periodic trends LW PAP

General Chemistry:

An Integrated Approach

Hill, Petrucci, 4 th Edition

Chapter 8

Electron Configurations, Atomic

Properties, and the Periodic Table

Mark P. Heitz

State University of New York at Brockport

© 2005, Prentice Hall, Inc.

Multielectron Atoms

Electrons are attracted to the nucleus while simultaneously repelling one another

In the hydrogen atom, all subshells of a principal shell are at the same energy level recall E n

= –B/ n 2

Chapter 8: Electron Configurations, Atomic

Properties and the Periodic Table

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2

Multielectron Atoms

In a multielectron atom the various subshells of a principal shell are at different energy levels, but all orbitals within a subshell are at the same energy level

The increasing energy order of subshells is generally:


Properties and the Periodic Table s < p < d < f

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Multielectron Atoms

Orbital energies are lower in multielectron atoms than in the hydrogen atom

In higher numbered principal shells of a multielectron atom, some subshells of different principal shells have nearly identical energies



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Electron Configurations

Electron configuration describes the distribution of electrons among the various orbitals in the atom

The spdf notation uses numbers to designate a principal shell and the letters to identify a subshell ; a superscript number indicates the number of electrons in a designated subshell



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Formula Mass

An orbital diagram uses boxes to represent orbitals within subshells and arrows to represent electrons:

Each box has arrows representing electron spins; opposing spins are paired together



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Rules for Electron Configurations

Electrons occupy the lowest available energy orbitals

Pauli exclusion principle – no two electrons in the same atom may have the same four quantum numbers

Orbitals hold a maximum of two electrons spins must be opposed



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For orbitals of identical energy, electrons enter empty orbitals whenever possible –

Hund’s rule

Electrons in half-filled orbitals have parallel spins



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Capacities of shells ( n ) and subshells ( l )



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Subshell filling order ...

Each subshell must be filled before moving to the next level

1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 ...

Illustration



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The Aufbau Principle

A hypothetical building up of an atom from the one that precedes it in atomic number

( Z = 1) H 1 s 1

( Z = 2) He 1 s 2

( Z = 3) Li 1 s 2 2 s 1

( Z = 3) Li 1 s 2 2 s 1 

[He]2s 1

Abbreviated electron configuration

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11

The Aufbau Principle ...

[He]2p 2

[He]2p 3

[He]2p 4

[He]2p 5

[He]2p 6

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12 Chapter 8: Electron Configurations, Atomic


Main Group and Transition

Elements

Elements in which the orbitals being filled in the aufbau process are either s or p orbitals of the outermost shell are called main group elements

“A” group designation on the periodic table

The first 20 elements are all main group elements

In transition elements , the subshell being filled in the aufbau process is in an inner principal shell

Fourth period transition elements have n = 4 as their outermost shell as the 3 d subshell fills



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Main Group and Transition

Elements

Completely filled and halffilled sublevels are more energetically favorable configurations

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Periodic Relationships

The valence shell is the outermost occupied shell

The period number = principal quantum number, n, of the electrons in the valence shell



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For main group elements the number of valence shell electrons is the same as the periodic table “A” group number



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We can deduce the general form of electron configurations directly from the periodic table



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17

Valence Electrons and Core

Electrons

Valence electrons are those with the highest principal quantum number

Sulfur has six valence electrons



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Valence Electrons and Core

Electrons

Electrons in inner shells are called core electrons

Sulfur has 10 core electrons



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Electron Configurations of Ions

Anions: gain e

– to complete the valence shell

Example:

-



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Cations: lose e

– to attain a complete valence shell

Example:

( Z = 11) Na

( Z = 11) Na +



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Cations formed from transition metals lose e

– from the highest principal energy level ( n )



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Magnetic Properties

Diamagnetism is the weak repulsion associated with paired electrons

Paramagnetism is the attraction associated with unpaired electrons

Ferromagnetism is the exceptionally strong attractions of a magnetic field for iron and a few other substances

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Periodic Trends

• Use must justify the trend across the period, you cannot simply state the trend.

• A trend is an observation, not an explanation !

• You should state the trend in your answer, but you must also go further by explaining what causes the observed trend!



24

Periodic Atomic Properties of the Elements

Periodic law states that certain sets of physical and chemical properties recur at regular intervals when the elements are arranged according to increasing atomic number

•

Consider atomic radii : distance between the nuclei of two atoms

•

The distance between the nucleus and the outer edge of the electron cloud



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Atomic Radii

• Atomic Radii decrease as atomic numbers increase in an given period (going across).

– A proton and electron are added so the effective nuclear charge increases because each proton has more of an effect than each additional electron

• As that attraction between the nucleus and electrons increases , and the atomic radius decreases

• Atomic Radii increase Sgoing down

– In going from top to bottom of a group, the valence electrons are assigned to orbitals with increasingly higher values of n (prin. Quantum number)

• The underlying electrons requires some space, so the electrons of the outer shell must be further (Your are adding energy levels )

26

Atomic Radii

•Z eff effective nuclear charge: the nuclear charge experienced by a particular electron in a multielectron atom

– Increases the attraction of the nucleus and pulls the electron cloud closer to the nucleus resulting in a smaller atomic radius

•Atomic radii of transtion metals trend a little differently

•Exceptions in atomic radii also exist in the lanthanide and actinide series because of how the f subshells are uniquely filled by electrons

27

Z eff

& Shielding

– The order in which electrons are assigned to subshells in an atom, as well as other properties are because of Z eff

– Shielding: electrons closer to the nucleus screen or shield the effect of nuclear charge on valence electrons

• the number of shielding electrons increases when you reach the end of the periodic table and go on to the next period.

• Shielding increases in steps as you start a new period or go down a group

•

Video



28

Transition Metal Atomic Trends

• From left to right across a period, the radii initially decrease, then size remains almost the same, then slightly increases toward the end.

•The small increase in atomic radii is because of the d subshell is filled with electrons and thus the ele-eletron repulsions cause the size to increase

29

Atomic Radii Properties

• The increased number of energy levels (n) increases the distance over which the nucleus must pull and therefore reduces the attraction for electrons

• Full energy levels provide shielding between the nucleus and valence electrons, so you see an increase in shielding as the level gets full

Illustration



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Ionic Radii

•

The ionic radius of each ion is the portion of the distance between the nuclei occupied by that ion

•

If the size of an atom is determined by the outermost electrons, what happens if you remove or add an electron?



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Ionic Radii

Cations are smaller than the atoms from which they are formed

– the nucleus attracts the remaining electrons more strongly

Anions are larger than the atoms from which they are formed

– the greater number of electrons repel more strongly

Think of the proton/electron ratio,

-as electrons are lost, the ratio of p+/e- increases and so the electrons are held closer vv.

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32

Isoelectronic Configurations

Isoelectronic species are elements that all have the same number of electrons

For isoelectronic species, the greater the nuclear charge, the smaller the species

Effective nuclear charge



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Atomic and Ionic Radii



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34

Ionization Energy

Ionization energy is the energy required to remove an electron from a ground state atom in the gaseous state

• to remove an electron, energy must be supplied to overcome the attraction of the nuclear charge (endothermic, always +)

Continual removal of electrons increases ionization energy greatly

Illustration

B



B + + e

–

I = 801 kJ mol

–1

B +



B +2 + e

–

I = 2427 kJ mol

–1

B +2



B +3 + e

–

I = 3660 kJ mol

–1

B +3



B +4 + e

–

I = 25,025 kJ mol

–1

B +4



B +5 + e

–

I = 32,822 kJ mol

–1



35

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Ionization energy

• First ionization energy- energy is increased with each successive removal because the electron is being removed from an increasingly positive ion

– The remaining electrons are held more tightly

– Notice the large jump at the 3 rd level for Mg.

– There is a large increase as you remove electrons from lower (inner) energy subshells

36

First Ionization Energies

Illustration



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Ionization energy

• Ionization energy increases as atomic number increases in any given period

– Z eff increases the attraction of the nucleus and holds the electrons more tightly

• Exceptions: group II to III, IE drops because the p electrons do not penetrate the nuclear region as well as s electrons so aren’t as tightly held

• Drop in IE also occurs between V & VI because of increased repulsion created by the first pairing of electrons, that is stronger than the increase in Z eff

, lowering the energy required to remove the electron



38

Ionization energy

• Ionization energy decreases as atomic number increases down a column or group

– The increased number of energy levels (n) increases the distance over which the nucleus must pull, reducing the attraction for electrons

– A full energy level provides some shielding between the nucleus and valence electrons



39

Electron Affinity

Electron affinity is the energy change that occurs when an electron is added to a gaseous atom

-How much an atom ‘likes’

Electrons (+ or -)

-the more negative it is the higher the EA

( energy is flowing out of the system)

Electron affinities are expressed as negative because the process is exothermic



Illustration

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Electronegativity

• A measure of the attraction of an atom for the pair of outer shell electrons in a covalent bond with another atom

– Pattern is same as electron affinity for same reasons

– Both are attraction nucleus has for electrons, one in forming an ion (EA) and one in forming a molecule (EN)

– Fluorine is the most electronegative. The closer

41 it is to fluorine, the more electronegative it is.

Metals, Nonmetals, and Metalloids

Metals have a small number of electrons in their valence shells and tend to form positive ions

Except for hydrogen and helium, all s -block elements are metals

All d - and f -block elements are metals



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Metals, Nonmetals, and Metalloids

Atoms of a nonmetal generally have larger numbers of electrons in their valence shell than do metals, and many tend to form negative ions

Nonmetals are all p block elements and include hydrogen and helium

Metalloids have properties of both metals and nonmetals



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Metals

• Metals react by losing electrons

– A loosely held electron will result in a more reactive metal

– This is tied directly to ionization energy

– With an increased # of energy levels ( n), comes increased distance from the nuclear attraction and thus a more loosely held electron available for reactions



44

Non-metals

• Non-metals tend to gain electrons, a strong nuclear attraction will result in a more reactive non-metal

• This means that an atom with the highest

Z eff and the least number of energy levels should be the most reactive nonmetal (F) because its nucleus exerts the strong pull



45

A Summary of Periodic Trends



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The Inert (Noble) Gases

The six noble gases, He, Ne, Ar, Kr, Xe, and Rn, rarely enter into chemical reactions

All have complete octets ...

= stability !



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“Periodic” Behavior of Elements

Flame tests: elements with low first ionization energies are excited in a flame

Atoms emit energy when electrons fall from higher to lower energy states

FlameTests



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“Periodic” Behavior of Elements halogens (Group 7A) are good oxidizing agents

When Cl

2 is bubbled in a solution containing iodide ions, chlorine oxidizes I

– to I

2



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The sBlock Metals as

Reducing Agents

2 K + 2 H

2

O



2 K + + 2 OH

–

+ H

2

Recall activity series ...

H + is reduced by these metals



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Acidic, Basic, and Amphoteric

Oxides

Acidic oxides are oxides that produce acids by reacting the oxide with water e.g., SO

3

+ H

2

O



H

2

SO

4

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Oxides

Basic oxides are oxides that produce bases by reacting with water e.g., MgO + H

2

O



Mg(OH)

2

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Oxides

Oxides that can react with either acids or bases are amphoteric oxides e.g., Al

2

O

3

Behavior of Oxides



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Summary of Concepts

• The wave-mechanical treatment of the hydrogen atom can be extended to multielectron atoms, but with two differences

• Electron configuration is the distribution of electrons in orbitals among the subshells and principal subshells

• There are two types of electron configuration notation: spdf and orbital



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Summary of Concepts

• The aufbau principle describes a process of hypothetically building up an atom from the atom of the preceding atomic number

• Elements in similar electron configurations fall in the same group of the periodic table

• An atom with all the electrons paired is diamagnetic; an atom with one or more unpaired electrons is paramagnetic

• Certain atomic properties, such as atomic radius, ionic radius, ionization energy, and electron affinity, vary periodically with increasing atomic number


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Summary of Concepts

• The regions of the periodic table ascribed to metals, nonmetals, metalloids, and the noble gases are related to the value of atomic properties



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