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.
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.

6. 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.

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

8. 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.

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

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

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

13. 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).

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

15. 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.

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

17. The elements arranged vertically in the periodic table share chemical properties.
They are also organized in horizontal rows, or periods.
The length of each period is determined by the number of electrons that can occupy the sublevels in that period.

18. The period (energy level) 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 (energy level) of the periodic table.

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

20. 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.

21. 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.

22. Alkaline-Earth Metals – The elements in Group 2 of the periodic table (beryllium, magnesium, calcium, strontium, barium and radium).
Slightly less reactive than the alkali metals.
Still too reactive to be found in nature as free elements.

23. 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.

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

25. 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).

26. Halogens – the elements of Group 17 (fluorine, chlorine, bromine, iodine and astatine).
Fluorine and chlorine are gases. Bromine is a liquid and iodine is a solid.

27. 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)

28. 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.

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

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

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

32. Atomic Radii
Ideally, the size of an atom is defined by the edge of its orbital.
However, this boundary and there is no set outer edge.
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.

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

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

37. 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.

38. Group Trends
There is an increase down a group.
This trend is because as electrons are added for each row of the periodic table, they are further from the nucleus.
They are in higher energy levels.
Some exceptions do occur.

39. 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 has the largest atomic radius and why?

40. 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.

42. 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.

43. 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.

44. 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.

45. Classwork
Review practice problems 1and 2 on page 142.

46. 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.

48. Period Trends
Among the elements of each period, the halogens (Group 17) gain electrons most readily.
As the number of positive protons increase across a period, it is easier to add electrons. .

49. Group Trends Electron affinities within groups generally decrease.
The electron being added is further from the positive nucleus so the attraction to the is less.

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

51. Electronegativity scale of 0 to 4 0: least electronegative
4: most electronegative

52. 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.

53. 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).

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

55. 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.

56. Cations and Anions When atoms lose electrons to form cations,
they always become smaller.
When atoms gain electrons to become anions, they always become larger.

57. 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. (Same row of the periodic table)
Valence electrons – the electrons available to be lost, gained or shared in the formation of chemical compounds.

58. Homework End of chapter problems – page 157-158
Questions 32, 34, and 37.
Due: