1. The periodic law
Chapter 5

2. Why do we need a table?
To organize the elements
To show trends

3. Periodic
A repeating pattern

4. Mendeleev’s table 1869 – Dmitri Mendeleev – Russian
Arranged the elements in order of increasing mass and noticed that chemical properties were periodic
Put the elements into groups according to properties

5. Mendeleev vs. Meyer
1860s Mendeleev and German Lothar Meyer each made an eight column table.
Mendeleev left some blanks in his table in order for all the columns to have similar properties – he predicted elements that hadn’t been discovered yet.

6. Why similar properties?
Why did they group according to properties and mass and not atomic number or number of outer level electrons?

7. Mendeleev’s predictions
Germanium
Mendeleev’s predictions
Actual element
Atomic mass = 72
Atomic mass = 72.60
High melting point
Melting point = 958 °C
Density = 5.5 g/cm3
Density = 5.36 g/cm3
Dark gray metal
Gray metal
Mendeleev’s blank spots and his ability to predict future elements helped his table win acceptance.

8. Mendeleev’s table Elements arranged in order of increasing mass.
Properties are repeated in an orderly, periodic, fashion.
Mendeleev’s periodic law – the properties of the elements are a periodic function of their masses.

10. Mass mistakes?
In order for Mendeleev to arrange his elements by properties, he had to put tellurium and iodine in the wrong order.
He explained this by assuming that their masses hadn’t been measured very accurately.

11. More mass mistakes? Nickel and cobalt Argon and potassium
Better mass measurements just confirmed the discrepancy

12. Explanation 1913 – Henry Moseley
X-ray experiments revealed the atomic number was the number of protons
Modern periodic law – the properties of the elements are a periodic function of their atomic numbers

13. Modern 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.

14. Noble gases Not discovered on Earth until 1894 - 1900.
Group 18 was added to the table

15. Lanthanides Hard to separate All have similar properties
Added to the table in the early 1900s

16. Actinides
Discovered later
Also all have similar properties

17. Periodicity
Elements in the same group (column) have similar properties.

18. Chemical properties of an element
Are governed by the electron configuration of an atom’s highest energy level

19. Period length
Determined by the number of electrons than can occupy the sublevels being filled in that period.
Table 5-1

20. Full periodic table
Table with f-block in place

21. 1st period 1s sublevel being filled
1s can hold 2 electrons, so there are 2 elements in the 1st period.

22. 2nd and 3rd periods 2s and 2p or 3s and 3p being filled
s and p sublevels can hold 8 total, so there are eight elements in these periods

23. 4th and 5th periods Add d sublevels, which can hold 10 electrons
Need to fill 4s, 3d, and 4p – 18 electrons
18 elements in each period

24. 6th and 7th periods Add f-block, which holds 14 electrons
Fill 6s, 5d, 4f, 6p
Need 32 electrons
32 elements in each period

25. Figure 5-5
Shows blocks

26. Electron configurations
Elements in columns 1, 2, and have their last electron added in an s or p orbital.
Elements in columns 3-12 have their last electron added in a d level.

27. The s-block elements: Groups 1 and 2
Chemically reactive metals
Group 1
Have 1 electron in outer s orbital
Coefficient represents period
Row 2: 2s1, Row 3: 3s1, etc. (ns1)
Group 2
Have 2 electrons in outer s orbital
Row 2: 2s2, Row 3: 3s2, etc. (ns2)

28. Alkali metals Metals in group 1 Have silvery appearance
Soft enough to cut with a knife
Not found alone in nature
React violently with nonmetals
Melting point decreases as you go down the table

29. Alkaline-earth metals
Group 2
Harder, denser, and stronger than alkali metals
Higher melting points than alkalis
Less reactive
Not found alone in nature

30. Hydrogen and helium Hydrogen Helium
Located above group 1 because of its electron configuration
Not really in group 1, because its properties don’t match
Helium
Has an electron configuration like group 2 elements
In group 18 because it is unreactive

31. Discuss
Page 133
Sample problem 5-1 and practice problems

32. Discuss
Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [Rn] 7s1 is located.
Group 1, 7th period, s block
Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [He]2s2 is located.
Group 2, second period, s block

33. d-block elements: Groups 3-12
End in d1 to d10.
Coefficients are one less than the period
Example: Fe is in the 6th column of transition elements in the 4th period, ends in 3d6

34. Transition elements Groups 3-12 Typical metallic properties
Good conductors
High luster
Less reactive than alkalis and alkaline-earths
Some are unreactive enough to appear in nature

35. p-block elements: groups 13-18
End in p1 to p6.
Coefficients are the same as the period
ns2np1
Always have a full s-sublevel

36. p-block elements Properties vary greatly
Includes all nonmetals except hydrogen and helium
Solids, liquids and gases
Includes all the metalloids
Between metals and nonmetals
Brittle solids
Semiconductors – can conduct under certain conditions
Includes some metals
Less reactive than alkalis and alkaline-earths

37. Halogens
Group 17
Most reactive nonmetals
Form compounds called salts

38. f-block elements Lanthanides and actinides
Endings are f1 to f14
Coefficients are two less than the period
All actinides are radioactive
Those after neptunium are synthetic

39. Discuss
Sample problems and practice problems on pages 136, 138, and 139
With your group first, then join with another group.
Do you have any questions?

40. Atomic radius
Ideally, the distance from the center of the atom to the edge of it’s orbital.
But, atoms are “fuzzy”, not clearly defined.
Defined as one-half the distance between the nuclei of identical atoms that are bonded together.

41. Period trends – see figure 5-13
As we move from left to right across the table, we gain protons.
There is a greater positive charge on the nucleus.
This greater charge pulls harder on the outer electrons, pulling them in closer.
The atom gets smaller.

42. Group trends
As we move down the table, the principle quantum number increases.
When the principle quantum number increases, the electron cloud gets bigger.
The size of the atoms gets bigger.

43. Discuss
Which of the elements Li, Rb, K, and Na has the smallest atomic radius? Why?
Li, it is highest on the table
Which of the elements Zr, Rb, Mo, and Ru has the largest atomic radius? Why?
Rb, it is farthest to the left on the table

44. Ion
An atom or group of bonded atoms that has a positive or negative charge

45. Ionization
Any process that makes ions

46. Ionization energy (IE)
First ionization energy (IE1) – the energy required to remove the most loosely held electron.
Measured in kJ/mol

47. Ionization energy – see figure 5-15
Experimentally determined.
From isolated atoms in the gas phase
Tends to increase as you move across a row from left to right
Why group 1 is most reactive
Caused by higher charge
Tends to decrease as you move down a column
Electrons are farther from nucleus
Shielding from inner electrons

48. Other Ionization Energies – see Table 5-3
Energy required to remove other electrons from positive ions.
IE2, IE3, etc
Get higher as you remove more electrons
Less shielding

49. Noble Gases Have High ionization energies
When a positive ion of another element reaches a noble gas configuration, its ionization energy goes up.
Example: When K loses one electron, it has Ar’s electron configuration
This makes it stable
Its IE2 is much higher than its IE1

50. Discuss
State in words the general trends in ionization energies down a group and across a period of the periodic table.

51. Electron affinity
The energy change that occurs when an electron is gained by a neutral atom
Most atoms release energy
Represented by a negative number
Some atoms gain energy
Represented by a positive number
These ions will be unstable
KJ/mol

52. Period trends – see figure 5-17
Group 17 has most negative electron affinity.
Tends to get more negative (release more energy) as we move to the right
Exceptions:
groups with full or half-full sublevels are more stable

53. Group trends Not as regular
Usually, electrons add with greater difficulty as we move down

54. Adding additional electrons
Second electron affinities are all positive because it is more difficult to add electrons to a negative ion.
If a noble gas configuration has been reached, it is even more difficult.

55. Discuss
State in words the general trends in electron affinities down a group and across a period of the periodic table.

56. Ionic Radii Cation – a positive ion Anion – a negative ion
Ionic radius smaller than atomic radius
Anion – a negative ion
Ionic radius is larger

57. Period Trends – see figure 5-19
Metals form cations by losing electrons
Ions are smaller
Radius decreases as we move across
Nonmetals form anions by gaining electrons
Ions are larger

58. Group trends
Ionic radius increases as you go down the table

59. Valence electrons
Available to be lost, gained or shared in the formation of chemical compounds
In highest energy levels
For s-block, the group number is the same as the number of valence electrons
For the p-block, the group number is 10 more than the number of valence electrons

60. Electronegativity
The measure of the ability of an atom in a compound to attract electrons
The atom with higher electronegativity pulls the electrons closer to itself

61. Electronegativity trends (figure 5-20)
Increases left to right across the rows
Decreases down the columns

62. Discuss
Explain why elements with high (more negative) electron affinities are also the most electronegative.

63. d- and f-block elements
Properties vary less and with less regularity than others
Atomic radii
d-block
Usual patterns
f-block (unusual)
Increase across periods
Decrease down groups

64. d- and f-block elements
Ionization energy
Increase across periods
d-block increases down groups (unusual)
f-block decreases down groups
Ionic radii
Cations have smaller radii
Electronegativity
d-block follows normal rules
f-block all have similar electronegativities

65. Discuss
Among the main-group elements, what is the relationship between group number and the number of valence electrons?
In general, how do the periodic properties of the d-block elements compare with those of the main-group elements?

66. Prelab notes Precipitate – solid that falls out of a solution
The formation of a precipitate indicates there has been a chemical change.
This means that there were ions present that were free to react.