1. In 1869, Mendeleev created the first periodic table of the elements.
CHEMISTRY CHAPTER 5
THE PERIODIC LAW
SECTION 1. HISTORY OF THE
PERIODIC TABLE
Mendeleev and Chemical Periodicity
In 1869, Mendeleev created the first periodic table of the elements.
He arranged the 57 known elements in order of atomic mass.

2. He noticed that certain similarities in their chemical properties appeared at regular intervals.
Repeating patterns are referred to as periodic.
He left some gaps in his table and predicted that elements would be discovered to fill them.
By 1886, all three of these elements had been discovered.

3. Properties of Some Elements Predicted By Mendeleev

4. Moseley and the Periodic Law
In 1911, the English scientist Henry Moseley discovered that the elements fit into patterns better when they were arranged according to atomic number, rather than atomic mass.
(Why didn’t Mendeleev arrange by atomic numbers?)
The Periodic Law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

5. Periodicity of Atomic Numbers

6. The Periodic Table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.
It is explained by electron configurations.

7. Periodic Table Overview 75020

8. SECTION 2. ELECTRON CONFIGURATION AND THE PERIODIC TABLE
Periods and Blocks of the Periodic Table
Group = vertical column. Elements in the same group have similar chemical properties.
Period = horizontal row. In each new period electrons begin to be added to a new main energy level.

9. The length of each period is determined by the number of electrons that can occupy the sublevels being filled in that period.
Ex.: in Period 4, electrons are added to the 4s, 3d, and 4p sublevels, for a total of 18 electrons. The period has 18 elements.

10. Relationship Between Periodicity and Electron Configurations

11. The periodic table is divided into four blocks, the s, p, d, and f blocks. The name of each block is determined by the electron sublevel being filled in that block.

12. Relating Period Length and Sublevels Filled 75084

13. Blocks of the Periodic Table Based on Sublevel 75085

15. Group 1 = alkali metals
lithium, sodium, potassium, rubidium, cesium, and francium
The most reactive metals
Electron configuration ends in ns1
Ex: Li = 1s22s K = [Ar]4s1
Form +1 ions by losing one electron
Hydrogen has an electron configuration of 1s1, but does not share the same properties as the elements of Group 1. It has unique properties.

16. Group 2 = alkaline-earth metals.
beryllium, magnesium, calcium, strontium, barium, and radium
Reactive, but not as much as group 1.
Still too reactive to be found in nature in pure form.
Electron configuration ends in ns2.
Ex. – Mg = [Ne]3s2
Form +2 ions by losing 2 electrons

17. Like the Group 2 elements, helium has an ns2 group configuration
Like the Group 2 elements, helium has an ns2 group configuration. Yet it is part of Group 18.
Because its highest occupied energy level is filled by 2 electrons, helium possesses special chemical stability.

18. The p-block elements consist of all the elements of Groups 13– 18 except helium.
The p-block elements together with the s-block elements (Groups 1-2) are called the main-group elements.
The properties of elements of the p block vary greatly.

19. At its right-hand end, the p block includes all of the nonmetals except hydrogen and helium.
All of the metalloids are also in the p block.
At the left-hand side and bottom of the block, there are eight p-block metals.

20. Group 13: metals form +3 ions
Group 13: metals form +3 ions. Aluminum is the most abundant in this group.
Electron configuration has ns2np1. Ex:
Al: [Ne]3s23p1
Ga: [Ar]3d104s24p1
Group 14: carbon family.
Metals (Sn, Pb) are fairly stable.
Electron configuration has ns2np2.
Carbon uses these 4 electrons to form 4 covalent bonds.

21. Group 15: nitrogen family.
Nitrogen and phosphorus can form -3 ions.
Electron configuration has ns2np3.
Nitrogen (in the form of N2) is the most abundant gas in the atmosphere.
Group 16: oxygen family.
Nonmetals form -2 ions.
Electron configuration has ns2np4.
Oxygen is the most abundant element in Earth’s crust.

22. The elements of Group 17 are known as the halogens (salt- formers).
fluorine, chlorine, bromine, iodine, and astatine
Most reactive nonmetals.
React vigorously with most metals to form salts.
Electron configuration: ns2np5
Form -1 ions.

23. The metalloids, or semiconducting elements, form a “staircase” located between nonmetals and metals in the p block.
Boron, silicon, germanium, arsenic, antimony, tellurium, polonium (text does not include Po)

24. The metals of the p block are generally harder and denser than the s-block alkaline-earth metals, but softer and less dense than the d-block metals.

25. The d sublevel first appears when n = 3.
The 3d sublevel is slightly higher in energy than the 4s sublevel, so these are filled in the order 4s3d.

26. After 2 electrons are added to the ns level of period n, the next 10 elements have electrons added to the d orbitals of the previous main energy level. Example: n = 4
Ca = [Ar]4s (group 2)
Sc = [Ar]3d14s2 (group 3)
Ti = [Ar]3d24s (group 4) .
Zn = [Ar]3d104s2 (group 12)
(Then, in group 13, the p orbitals begin to be filled:
Ga = [Ar]3d104s24p1 (group 13)

27. Elements of the d-block, groups 3-12, are often referred to as the transition elements.
They are metals with typical metallic properties.
Cu, Ag, Au (group 11) are referred to as the coinage metals.

28. Beginning with period 6, electrons are added to f orbitals.
Lanthanides and Actinides = f Block = Inner Transition Elements
Beginning with period 6, electrons are added to f orbitals.
Lanthanides = elements 58-71, following element 57 (lanthanum). These have electrons added to the 4f orbitals.

29. Actinides = elements 90-103, following element 89 (actinium)
Actinides = elements , following element 89 (actinium). These have electrons added to the 5f orbitals. They are all radioactive, and most are synthetic.

30. The lanthanides and actinides are usually represented at the bottom of the periodic table rather than in their proper sequences.
There is no strict rule as to which group they belong to. Sometimes they are all considered to be part of group 3.

31. Rare Earths
In chemistry, “earth” is an old name used for certain compounds (oxides) that were difficult to separate into elements (and were therefore thought to be elements). Examples are the alkaline earths, from which the alkaline earth elements (group 2) were later isolated.

32. The term rare earths elements (or usually just “rare earths”) is used for scandium, yttrium, lanthanum, and the lanthanides. Rare earths are critical for modern electronic products. Currently there is concern because more than 95% of production is in China.

33. New York Times

34. New Elements Named or Discovered Since Textbook Published
A new element is not given an official name until its discovery is confirmed.
111 - Roentgenium (discovered 1994; name accepted 2004)
112 - Copernicum (discovered 1996; name accepted 2009)

35. 113 – Discovery reported 2003, accepted 2015, proposed name nihonium
114 - Flerovium (discovered 1999; name accepted 2012)
115 – Discovery reported 2003, confirmed 2013, proposed name moscovium

36. 116 - Livermorium (discovered 2000, name accepted 2012)
117 – Discovery reported 2010, accepted 2015, proposed name tennessine
118 – Discovery reported 2002, recognized 2015, proposed name oganesson

37. https://iupac.org/iupac-is-naming-the-four-new-elements-nihonium-moscovium-tennessine-and-oganesson/

38. SECTION 3. ELECTRON CONFIGURATION AND PERIODIC PROPERTIES
This section discusses how several properties vary across periods and down groups:
Atomic radii
Ionization energy
Electron affinity
Electronegativity

39. Atomic Radii
The boundaries of an atom are fuzzy, and an atom’s radius can vary under different conditions.
To compare different atomic radii, they must be measured under specified conditions.
Atomic radius may be defined as one-half the distance between the nuclei of identical atoms that are bonded together.

40. Atomic Radius

41. Atomic Radii: Periodic Trends
Atoms tend to be smaller the farther to the right they are found across a period.
The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus, which attracts electrons toward the nucleus.

42. Atoms tend to be larger the farther down in a group they are found.
The trend to larger atoms down a group is caused by electrons occupying higher main energy levels.

44. Sample Problem
Of the elements magnesium, Mg, chlorine, Cl, sodium, Na, and phosphorus, P, which has the largest atomic radius? Explain your answer in terms of trends of the periodic table.

45. Solution
Sodium has the largest atomic radius
All of the elements are in the third period. Of the four, sodium has the lowest atomic number and is the first element in the period. Atomic radii decrease across a period.

46. Sodium (Na), for example, easily loses an electron to form Na+.
Ions
An ion is an atom or group of bonded atoms that has a positive or negative charge.
Sodium (Na), for example, easily loses an electron to form Na+.
A cation is a positive ion. An anion is a negative ion.

47. Ion

48. The electron cloud becomes smaller.
Ionic Radii
The formation of a cation by the loss of one or more electrons always leads to a decrease in atomic radius.
The electron cloud becomes smaller.
The remaining electrons are drawn closer to the nucleus by its unbalanced positive charge.
The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius.

49. Comparing Cations and Anions 75095

50. Cationic and anionic radii decrease across a period.
The electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same main energy level.
The outer electrons in both cations and anions are in higher energy levels as one reads down a group.
There is a gradual increase of ionic radii down a group.

52. Ionization Energy
Any process that results in the formation of an ion is referred to as ionization.
The energy required to remove one electron from a neutral atom of an element is the ionization energy, IE (or first ionization energy, IE1).

53. In general, ionization energies of the main-group elements increase across each period.
This increase is caused by increasing nuclear charge. A higher charge more strongly attracts electrons in the same energy level.
It is much easier to remove electrons from metals than from nonmetals (nonmetals have higher ionization energies).

54. Among the main-group elements, ionization energies generally decrease down the groups.
In larger atoms the electrons are farther from the nucleus and are more easily removed.

55. Ionization

57. Electron affinity generally increases across periods.
The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity.
Electron affinity generally increases across periods.
Increasing nuclear charge along the same sublevel attracts electrons more strongly

58. Electron affinity generally decreases down groups.
The larger an atom’s electron cloud is, the farther away its outer electrons are from its nucleus.

59. Electron Affinity

60. Valence Electrons
Chemical compounds form because electrons are lost, gained, or shared between atoms.
The electrons that interact in this manner are those in the highest energy levels.

61. The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons.
For the main group elements, these are the electrons in the outermost s and p orbitals.

62. Group
Valence e- 1 15 5 2 16 6 13 3 17 7 14 4 18 8
So for groups 13-18, the number of valence electrons is the group number minus 10.

63. Valence Electrons

64. Valence electrons hold atoms together in chemical compounds.
Electronegativity
Valence electrons hold atoms together in chemical compounds.
In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another.

65. Electronegativity is the ability of an atom in a chemical compound to attract electrons from another atom in the compound.
Electronegativities tend to increase across periods, and decrease or remain about the same down a group.

66. Electronegativity

67. The most electronegative elements are oxygen, nitrogen, and the halogens.
F = 4.0 O = 3.5 N = 3.0
Alkali and alkaline-earth elements are least electronegative:
Fr, Cs = 0.7; Rb = 0.8; Na, Ra, Ba = 0.9

68. Achieving Octets
Main group elements tend to gain, lose, or share electrons so as to get a set of 8 electrons in their highest occupied energy level (an octet).
Groups 1 and 2 can get octets by losing electrons to form cations:
Na: 1s22s22p63s1 →
1s22s22p6 (Na+) + 1e-
Ca: [Ar]4s2 →
[Ar] (or [Ne]3s23p6) (Ca2+) + 2e-

69. Groups 16 and 17 can get octets by gaining electrons to form anions:
Br: [Ar]4s24p5 + 1e- →
[Ar]4s24p6 (or [Kr]) (Br-)
O: 1s22s22p4 + 2e- →
1s22s22p6 (or [Ne]) (O2-)
Noble gases already have octets, so they are unreactive.