1. Chapter 5
The Periodic Law

2. Objectives: Describe the periodic tables of Moseley and Mendeleev.
Identify the various families of elements on the periodic table.
State the trends in atomic radius, ionization energy, electron affinity and ion size with a group or period on the periodic table.
Identify the relationship between these trends and the structure of the atom.

3. History of the Periodic Table
Chapter 5
Section 1
History of the Periodic Table

4. By 1860, more than 60 elements had been discovered.
Chemists needed to learn the properties of these elements as well as those of the compounds they formed.
At the time there was no method for accurately determining an element’s atomic mass.
Things were disorganized and confusing among chemists.

5. Dmitri Mendeleev
Mendeleev hoped to organize the elements according to their properties.
He arranged the elements in order of increasing atomic mass and noticed certain similarities in their chemical properties would appear at regular intervals.
In 1869 Mendeleev created the first periodic table grouping the elements with similar properties.

7. Henry Moseley
In 1911 Moseley revised the periodic table of Mendeleev by arranging the elements in order of increasing atomic number (number of protons) and not atomic mass.
The elements were still grouped by similar physical and chemical properties.

8. Modern Periodic Table
The periodic table has undergone extensive changes since Mendeleev’s time.
More than 40 new elements have been discovered or synthesized.

9. Periodic Table – an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.

10. The elements arranged vertically in the periodic table share chemical properties.
Example:
F, Cl, Br and I all react in a similar manner.

11. Later additions to the periodic table include:
Noble Gases ( ) – group 18 elements that are characterized by their relative un-reactivity.
The Lanthanides (early 1900’s) – the 14 elements with atomic numbers from 58 (cerium, Ce) to 71 (lutetium, Lu)

12. The Actinides – the 14 elements with atomic numbers from 90 (thorium, Th) to 103 (lawrencium, Lr).
The lanthanides and actinides belong in periods 6 and 7 of the periodic table.
To save space and to group them together they are set off below the main portion of the periodic table.

13. Periodic Table Elements are classified in three major groups:
Metals
Nonmetals
Metalloids
Use the periodic table to distinguish the classes of elements

15. Physical Properties of Metals
Malleability
Ductility
Luster
Heat conductors and electrical conductors
Solids
Ex. Iron (Fe), tin (Sn), zinc (Zn), and copper (Cu)

16. Properties of Nonmetals
Dull in appearance
Brittle
Do not conduct electricity
Ex. Carbon (C), Oxygen (O)
and Sulfur (S)
Solids, liquids or gases

17. The metalloids divide the metals from the nonmetals.
They are mostly brittle solids with some properties of metals and some of nonmetals.
The electrical conductivity falls between the metals and nonmetals.
The metals of the p block are reactive enough to be found in nature only as compounds and not free elements (except bismuth).

18. Metalloids Properties of metals and nonmetals
Ex. Silicon (Si) and Germanium (Ge)
Common in the computer
industry

19. Summary – Section 5.1 Recognize the work of Mendeleev and Moseley.
How the modern periodic table is arranged with respect to the elements.
Know the three sets of elements added to the periodic table after Mendeleev.

20. Electron Configuration
Chapter 5
Section 2
Electron Configuration
and the
Periodic Table

22. In the first period, the 1s sublevel is filled.
The 1s sub level can hold a total of two electrons.
Therefore, the first period consists of two elements - hydrogen and helium.
The second period can held 8 electrons and therefore consists of 8 elements.
The third period is similar to the second period – 8 elements.

23. When you get to the fourth and fifth rows you have to include the transition metals.
There are 10 transition metals in each of these rows.
The total number of elements in the fourth and fifth rows are therefore 18 elements.

24. When you get to the sixth and seventh rows you have to include the lanthanides and actinides.
There are 14 elements in each of these groups.
The total number of elements in the six and seventh rows can therefore include 32 elements.

25. The periodic table is divided into 4 blocks

26. The period of an element can be determined from the element’s electron configuration.
For example, arsenic – As, has the electron configuration of:
[Ar]3d104s24p3
The 4 in 4p3 indicates that arsenic is in the fourth period of the periodic table.

27. Without looking at the periodic table give the period number for the following electron configuration:
[Xe]6s2

28. The s-Block Elements
The elements of the s block are chemically reactive metals.
The Group 1 metals are more reactive than the Group 2 metals.
The outermost energy level in an atom of the Group 1 elements contain a single electron.
The ease with which the single electron is lost make the Group 1 metals extremely reactive.

29. Alkali Metals – The elements in Group 1 of the periodic table (lithium, sodium, potassium, rubidium, cesium and francium).
Because the alkali metals are so reactive they are not found in nature as free elements.
They combine vigorously with most nonmetals.
They react strongly with water to produce hydrogen gas.

30. Alkaline-Earth Metals – The elements in Group 2 of the periodic table (beryllium, magnesium, calcium, strontium, barium and radium).
Atoms of the alkaline-earth metals contain a pair of electrons in their outermost shell.
Slightly less reactive than the alkali metals.
Still too reactive to be found in nature as free elements.

31. Classwork
Review the practice problems 1 and 2a-d on page 133.

32. The d-Block Elements
The elements of the d block are known as the transition metals.
They are good conductors of electricity and have a high luster.
They are typically less reactive than the alkali and alkaline-earth metals.
Palladium, platinum and gold are among the least reactive of all elements.

33. Classwork
Review the practice problems 1 and 2a-b on page 136.

34. The p-Block Elements
The elements of the p-block and s-block are known as the main-group elements.
The p-block consists of the elements in Groups (except helium).
The properties of the p-block vary greatly.
The p-block contains metals (Al), metalloids (Si) and nonmetals (Br).

35. The electron configuration of bromine is:
[Ar]3d104s24p5

36. Halogens – the elements of Group 17 (fluorine, chlorine, bromine, iodine and astatine).
The halogens are the most reactive nonmetals.
They react vigorously with most metals to form compounds known as salts (NaCl).
Fluorine and chlorine are gases. Bromine is a liquid and iodine is a solid.

37. The Group 18 elements (Noble Gases) undergo few chemical reactions.
This stability results from the gases electron configuration.
Their highest occupied levels are completely filled with electrons (octet).
An atom’s electron configuration governs the atom’s chemical properties.

38. Classwork
Review the practice problems 1a-b and 2a-b on page 138.

39. Homework Section Review – page 139 Questions 1, 2, 4 and 5
End of chapter problems – page
Questions 27, 28 and 29
Due:

40. Electron Configuration
Chapter 5
Section 3
Electron Configuration
and the
Periodic Properties

41. So far you have learned that the elements are arranged in the periodic table according to their atomic number.
There is also a correlation between the arrangement of the elements and their electronic configuration.
We will look at the relationship between the electron configurations and the periodic trends of the elements.

42. Atomic Radii
Ideally, the size of an atom is defined by the edge of its orbital.
However, this boundary varies under different conditions.
One way to express the atomic radius is to measure the distance between the nuclei of two identical atoms that are bonded together, then divide this distance by two.

43. Atomic radius may be defined as one-half the distance between the nuclei of identical atoms that are bonded together.

44. Period Trends
There is a gradual decrease across a row.

45. There is a gradual decrease across a row.
The trend to smaller atoms across a period is caused by increasing positive charge of the nucleus. Adding of protons or increasing atomic number.
The electrons are pulled closer to the nucleus.
The increase pull results in a smaller atomic radius.

46. Group Trends
There is an increase down a group.

47. There is an increase down a group.
As electrons are added to sublevels in higher energy levels located further from the nucleus, the size of the atoms increase.
An exception is between aluminum (radius – 143 pm) and gallium (radius – 135 pm).
This is due to the gallium being proceeded by the 10 d-block elements.
The nuclear charge is considerably higher.

48. Problem
Of the elements magnesium-Mg, chlorine-Cl, sodium-Na, and phosphorus-P, which has the largest atomic radius and why?
Of the elements calcium-Ca, beryllium-Be, barium-Ba and strontium-Sr, which as the largest atomic radius and why?

49. Ionization Energy
Ionization energy (IE) – the energy required to remove one electron from a neutral atom of an element.
A + energy A+ + e-
An ion is an atom or group of bonded atoms that has a positive or negative charge.

52. Period Trends
In general, ionization energies of the elements increase across a period.
Group 1 elements – have the lowest ionization energies. Therefore they lose electrons most easily. Very reactive.
Group 18 elements - have the highest ionization energies . They do not lose electrons easily. Very low reactivity.

53. The increase is due to increasing nuclear charge (more protons going across a period).
A higher positive charge more strongly attracts electrons in the same energy level.
Therefore, it is tougher to remove an electron from an atom.
Increasing nuclear charge is responsible for both an increasing ionization energy and decreasing atomic radius across a period.

54. Group Trends Ionization energies generally decrease down a group.
Electrons removed from atoms of the elements down a group are farther from the nucleus.
Also, the electrons from the lower energy levels shield the outer electrons.
Therefore, they are removed more easily.

55. Classwork Review practice problems 1and 2 on page 142.
Review practice problems 1 (a-c) on page 146.

56. Electron Affinity Neutral atoms can also acquire electrons.
Electron Affinity – the energy change that occurs when an electron is acquired by a neutral atom.
A + e A- + energy
The quantity of energy released is represented by a negative number.

58. Period Trends
Among the elements of each period, the halogens (Group 17) gain electrons most readily.
The ease with which halogen atoms gain electrons is a major reason for the high reactivities of the Group 17 elements.
As electrons add to the same p sublevel of atoms, electrons affinities increase.

59. Period Trends
An exception to this trend occurs between Group 14 and 15.
Compare the electron affinities of carbon ([He] 2s2 2p2 ; ) and nitrogen
[He] 2s2 2p3; 0).
Adding an electron to nitrogen is more difficult because it forces the electron to pair with another electron.

60. Group Trends
Trends for electron affinities within groups are not as regular as trends for ionization energies and atomic radii.

61. Cations and Anions Cation (A+) – a positive ion that results with
loss of an electron.
Anion (A-) – a negative ion that results with
the gain of an electron.

62. Valence Electrons
Chemical compounds form because electrons are lost, gained or shared between atoms.
The electrons that interact are those in the highest energy levels.
Valence electrons – the electrons available to be lost, gained or shared in the formation of chemical compounds

63. Valence electrons are often located in incompletely filled main-energy levels.
For example, the electron lost in Na to form Na+ is from the 3s sublevel.
For the main group elements, the valence electrons are the electrons in the outermost s and p sublevels.
The inner electrons are in filled energy levels and are held to tightly to be involved.

65. Electronegativity
Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons.
In many compounds the negative charge of the valence electrons are not shared evenly between atoms.
The uneven concentration of charge has an effect on the chemical properties of a compound.

67. Period Trends Electronegativities tend to increase across a period.
The Group 1 and 2 elements have the lowest electronegativities.
Groups 16 and 17 have the highest electronegativities.

68. Group Trends Electronegativities tend to decrease down a group.
The combination of these trends in electronegativities results in the highest values belonging to the elements in the upper right of the periodic table (fluorine).

69. Problem
Of the elements gallium-Ga, bromine-Br, and calcium-Ca, which has the highest electronegativities and why?

70. Homework End of chapter problems – page 157-158
Questions 32, 34, 36 a-b, 37 and 46.
Due: