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- Chapter 7 -

Periodic Properties

of the Elements

Summary

7.1 – Development of the periodic table

7.2 – Effective nuclear charge

7.3 – Size of atoms and ions

7.4 – Ionization energy

7.5 – Electron affinities

7.6 – Metals, Nonmetals, and Metalloids

7.7 – Group trends for the active metals

7.8 – Group trends for selected Nonmetals

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Development of Periodic Table

The periodic table is the most significant tool that chemists use for organizing elements and recalling chemical facts.

Elements in the same column contain the same number of outer-shell electrons or valence electrons and have consequently similar chemical properties.

Development of Periodic Table

Mendeleev and Meyer arranged the elements in order of increasing atomic weight.

Certain elements were missing from their scheme.

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Development of Periodic Table

in 1871 Mendeleev noted that As properly belonged underneath P and not Si, which left a missing element underneath Si. He predicted a number of properties for this missing element (which he called eka-silicon or Germanium) with chemical properties similar to those of silicon.

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Periodic Trends

In this chapter, we will rationalize observed trends in:

• Sizes of atoms and ions.

• Ionization energy.

• Electron affinity.

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.

The electron is attracted to the nucleus, but repelled by core electrons that shield or screen it from the full nuclear charge.

This shielding is called the screening effect.

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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.

Na (Z=11)

The effective charge Zeff = Z − S

= 11+ -10 + =1+

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Sizes of Atoms

Consider a collection of argon atoms in the gas phase.

During collisions electron clouds cannot penetrate each other to a significant extent.

The apparent radius is determined by half of the closest distances separating the nuclei during such collisions.

This radius is the nonbonding radius.

Sizes of atoms in molecules

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

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Example

The distance separating the iodine (I2) nuclei is 2.66 A° , thus the radius is 1.33 A°.

Sizes of Atoms

Knowing the atomic radii allows the estimation of the bond lengths between different elements in molecules. In the compound CCl4 the measured length of C-Cl bond is 1.77 A° which is very close to the sum of (0.77A°+ 0.99 A°) for C and Cl respectively

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Sizes of Atoms

Bonding atomic radius tends to decrease from left to right across a row

due to increasing ZEFF which draws the electrons closer to the nucleus causing the atom to decrease in size.

Bonding atomic radius tends to increase from top to bottom of a column due to increasing value of n

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Sizes of Ions

The radii of ions are based on the distances between ion nuclei in an ionic compound.

Ionic size depends upon:

• Nuclear charge.

• Number of electrons.

• Orbitals in which electrons reside.

Sizes of Cations

Cations are smaller than their parent atoms.

• The outermost electron is removed and repulsions are reduced.

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Sizes of Anions

Anions are larger than their parent atoms.

• Electrons are added and repulsions are increased.

Sizes of Ions

Ions increase in size as you go down a column.

• Due to increasing value of n.

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Sizes of Ions

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

In an isoelectronic series, ionic size decreases with an increasing nuclear charge.

Ionization Energy Amount of energy required to

remove an electron from the ground state of a gaseous atom or ion.

• First ionization energy is the energy required to remove the first electron.

• Second ionization energy is the energy required to remove the second electron, etc.

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Ionization Energy

More energy is needed to remove each successive electron since the effective nuclear charges increase (more attractions).

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.

Trends in First Ionization Energies

Within each group the ionization energy generally decreases with increasing atomic number. The alkali metals have the lowest first ionization energies

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Irregularities in First Ionization

Energies

There are two apparent irregularities:

The first occurs between Groups IIIA and IIA. Electrons are more easily removed from p-orbitals than a s-orbital.

The second occurs between Groups VA and VIA.

• Electron removed comes from doubly occupied orbital. Repulsion from other electron in the same orbital helps in its removal

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Electron Affinity

Energy change accompanying addition of an electron to a gaseous atom:

Cl (g) + e- Cl− (g)

Trends in Electron Affinity

In general, electron

affinity becomes more exothermic as you go from left to right across a row. The greater the attraction between atom and added electron the more negative is the electron affinity. A positive value means that an electron cannot be attached to an atom.

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Trends in Electron Affinity

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

Trends in Electron Affinity

The first occurs between Groups IA and IIA.

• Added electron must go in p-orbital, not s-orbital (highly energetic subshell which is empty in the neutral atom .

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Trends in Electron Affinity

The second occurs between Groups IVA and VA.

• Group VA has no empty orbitals.

• Extra electron must go into occupied orbital, creating repulsion.

Properties of Metal, Nonmetals,

and Metalloids

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Metals versus Nonmetals

Metals tend to form cations.

Nonmetals tend to form anions.

Metals versus Nonmetals

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Metals

Tend to be malleable and good conductors of heat and electricity.

Compounds formed between metals and nonmetals tend to be ionic.

Metal oxides tend to be basic.

Nonmetals

brittle to hard substances that are poor conductors of heat and electricity.

Tend to gain electrons in reactions with metals to acquire noble gas configuration.

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Nonmetals

Substances containing only nonmetals are molecular compounds.

Most nonmetal oxides are acidic.

Metalloids

Have some characteristics of metals, some of nonmetals.

For instance, silicon looks shiny, but is brittle and fairly poor conductor.

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Group Trends

Alkali Metals

Soft, metallic solids.

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Alkali Metals

Found only as compounds in nature.

Have low densities and melting points.

Also have low ionization energies.

Alkali Metals

Their reactions with water are famously exothermic.

Used in distillation stills.

Story of the guy burning his hand in Notts.

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Alkaline Earth Metals

Have higher densities and melting points than alkali metals.

Have low ionization energies, but not as low as alkali metals.

Alkaline Earth Metals

Beryllium does not react with water but others react readily with water.

Reactivity tends to increase as you go down a group.

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Group 6A

Oxygen, sulfur, and selenium are nonmetals.

Tellurium is a metalloid.

The radioactive polonium is a metal.

Oxygen

Two allotropes: • O2

• O3, ozone

Three anions: • O2−, oxide

• O22−, peroxide

• O2−, superoxide

Tends to take electrons from other elements (oxidation)

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Sulfur

Weaker oxidizing agent than oxygen.

Most stable allotrope is S8, a ringed molecule. 30 allotropes

Group VIIA: Halogens

Name comes from the Greek halos and gennao: “salt formers”

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Group VIIA: Halogens

Large, negative electron affinities

• Therefore, tend to oxidize other elements easily

React directly with metals to form metal halides

Chlorine added to water supplies to serve as disinfectant

Group VIIIA: Noble Gases

Xe forms three compounds:

• XeF2

• XeF4 (at right)

• XeF6

Kr forms only one stable compound:

• KrF2

The unstable HArF was synthesized in 2000.

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