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3 Interpreting the Periodic Table

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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 p 11 e 11 p 10 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 p 17 e 17 p 18 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.

11 Physical state of elements

12 Chemical Periodicity

13 Periodic Properties of the Elements Periodic Properties of the Elements Atomic 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.

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

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

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

17 Ionization Energy First ionization energy (IE 1 ) 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 -

18 Ionization Energy Second ionization energy (IE 2 ) The amount of energy required to remove the second electron from a gaseous 1+ ion. Symbolically: ion + + energy  ion 2+ + e - Mg kJ/mol  Mg 2+ + e - Atoms can have 3 rd (IE 3 ), 4 th (IE 4 ), etc. ionization energies.

19 Ionization Energy Periodic trends for Ionization Energy: 1. IE 2 > IE 1 It always takes more energy to remove a second electron from an ion than from a neutral atom. 2. IE 1 generally increases moving from left to right in the same period. 3. IE 1 generally decreases moving down a group. IE 1 for Li > IE 1 for Na, etc.

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

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

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

23 Ionization Energy Group and element IA Na IIA Mg IIIA Al IVA Si IE 1 (kJ/mol) IE 2 (kJ/mol) IE 3 (kJ/mol) IE 4 (kJ/mol) ,55011,

24 Ionization Energy The reason Na forms Na + and not Na 2+ is that the energy difference between IE 1 and IE 2 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 Mg 2+ and not Mg 3+. Al forms Al 3+.

25 Ionization Energy Example: What charge ion would be expected for an element that has these ionization energies? IE 1 (kJ/mol)1680 IE 2 (kJ/mol)3370 IE 3 (kJ/mol)6050 IE 4 (kJ/mol)8410 IE 5 (kJ/mol)11020 IE 6 (kJ/mol)15160 IE 7 (kJ/mol)17870 IE 8 (kJ/mol)92040 Notice that the largest increase in ionization energies occurs between IE 7 and IE 8. Thus this element would form a 1- ion.

26 Ionization Energy Example: What charge ion would be expected for an element that has these ionization energies? IE 1 (kJ/mol)1680 IE 2 (kJ/mol)3370 IE 3 (kJ/mol)11586 IE 4 (kJ/mol)12410 IE 5 (kJ/mol)13020 IE 6 (kJ/mol)15160 IE 7 (kJ/mol)17870 IE 8 (kJ/mol)18040 Notice that the largest increase in ionization energies occurs between IE 2 and IE 3. Thus this element would form a 2+ ion.

27 Ionization Energy Example: What charge ion would be expected for an element that has these ionization energies? IE 1 (kJ/mol)1680 IE 2 (kJ/mol)3370 IE 3 (kJ/mol)4586 IE 4 (kJ/mol)5410 IE 5 (kJ/mol)6020 IE 6 (kJ/mol)7160 IE 7 (kJ/mol)17870 IE 8 (kJ/mol)18040 Notice that the largest increase in ionization energies occurs between IE 6 and IE 7. Thus this element would form a 2- ion.

28 Ionic Radii Cations (positive ions) are always smaller than their respective neutral atoms. ElementLiBe Atomic Radius (Å) IonLi + Be 2+ Ionic Radius (Å)

29 Ionic Radii Cations (positive ions) are always smaller than their respective neutral atoms. ElementNaMgAl Atomic Radius (Å) IonNa + Mg 2+ Al 3+ Ionic Radius (Å)

30 Ionic Radii Anions (negative ions) are always larger than their neutral atoms. ElementNOF Atomic Radius(Å) IonN 3- O 2- F 1- Ionic Radius(Å)

31 Ionic Radii

32 Ionic Radii Cation (positive ions) radii decrease from left to right across a period. Increasing nuclear charge attracts the electrons and decreases the radius. IonRb + Sr 2+ In 3+ Ionic Radii(Å)

33 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. IonN 3- O 2- F 1- Ionic Radii(Å)

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

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

36 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.

37 Electronegativity

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

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

See animation for Summary of Trends in the Periodic Table