chapter seven: atomic structure and periodicity. electrons in atoms are arranged as shells (n)...

Chapter Seven:Atomic Structure and Periodicity

Electrons in atoms are arranged as

SHELLS (n)

SUBSHELLS (l)

ORBITALS (ml)

Arrangement of Electrons in Atoms

When n = 1, then l = 0

This shell has a single orbital (1s) to which 2 electrons can be assigned.

When n = 2, then l = 0, 1

2s orbital 2 electrons three 2p orbitals 6 electrons TOTAL = 8 electrons

Electrons in Atoms

The number of shells increases with the shell value (n).

Therefore each higher shell holds more electrons. (more orbitals)

Electrons in Shells

Electrons in an atom are arranged by: (quantum number)

Principle energy levels (shells) n

angular energy levels (sub-shells) l

oriented energy levels (orbitals) ml

Within an individual orbital, there may only be two electrons differentiated by their spin.

Spin states (electrons) ms

Quantum Numbers

Based on theoretical and experimental studies of electron distributions in atoms, chemists have found there are two general rules that help predict these arrangements:

1. Electrons are assigned to subshells in order of increasing “n + l” value.

2. For two subshells with the same value of “n + l” electrons are assigned first to the subshell of lower n.

Atomic Subshell Energies & Electron Assignments

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

• Z* is the net charge experienced by a particular electron in a multi-electron atom resulting from a balance of the attractive force of the nucleus and the repulsive forces of other electrons.

• Z* increases across a period owing to incomplete screening by inner electrons

• This explains why E(4s electron) < E(3p electron)• Z* [ Z (no. inner electrons) ]

Charge felt by 2s electron in: Li Z* = 3 2 = 1 Be Z* = 4 2 = 2 B Z* = 5 2 = 3and so on!

Effective Nuclear Charge, Z*

Electron Configurations & the Periodic Table

Electron Filling Order

• 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

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

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.

Filling Rules

• Aufbau Principle:– Electrons fill the lowest energy levels first.

• Pauli Exclusion Principle:– Electrons can fill two/orbital, providing they have opposite

spins ( )

• Hund’s Rule:– Orbitals of the same energy level (p,d,f) distribute

electrons 1/orbital before adding the second.

Who are they?

1. 1s22s22p63s2

2. [Kr]5s24d105p5

3. 1s22s22p3

4. [Rn]7s1

5. [Ar]4s23d10

Who are they?

1. 1s22s22p63s2 Magnesium2. [Kr]5s24d105p5 Iodine3. 1s22s22p3 Nitrogen4. [Rn]7s1 Francium5. [Ar]4s23d10 Zinc

What’s wrong?

Mg: [Ar]3s2

Fe: 1s22s22p63s23d64s24d6

Al: [Ne] 3s3

Mistakes!

Mg: [Ar]3s2

Fe: 1s22s22p63s23p64s24d6

Al: [Ne] 3s3

Corrections

Mg: [Ne]3s2

Fe: 1s22s22p63s23p64s23d6

Al: [Ne] 3s23p1

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

• 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:

• Ions that contain the same number of electrons.

• Example: Al3+, Mg 2+, Na+, F-, O2-, N3-

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

• Diamagnetic Substances: Are NOT attracted to a magnetic field

• Paramagnetic Substances: ARE attracted to a magnetic field.

• Substances with unpaired electrons are paramagnetic.

Electron Spin & Magnetism

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

Practice: Electron Configuration

• Write the following electron configurations1. Silicon- orbital notation2. Strontium – long (spdf) notation3. Bismuth – noble gas notation4. Silver – noble gas notation5. O 2- - orbital notation6. Fe2+ - noble gas notation

Practice #2: Electron Configuration

• Write the following electron configurations1. Chromium- long (spdf) notation2. Sulfur – orbital notation3. Ag+ – noble gas notation4. Americium – noble gas notation5. Mg 2+ - orbital notation6. Krypton – noble gas notation

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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