1. Unit 1: Structure of Matter
Sections: 2.4, 8
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2. When the Elements Were Discovered

3. The Modern Periodic Table
Group
Period
Metal: good conductor of heat and electricity
Nonmetal: usually a poor conductor of heat and electricity
Metalloid: has properties that are intermediate between those of metals and nonmetals

4. Classification of the Elements

5. The Modern Periodic Table
Alkali Earth Metal
The Modern Periodic Table
Halogen
Noble Gas
Alkali Metal

6. Ground State Electron Configurations of the Elements
ns2np6
Ground State Electron Configurations of the Elements
ns1
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns2
d10 d1 d5
Valence electrons: the outermost electrons 4f 5f
Core electrons: all nonvalence electrons in an atom

7. Review of Concepts
In viewing the periodic table, do chemical properties change more markedly across a period or down a group?

8. Group 1A Elements (ns1, n  2) Alkali Metals
M M+1 + 1e-
2M(s) + 2H2O(l) MOH(aq) + H2(g)
Low IE = very reactive
Never found in the pure state
4M(s) + O2(g) M2O(s)
Increasing reactivity

9. Group 1A Elements (ns1, n  2)

10. Group 2A Elements (ns2, n  2) Alkaline Earth Metals
M M+2 + 2e-
Be(s) + 2H2O(l) No Reaction
Mg(s) + 2H2O(g) Mg(OH)2(aq) + H2(g)
M(s) + 2H2O(l) M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba
Increasing reactivity

11. Group 2A Elements (ns2, n  2)

12. Comparison of Group 1A and 1B
The metals in these two groups have similar outer electron configurations, with one electron in the outermost s orbital.
Chemical properties are quite different due to difference in the ionization energy.
Lower IE1, more reactive

13. Group 3A Elements (ns2np1, n  2)

14. Group 3A Elements (ns2np1, n  2)

15. Group 4A Elements (ns2np2, n  2)

16. Group 4A Elements (ns2np2, n  2)

17. Group 5A Elements (ns2np3, n  2)

18. Group 5A Elements (ns2np3, n  2)

19. Group 6A Elements (ns2np4, n  2)

20. Group 6A Elements (ns2np4, n  2)

21. Group 7A Elements (ns2np5, n  2) Halogens
X + 1e X-1
X2(g) + H2(g) HX(g)
Increasing reactivity

22. Group 7A Elements (ns2np5, n  2)

23. Group 8A Elements (ns2np6, n  2)
Noble Gas
Completely filled ns and np subshells.
Highest ionization energy of all elements.
No tendency to accept extra electrons.

25. Compounds of the Noble Gases
A number of xenon compounds XeF4, XeO3, XeO4, XeOF4 exist.
A few krypton compounds (KrF2, for example) have been prepared.

26. Practice Exercise
8.1
An atom of a certain element has 20 electrons. Without consulting a periodic table, answer the following questions:
What is the ground-state electron configuration of the element?
(b) How should the element be classified?
(c) Is the element diamagnetic or paramagnetic?

27. Representing Free Elements in Chemical Equations
Metals
Nonmetals
Do not exist in discrete molecular units
Use empirical formulas in chemical equation
Empirical formulas are the same as the symbols that represent the elements
No single rule
Complex three-dimensional network of atoms, use empirical formula
Carbon
Metalloids
Diatomic and polyatomic molecules use molecular formulas
HONClBrIF
P4, S8
Noble gases (monatomic species) use their symbols

28. Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Atoms lose electrons so that cation has a noble-gas outer electron configuration.
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
H 1s1
H- 1s2 or [He]
Atoms gain electrons so that anion has a noble-gas outer electron configuration.
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]

29. Cations and Anions Of Representative Elements+1 +2 +3 -3 -2 -1

30. Isoelectronic: have the same number of electrons, and hence the same ground-state electron configuration
Na+: [Ne]
Al3+: [Ne]
F-: 1s22s22p6 or [Ne]
O2-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

31. Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.
Fe: [Ar]4s23d6
Mn: [Ar]4s23d5
Fe2+: [Ar]4s03d6 or [Ar]3d6
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5
*Most transition metals can form more than one cation and frequently the cations are not isoelectronic with the preceding noble gas
*The order of electron filling does not determine or predict the order of electron removal for transition metals.

32. Identify the elements that fit the following descriptions:
Review of Concepts
Identify the elements that fit the following descriptions:
An alkaline earth metal ion that is isoelectronic with Kr
An anion with a -3 charge that is isoelectronic with K+
An ion with a +2 charge that is isoelectronic with Co3+

33. Effective nuclear charge (Zeff) is the “positive charge” felt by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons
Zeff
Core Z
Radius (pm) Na Mg Al Si 11 12 13 14 10 1 2 3 4
186
160
143
132

34. Trend in Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff
*core electrons shield valence electrons much more than valence electrons shield one another

35. Atomic Radii covalent radius metallic radius
Atomic radius is one-half the distance between the two nuclei in two adjacent metal atoms or in a diatomic molecule (or one-half the distance between the nuclei in two neighboring atoms).

37. Trends in Atomic Radii

38. Practice Exercise
8.2
Referring to a periodic table, arrange the following atoms in order of decreasing atomic radius: C, Li, Be.
N < P < Si

39. Compare the size of each pair of atoms listed here:
Review of Concepts
Compare the size of each pair of atoms listed here:
Be, Ba
Al, S
12C, 13C

40. Video

41. Comparison of Atomic Radii with Ionic Radii

42. Cation is always smaller than atom from which it is formed.
Anion is always larger than atom from which it is formed.

43. The Radii (in pm) of Ions of Familiar Elements

44. Practice Exercise
8.3
For each of the following pairs, indicate which one of the two species is smaller:
K+ or Li+
Au+ or Au3+
P3− or N3−

45. Review of Concepts
Identify the spheres shown here with each of the following: S2−, Mg2+, F−, Na+

46. IE1 first ionization energy
Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.
*the magnitude of ionization energy is a measure of how “tightly” the electron is held in the atom
IE1 + X (g) X+(g) + e-
IE1 first ionization energy
IE2 + X+(g) X2+(g) + e-
IE2 second ionization energy
IE3 + X2+(g) X3+(g) + e-
IE3 third ionization energy
IE1 < IE2 < IE3
*Chemical properties of any atom are determined by the configuration of the atom’s valence electrons. The stability of these outermost electrons is reflected directly in the atom’s ionization energies.

47. *Note: while valence electrons are relatively easy to remove from the atom, core electrons are much harder to remove. Thus, there is a large jump in ionization energy between the last valence electron and the first core electron.

48. Variation of the First Ionization Energy with Atomic Number
Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell

49. General Trends in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy

50. Exceptions Between Group 2A and 3A (ex: Be and B)
Group 3A is lower b/c of the single electron in the outermost p subshell which is shielded
Between Group 5A and 6A (ex: N and O)
Group 6A doubles up one electron, the proximity of two electrons in the same orbital results in greater electrostatic repulsion and lower energy

51. Practice Exercise
8.4
Which atom should have a larger first ionization energy: nitrogen or phosphorus?
Which atom should have a smaller second ionization energy: sodium or magnesium?

52. Review of Concepts
Label the plots shown here for the firs, second, and third ionization energies for Mg, Al, and K.

53. *lg. pos. electron affinity means that the negative ion is very stable
Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.
*or think of it as the energy that must be supplied to remove an electron from the anion.
X (g) + e− X− (g)
F (g) + e− F− (g)
DH = -328 kJ/mol
EA = +328 kJ/mol
*lg. pos. electron affinity means that the negative ion is very stable
O (g) + e− O− (g)
DH = -141 kJ/mol
EA = +141 kJ/mol
*The more positive the electron affinity is of an element, the greater the affinity of an atom of the element to accept an electron.

54. Variation of Electron Affinity With Atomic Number (H – Ba)

55. Practice Exercise
8.5
Why are the electron affinities of the alkaline earth metals, shown in Table 8.3, either negative or small positive values?
Is it likely that Ar will form the anion Ar−

56. Review of Concepts
Why is it possible to measure the successive ionization energies of an atom until all the electrons are removed, but it becomes increasingly difficult and often impossible to measure the electron affinity of an atom beyond the first stage?

57. Diagonal Relationships on the Periodic Table
*similarities between pairs of elements in different groups and periods of the periodic table