valence electrons

• Ions are formed when atoms either:1. Cation (+): Give up (lose) electrons2. Anion (-): Gain electrons Cation=

• Overall goal: stable noble gas notation

• Elements with < 4 valence electrons- form cations• Elements with > 4 valence electrons for anions.• Non-metals with 4 valence electrons- do not form ions• Noble gases (8 valence electrons) – are unreactive

Ions & Electron Configuration

• Atoms or groups of atoms that carry a charge• Cations- positive charge

– Formed when an atom loses electron(s)– Na Na+ + e-– Na+ : 1s22s22p6 (10 e-) = [Ne]

• Anions –negative charge– Formed when an atom gain electrons(s)– F + e- F-– Cl- : 1s22s22p6 (10 e-) = [Ne]

Isoelectronic series:

Transition Metals & Ions

• Transition metals: multi-valent ions

(more than one possible charge)

• Electrons are removed from the highest quantum number first:

• Example: Cu+ & Cu 2+

Cu+: [Ar] 3d10

Cu2+ : [Ar] 3d9

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Development of Periodic Table

• Elements in the same group generally have similar chemical properties.

• Physical properties are not necessarily similar, however.

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Periodic Trends

• In this chapter, we will rationalize observed trends in– Sizes of atoms and ions.– Ionization energy.– Electron affinity.

• Atomic and ionic size• Ionization energy• Electron affinity

Higher effective nuclear charge

Electrons held more tightly

Larger orbitals.Electrons held less

tightly.

General Periodic Trends

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Effective Nuclear Charge

• 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.

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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, usually close to the number of inner electrons.

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What Is the Size of an Atom?

The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

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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).

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Ionization Energy

• The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.– The first ionization energy is that energy

required to remove first electron.– The second ionization energy is that energy

required to remove second electron, etc.

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Ionization Energy• It requires more energy to remove each successive

electron.• When all valence electrons have been removed, the

ionization energy takes a quantum leap.

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Trends in First Ionization Energies

• As one goes down a column, less energy is required to remove the first electron.– For atoms in the same

group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

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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.

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Trends in First Ionization Energies

However, there are two apparent discontinuities in this trend.

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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.

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Electron Affinity

Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom:

Cl + e− Cl−

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Trends in Electron Affinity

In general, electron affinity becomes more exothermic as you go from left to right across a row.

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Trends in Electron Affinity

There are again, however, two discontinuities in this trend.

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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.

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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.

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Sizes of Ions

• Ionic size depends upon:– The nuclear charge.– The number of

electrons.– The orbitals in which

electrons reside.

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Sizes of Ions

• Cations are smaller than their parent atoms.– The outermost

electron is removed and repulsions between electrons are reduced.

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Sizes of Ions

• Anions are larger than their parent atoms.– Electrons are added

and repulsions between electrons are increased.

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Sizes of Ions

• Ions increase in size as you go down a column.– This is due to

increasing value of n.

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Sizes of Ions

• In an isoelectronic series, ions have the same number of electrons.

• Ionic size decreases with an increasing nuclear charge.

Problem:Rank the following ions in order of decreasing size?

Na+, N3-, Mg2+, F– O2–

Problem:Rank the following ions in order of decreasing size?

Na+, N3-, Mg2+, F– O2–

Ion # of protons # of electrons ratio of e/p

Na+

N3-

Mg2+

F–

O2–

Problem:Rank the following ions in order of decreasing size?

Na+, N3-, Mg2+, F– O2–

Ion # of protons # of electrons ratio of e/p

Na+

N3-

Mg2+

F–

O2–

11

7

12

9

8

Problem:Rank the following ions in order of decreasing size?

Na+, N3-, Mg2+, F– O2–

Ion # of protons # of electrons ratio of e/p

Na+

N3-

Mg2+

F–

O2–

11

7

12

9

8

10

10

10

10

10

Problem:Rank the following ions in order of decreasing size?

Na+, N3-, Mg2+, F– O2–

Ion # of protons # of electrons ratio of e/p

Na+

N3-

Mg2+

F–

O2–

11

7

12

9

8

10

10

10

10

10

0.909

1.43

0.833

1.11

1.25

Na+

N3-

Mg2+

F–

O2–

Ion

0.909

1.43

0.833

1.11

1.25

e/p ratio Since N3- has the highest ratio of electrons to protons, it must have the largest radius.

Na+

N3-

Mg2+

F–

O2–

Ion

0.909

1.43

0.833

1.11

1.25

e/p ratio Since N3- has the highest ratio of electrons to protons, it must have the largest radius.

Since Mg2+ has the lowest, it must have the smallest radius.

The rest can be ranked by ratio.

Na+

N3-

Mg2+

F–

O2–

Ion

0.909

1.43

0.833

1.11

1.25

e/p ratio Since N3- has the highest ratio of electrons to protons, it must have the largest radius.

Since Mg2+ has the lowest, it must have the smallest radius.

The rest can be ranked by ratio.

N3- > O2– > F– > Na+ > Mg2+

Decreasing size

Notice that they all have 10 electrons: They are isoelectronic (same electron configuration) as Ne.

Moving through the periodic table: Atomic radii Ionization

EnergyElectron Affinity

Down a group Increase Decrease Becomes less exothermic

Across a Period Decrease IncreaseBecomes

more exothermic

Summary of Periodic Trends