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Interpreting the Periodic Table

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Interpreting the Periodic Table

1. Typically they have a shiny luster.

2. Relatively high density.

3. Malleable ( they can be hammered into thin sheets).

4. Ductile (they can be drawn into a wire).

5. Good conductors of electricity.

6. Reflect light and heat.

7. High melting and boiling points so they are solids at room temperature (except Hg).

8. Lose electron(s) forming positive ions (cations).

9. Combine with non-metals.

10. Do not readily combine with each other.

General Properties of Metallic Elements:

Examples of Metals

Potassium, K reacts with water and must be stored in kerosene

Zinc, Zn, is more stable than potassium

Copper, Cu, is a relatively soft metal, and a very good electrical conductor.

Mercury, Hg, is the only metal that exists as a liquid at room temperature

11 p11 e

11 p10 e

• Metals lose electrons forming positive ions.

• The radius of the cation is always smaller than the radius of the parent atom.

1. Poor conductors of heat and electricity.

2. Not malleable or ductile; fragile.

3. Low densities.

4. Low melting and boiling points so they can be gases, liquids, or solids.

5. Gain electron(s) forming negative ions (anions).

6. Combine with metals.

7. Combine with each other to a limited extent.

General Properties of Non-metallic Elements:

17 p17 e

17 p18 e

• Non metals gain electrons forming negative ions.

• The radius of the anion is always larger than the radius of the parent atom.

Metalloids: elements that lie in the colored stair step line of the Periodic Table. They have both metallic and non-metallic properties.

These elements are weak conductors of electricity, which makes them useful semiconductors in the integrated circuits of computers.

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Physical state of elements

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Chemical Periodicity

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Periodic Properties of the ElementsAtomic Radii

Atomic radii describes the relative sizes of atoms. It is understood as the distance from the nucleus to the outermost occupied energy level.Atomic radii increase within a column going from the top to the bottom of the periodic table.Atomic radii decrease within a row going from left to right on the periodic table.

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Atomic Radii

Example: Arrange these elements based on their atomic radii. Se, S, O, Te

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Atomic Radii

Example: Arrange these elements based on their atomic radii. P, Cl, S, Si

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Atomic Radii

Example: Arrange these elements based on their atomic radii. Ga, F, S, As

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Ionization Energy

First ionization energy (IE1) The minimum amount of energy required to remove the

most loosely bound electron from an isolated gaseous atom to form a 1+ ion.

Symbolically:Atom(g) + energy ion+

(g) + e-

Mg(g) + 738kJ/mol Mg+ + e-

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Ionization Energy

Second ionization energy (IE2) The amount of energy required to remove the

second electron from a gaseous 1+ ion.

Symbolically: ion+ + energy ion2+ + e-

Mg+ + 1451 kJ/mol Mg2+ + e-

•Atoms can have 3rd (IE3), 4th (IE4), etc. ionization energies.

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Ionization EnergyPeriodic trends for Ionization Energy:

1. IE2 > IE1

It always takes more

energy to remove a second electron from an ion than from a neutral atom.

2. IE1 generally increases moving from left to right in the same period.

3. IE1 generally decreases moving down a group.

IE1 for Li > IE1 for Na, etc.

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Ionization Energy

Example: Arrange these elements based on their first ionization energies. Sr, Be, Ca, Mg

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Ionization Energy

Example: Arrange these elements based on their first ionization energies. Al, Cl, Na, P

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Ionization Energy

Example: Arrange these elements based on their first ionization energies. B, O, Be, N

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Ionization EnergyGroup

and element

IANa

IIAMg

IIIAAl

IVASi

IE1 (kJ/mol)

496 738 578 786

IE2

(kJ/mol)4562 1451 1817 1577

IE3

(kJ/mol)6912 7733 2745 3232

IE4

(kJ/mol)9540 10,550 11,580 4356

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Ionization Energy

The reason Na forms Na+ and not Na2+ is that the energy difference between IE1 and IE2 is so large. Requires more than 9 times more energy to

remove the second electron than the first one.

The same trend is persistent throughout the series. Thus Mg forms Mg2+ and not Mg3+. Al forms Al3+.

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Ionization Energy

Example: What charge ion would be expected for an element that has these ionization energies?

IE1 (kJ/mol) 1680

IE2 (kJ/mol) 3370

IE3 (kJ/mol) 6050

IE4 (kJ/mol) 8410

IE5 (kJ/mol) 11020

IE6 (kJ/mol) 15160

IE7 (kJ/mol) 17870

IE8 (kJ/mol) 92040

Notice that the largest increase in ionization energies occurs between IE7 and IE8. Thus this element would form a 1- ion.

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Ionization Energy

Example: What charge ion would be expected for an element that has these ionization energies?

IE1 (kJ/mol) 1680

IE2 (kJ/mol) 3370

IE3 (kJ/mol) 11586

IE4 (kJ/mol) 12410

IE5 (kJ/mol) 13020

IE6 (kJ/mol) 15160

IE7 (kJ/mol) 17870

IE8 (kJ/mol) 18040

Notice that the largest increase in ionization energies occurs between IE2 and IE3. Thus this element would form a 2+ ion.

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Ionization Energy

Example: What charge ion would be expected for an element that has these ionization energies?

IE1 (kJ/mol) 1680

IE2 (kJ/mol) 3370

IE3 (kJ/mol) 4586

IE4 (kJ/mol) 5410

IE5 (kJ/mol) 6020

IE6 (kJ/mol) 7160

IE7 (kJ/mol) 17870

IE8 (kJ/mol) 18040

Notice that the largest increase in ionization energies occurs between IE6 and IE7. Thus this element would form a 2- ion.

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Ionic Radii

Cations (positive ions) are always smaller than their respective neutral atoms.Element Li Be

Atomic Radius (Å)

1.52 1.12

Ion Li+ Be2+

Ionic Radius (Å)

0.90 0.59

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Ionic Radii

Cations (positive ions) are always smaller than their respective neutral atoms.

Element Na Mg Al

Atomic Radius (Å)

1.86 1.60 1.43

Ion Na+ Mg2+ Al3+

Ionic Radius (Å)

1.16 0.85 0.68

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Ionic Radii

Anions (negative ions) are always larger than their neutral atoms.

Element N O F

AtomicRadius(Å)

0.75 0.73 0.72

Ion N3- O2- F1-

IonicRadius(Å)

1.71 1.26 1.19

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Ionic Radii

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Ionic Radii

Cation (positive ions) radii decrease from left to right across a period. Increasing nuclear charge attracts the electrons and

decreases the radius.

Ion Rb+ Sr2+ In3+

IonicRadii(Å)

1.66 1.32 0.94

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Ionic Radii

Anion (negative ions) radii decrease from left to right across a period. Increasing electron numbers in highly charged ions

cause the electrons to repel and increase the ionic radius.

Ion N3- O2- F1-

IonicRadii(Å)

1.71 1.26 1.19

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Ionic Radii

Example: Arrange these elements based on their ionic radii. Ga, K, Ca

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Ionic Radii

Example: Arrange these elements based on their ionic radii. Cl, Se, Br, S

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Electronegativity

Electronegativity is a measure of the relative tendency of an atom to attract electrons to itself when chemically combined with another element. Electronegativity is measured on the Pauling

scale. Fluorine is the most electronegative element. Cesium and francium are the least

electronegative elements.For the representative elements, electronegativities usually increase from left to right across periods and decrease from top to bottom within groups.

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Electronegativity

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Electronegativity

Example: Arrange these elements based on their electronegativity. Se, Ge, Br, As

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Electronegativity

Example: Arrange these elements based on their electronegativity. Be, Mg, Ca, Ba

See animation for Summary of Trends in the Periodic Table

http://www.learnerstv.com/animation/animation.php?ani=56&cat=chemistry

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