1. Chemistry: Atoms First Julia Burdge & Jason Overby
Chapter 4
Periodic Trends
of the Elements
Kent L. McCorkle
Cosumnes River College
Sacramento, CA
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. Periodic Trends of the Elements4
4.1 Development of the Periodic Table
4.2 The Modern Periodic Table
Classification of Elements
4.3 Effective Nuclear Charge
4.4 Periodic Trends in Properties of Elements
Atomic Radius
Ionization Energy
Electron Affinity
Metallic Character
4.5 Electron Configuration of Ions
Ions of Main Group Elements
Ions of d-Block Elements
4.6 Ionic Radius
Comparing Ionic Radius with Atomic Radius
Isoelectronic Series

3. Development of the Periodic Table
4.1
In 1864, John Newlands noted that when the elements were arranged in order of atomic number that every eighth element had similar properties.
He referred to this as the law of octaves.
In 1869, Dmitri Mendeleev and Lothar Meyer independently proposed the idea of periodicity.
Mendeleev grouped elements (66) according to properties.
Mendeleev predicted properties for elements not yet discovered, such as Ga.

4. Development of the Periodic Table
However, Mendeleev could not explain inconsistencies such as argon coming before potassium in the periodic table, despite having a higher atomic mass.
In 1913, Henry Moseley discovered the correlation between the number of protons (atomic number) and frequency of X-rays generated.
Ordering the periodic table by atomic number instead of atomic mass enabled scientists to make sense of discrepancies.
Entries today include atomic number and symbol; and are arranged according to electron configuration.

5. Worked Example 4.1
What elements would you expect to exhibit properties most similar to those of chlorine?
Strategy Because elements in the same group tend to have similar properties, you should identify elements in the same group (7A) as chlorine.
Solution Fluorine, bromine, and iodine, the other nonmetals in Group 7A, should have properties most similar to those of chlorine.
Think About It Astatine (At) is also in Group 7A. Astatine, though, is classified as a metalloid, and we have to be careful comparing nonmetals to metalloids (or to metals). As a metalloid, the properties of astatine should be less similar to those of chlorine than the other members of Group 7A.

6. The Modern Periodic Table
4.2

7. Classification of Elements
The main group elements (also called the representative elements) are the elements in Groups 1A through 7A.

8. Classification of Elements
The noble gases are found in Group 8A and have completely filled p subshells.

9. The Modern Periodic Table
The transition metals are found in Group 1B and 3B through 8B.
Group 2B have filled d subshells and are not transition metals.

10. The Modern Periodic Table
The lanthanides and actinides make up the f-block transition elements.

11. The Modern Periodic Table
There is a distinct pattern to the electron configurations of the elements in a particular group.
For Group 1A: [noble gas]ns1
For Group 2A: [noble gas]ns2

12. The Modern Periodic Table
The outermost electrons of an atom are called the valence electrons.
Valence electrons are involved in the formation of chemical bonds.
Similarity of valence electron configurations help predict chemical properties.
For Group 1A: [noble gas]ns1
valence
core
For Group 2A: [noble gas]ns2
valence
core
For Group 7A: [noble gas]ns2np5
valence
core

13. Worked Example 4.2
Without using a periodic table, give the ground-state electron configuration and block designation (s-, p-, d-, or f-block) of an atom with (a) 17 electrons, (b) 37 electrons, and (c) 22 electrons. Classify each atom as a main group element or transition metal.
Strategy Use the figure at right to assign electrons to orbitals in the correct order. Recall that an s subshell has one orbital, a p subshell has three orbitals, and a d subshell has five orbitals. Remember, too, that each orbital can accommodate a maximum of two electrons.
Solution
1s22s22p63s23p5, p-block, main group
1s22s22p63s23p64s23d104p65s1, s-block, main group
1s22s22p63s23p64s23d2, d-block, transition metal
Think About It Consult the figure to confirm your answer.

14. Effective Nuclear Charge
4.3
Effective nuclear charge (Zeff) is the actual magnitude of positive charge that is “experienced” by an electron in the atom.
In a multi-electron atom, electrons are simultaneously attracted to the nucleus and repelled by one another.
This results in shielding, where an electron is partially shielded from the positive charge of the nucleus by the other electrons.
Although all electrons shield one another to some extent, the most effective are the core electrons.
As a result, the value of Zeff increases steadily from left to right because the core electrons remain the same but Z increases. Li Be B C N O F Z 3 4 5 6 7 8 9
Zeff
1.28
1.91
2.42
3.14
3.83
4.45
5.10

15. Effective Nuclear Charge
4.3
In general, the effective nuclear charge is given by
Z is the nuclear charge or simply the number of protons in the nucleus.
σ is the shielding constant.
Zeff increases from left to right across a period; changes very little down a column.
Zeff = Z – σ Li Be B C N O F Z 3 4 5 6 7 8 9
Zeff
1.28
1.91
2.42
3.14
3.83
4.45
5.10

16. Periodic Trends in Properties of Elements
4.4
Atomic radius is the distance between the nucleus of an atom and its valence shell.
(a) Atomic radius in metals, or metallic radius, is half the distance between the nuclei of two adjacent, identical metal atoms.
(b) Atomic radius in nonmetals, or covalent radius, is half the distance between adjacent, identical nuclei connected by a chemical bond.

17. Atomic radii (in picometers)
Atomic Radius
Atomic radii (in picometers)
The atomic radius increases from top to bottom down a group.
Increasing n, so outermost shell lies farther from the nucleus
Atomic radius decreases from left to right across a period.
Increasing Zeff which draws the valence shell closer to the nucleus

18. Atomic Radius
Atomic radius decreases left to right across a period due to increased electrostatic attraction between the effective nuclear charge and the charge on the valence shell.

19. Worked Example 4.3
Referring only to a periodic table, arrange the elements P, S, and O in order of increasing atomic radius.
Strategy Use the left-to-right (decreasing) and top-to-bottom (increasing) trends to compare the atomic radii of two of the three elements at a time.
Solution
O < S < P
Think About It Consult the atomic radii figure to confirm the order. Note that there are circumstances under which the trends alone will be insufficient to compare the radii of two elements. Using only a periodic table, for example, it would not be possible to determine that bromine (r = 114 pm) has a smaller radius than silicon (r = 118 pm).

20. Ionization Energy
Ionization energy (IE) is the minimum energy required to remove an electron from an atom in the gas phase.
The result is an ion, a chemical species with a net charge.
Sodium has an ionization energy of kJ/mol.
Specifically, kJ/mol is the first ionization energy of sodium, IE1(Na), which corresponds to the removal of the most loosely held electron.
Na(g) → Na+(g) + e−

21. Ionization Energy

22. Ionization Energy
In general, as Zeff increases, ionization energy also increases.
Thus, IE1 increases from left to right across a period.

23. Ionization Energy
Within a given shell, electrons with a higher value of l are higher in energy and thus, easier to remove.

24. Ionization Energy
Removing a paired electron is easier because of the repulsive forces between two electrons in the same orbital.

25. Ionization Energy
It is possible to remove additional electrons in subsequent ionizations, giving IE1, IE2, and so on.
IE1(Na) = 496 kJ/mol
IE2(Na) = 4562 kJ/mol
Na(g) → Na+(g) + e−
Na+(g) → Na2+(g) + e−

26. Ionization Energy
It takes more energy to remove the 2nd, 3rd, 4th, etc. electrons because it is harder to remove an electron from a cation than an atom.
It takes much more energy to remove core electrons than valence.
Core electrons are closer to nucleus.
Core electrons experience greater Zeff because of fewer filled shells shielding them from the nucleus.

27. Worked Example 4.4
Would you expect Na or Mg to have the greater first ionization energy (IE1)? Which should have the greater second ionization energy (IE2)?
Strategy Consider effective nuclear charge and electron configuration to compare the ionization energies. Na has one valence electron and Mg has two.
Solution
IE1(Mg) > IE1(Na) because Mg is to the right of Na in the periodic table (i.e., Mg has the greater Zeff, so it is more difficult to remove its electron).
IE2(Na) > IE2(Mg) because the second ionization of Mg removes a valence electron, whereas the second ionization of Na removes a core electron.
Think About It The first ionization energies of Na and Mg are 496 and 738 kJ/mol, respectively. The second ionization energies of Na and Mg are 4562 and 1451 kJ/mol, respectively.

28. Electron Affinity
Electron affinity (EA) is the energy released when an atom in the gas phase accepts an electron.
Cl(g) + e− → Cl−(g)

29. Electron Affinity
Like ionization energy, electron affinity increases from left to right across a period as Zeff increases.
Easier to add an electron as the positive charge of the nucleus increases.

30. Electron Affinity
It is easier to add an electron to an s orbital than to add one to a p orbital with the same principal quantum number.

31. Electron Affinity
Within a p subshell, it is easier to add an electron to an empty orbital than to add one to an orbital that already contains an electron.

32. Electron Affinity
More than one electron may be added to an atom.
While many first electron affinities are positive, subsequent electron affinities are always negative.
Considerable energy is required to overcome the repulsive forces between the electron and the negatively charged ion.
Process
Electron Affinity
O(g) + e− → O−(g)
EA1 = 141 kJ/mol
O− (g) + e− → O2−(g)
EA2 = −741 kJ/mol

33. Metallic Character
Metals tend to
Be shiny, lustrous, malleable, and ductile
Be good conductors of both heat and electricity
Have low ionization energies (commonly form cations)

34. Metallic Character
Nonmetals tend to
Vary in color and are not shiny
Be brittle, rather than malleable
Be poor conductors of both heat and electricity
Have high electron affinities (commonly form anions)

35. Metallic Character
Metalloids are elements with properties intermediate between those of metals and nonmetals.

36. Electron Configurations of Ions
4.5
To write the electron configuration of an ion formed by a main group element:
Write the configuration for the atom.
Add or remove the appropriate number of electrons.
Na: 1s22s22p63s1
Na+: 1s22s22p6
10 electrons total, isoelectronic with Ne
Cl: 1s22s22p63s23p5
Cl− : 1s22s22p63s23p6
18 electrons total, isoelectronic with Ar

37. Worked Example 4.7
Write electron configurations for the following ions of main group elements: (a) N3-, (b) Ba2+, and (c) Be2+.
Strategy First write electron configurations for the atoms. Then add electrons (for anions) and remove electrons (for cations) to account for the charge.
Solution
[He]2s22p6 or [Ne]
[Kr]5s24d105p6 or [Xe]
1s2 or [He]
Think About It Be sure to add electrons to form an anion, and remove electrons to form a cation.

38. Ions of d-Block Elements
Ions of d-block elements are formed by removing electrons first from the shell with the highest value of n.
For Fe to form Fe2+, two electrons are lost from the 4s subshell not the 3d.
Fe can also form Fe3+, in which case the third electron is removed from the 3d subshell.
Fe: [Ar]4s23d 6
Fe2+: [Ar]3d 6
Fe: [Ar]4s23d 6
Fe3+: [Ar]3d 5

39. Worked Example 4.8
Write electron configurations for the following ions of d-block elements: (a) Zn2+, (b) Mn2+, and (c) Cr3+.
Strategy First write electron configurations for the atoms. Then add electrons (for anions) and remove electrons (for cations) to account for the charge. The electrons removed from a d-block element must come first from the outermost s subshell, not the partially filled d subshell.
Solution
[Ar]3d10
[Ar]3d5
[Ar]3d3
Think About It Be sure to add electrons to form an anion, and remove electrons to form a cation. Also, double-check to make sure that electrons removed from a d-block elment come first from the ns subshell and then, if necessary, from the (n – 1)d subshell.

40. Ionic Radius
4.6
The ionic radius is the radius of a cation or an anion.
When an atom loses an electron to become a cation, its radius decreases due in part to a reduction in electron-electron repulsions in the valence shell.
A significant decrease in radius occurs when all of an atom’s valence electrons are removed.

41. Comparing Ionic Radius with Atomic Radius
When an atom gains one or more electrons and becomes an anion, its radius increases due to increased electron-electron repulsions.

42. O2− : 1s22s22p6 F−: 1s22s22p6 Ne: 1s22s22p6 Isoelectronic Series
An isoelectronic series is a series of two or more species that have identical electron configurations, but different nuclear charges.
O2− : 1s22s22p6
F−: 1s22s22p6
Ne: 1s22s22p6
isoelectronic

43. Worked Example 4.9
Identify the isoelectronic series in the following group of species, and arrange them in order of increasing radius: K+, Ne, Ar, Kr, P3-, S2-, and Cl-.
Strategy Isoelectronic series are species with identical electron configurations but different nuclear charges. Determine the number of electrons in each species. The radii of isoelectronic series members decreases with increasing nuclear charge.
Solution
The isoelectronic series includes K+, Ar, P3-, and Cl- (18 electrons each).
In order of increasing radius: K+ < Ar < Cl- < S2- < P3-.
Think About It With identical electron configurations, the attractive force between the valence electrons and the nucleus will be strongest for the largest nuclear charge. Thus, the larger the nuclear charge, the closer in the valence electrons will be pulled and the smaller the radius will be.

44. Summary of periodic trends
Effective Nuclear Charge