1. Chapter 6: The Periodic Table General Chemistry

2. Objectives
Periodicity of physical and chemical properties relates to atomic structure and led to the development of the periodic table.
The periodic table displays the elements in order of increasing atomic number.
Explain the relationship of an element’s position on the periodic table to its atomic number and mass.

3. Objectives
Use the periodic table to identify metals, nonmetals, metalloids, families (groups), periods, valence electrons, and reactivity with other elements in the table.
Relate the position of an element on the periodic table to its electron configuration.
Identify trends on the periodic table (ionization energy, electronegativity, electron affinity, and relative size of atoms and ions).

4. Review/Link to Previous Learning
In Chapter 4, we learned about electrons configurations of elements.
Discovered there is a pattern of electron configurations on the Periodic Table.
Are there other patterns on the Periodic Table? (yes)
In Chapter 5 we will learn how the Periodic Table is organized.

5. Collections Do you like to play cards? Do you have a stamp, baseball
card, or comic book collection?
How do you organize your collection?

6. Attempts at Organizing Elements
Early scientists knew about some properties of elements.
Is there a characteristic of elements that can organize them?

7. JOHAN DOBEREINER( )

8. Dobereiner’s Triads THE LAW OF TRIADS:
The atomic mass of the middle element of the triad
is equal to the mean of the atomic masses of the
other two elements.
EXAMPLE: Lithium Atomic Mass of 7
Sodium Atomic Mass of 23
Potassium Atomic Mass of 39
According to Dobereiner’s Law, the atomic mass of sodium
Should equal the arithmetic mean of lithium and potassium.
(7+39)/2 = 23, which is the mass of sodium.

9. Problems with Dobereiner’s
Law of Triads.
1) All the elements known at that time could not be
arranged in triads.
2) The law did not work for very low or very high
massed elements such as F, Cl, and Br.
3) As techniques improved for measuring atomic
masses accurately, the law became obsolete.
Dobereiner’s research made chemists look at groups of
elements with similar chemical and physical properties.

10. JOHN A.R. NEWLANDS ( )

11. Newland’s Law of Octaves
When placed in increasing order of their atomic
masses, every eighth element showed similar physical
and chemical properties.
Li Be B C N O F
Na Mg Al Si P S Cl
K Ca

12. Newland’s Law of Octaves
Problems with
Newland’s Law of Octaves
1) It was not valid for elements that had atomic
masses higher than Ca.
2) When more elements were discovered
(Noble gases) they could not be accommodated in his
table.
However, the modern periodic table does draw from the
concept of periods of eight.

13. DMITRI MENDELEEV ( )

14. Julius Lothar Meyer (1830-1895)

15. Mendeleev and Meyer
Published nearly identical schemes for classifying elements
Arranged elements by increasing atomic mass
Mendeleev generally given more credit
Published first
More successful at demonstrating value of table
Predicted discovery of new elements, properties of new elements

16. Properties of Some Elements Predicted By Mendeleev

17. Mendeleev’s Table: the first periodic table of the elements.
He arranged the table so that elements in the same column
have similar properties.

18. Problems with Mendeleev’s Table:
1) The positions of isotopes could not be
accommodated within the table.
2) In order to make the elements fit the requirements,
Mendeleev was forced to put an element of slightly
higher atomic weight ahead of one of slightly lower
atomic weight.

19. Henry Moseley (1887-1915) Developed concept of atomic number
amount of positive charge in the nucleus
Later determined that arranging periodic table according to increasing atomic number eliminated problems seen in Mendeleev’s table

20. Why is it the “periodic” table?
Periodic Law: when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern

21. Study Buddy Review
Describe the contribution each person below made to the development of the periodic table:
Johan Dobereiner
John Newland
Dmitri Mendeleev
Julius Meyer
Henry Moseley

22. Parts of the Periodic Table

23. Parts of Periodic Table
Groups/families: vertical columns
Alkali metals: 1A
Alkali earth metals: 2A
Boron, carbon families
chalcogens (oxygen family).
pnictogens (nitrogen family)
Halogens (fluorine family): 7A
Noble gasses: 8A/0

24. Horizontal rows are called periods
There are 7 periods

25. The elements in the A groups are called the representative elements
outer s or p filling
8A0 1A 2A 3A 4A 5A 6A 7A

26. Parts of Periodic Table
Metals: left of staircase
Luster, malleable, conduct, ductile
Nonmetals: right of staircase
Dull in appearance, nonconductor, brittle
Metalloids: elements adjacent to staircase (except Al, Po)
Some properties of both metals and nonmetals

27. The group B are called the transition metals
These are called the inner transition metals and they belong here

28. Study Buddy Review Identify the follow parts of the periodic table:
Halogens
family
Alkali metals
Metals
Inner transition metals
Noble gases
Metalloids
Period

29. Periodic Properties of Elements

30. Periodic Trends Atomic Radius Ionic Radius Ionization Energy
Electron Affinity
Electronegativity

31. Atomic Radius }
Radius
Atomic Radius = half the distance between two nuclei of a diatomic molecule.

33. Trends in Atomic Radius
Influenced by three factors:
1. Charge on nucleus
More charge pulls electrons in closer.
2. Energy Level
Higher E level is further away from nucleus
3. Shielding effect
The number of electrons between electrons and nucleus affects the pull felt by the outer electrons

34. Atomic Radius Group trendsH Li
As we go down a group...
each atom has another energy level
so the atoms get bigger. Na K Rb

35. Atomic Radius Periodic Trends
As you go across a period, the radius gets smaller.
The increasing number of protons in the nucleus pulls the electrons in more tightly Na Mg Al Si P S Cl Ar

36. Atomic Radius

37. Ionic Size Ion: electrically charged atom
Cation: positively charged ion
Anion: negatively charged ion
Ions aren't the same size as the neutral atoms they come from.
Compare the sizes of sodium and chloride ions with the sizes of sodium and chlorine atoms.

38. Positive ions are smaller than the atoms they come from.
The sodium ion loses a whole layer of electrons, and the remaining 10 electrons are being pulled in by the full force of 11 protons.
Negative ions are bigger than the atoms they come from.
Although the electrons are still all in the 3-level, the extra repulsion produced by the incoming electron causes the atom to expand. There are still only 17 protons, but they are now having to hold 18 electrons.

40. Study Buddy Review-A.R., I.R.
Describe the pattern for atomic radius
As you move across a period
As you move down a column
What charge does a cation have?
What charge does an anion have?
Which is larger than its parent atom, a cation or an anion?

41. First Ionization Energy
Ionization energy is the energy required to remove the first electron from an atom of an element
Elements want to have the e- configuration like that of a noble gas (filled)
Column 1A elements have need to LOSE one electron to have noble gas configuration so it is EASY to remove electron
Column 7A element need to GAIN one electron to have noble gas configuration, so it is HARD to remove electron

42. First Ionization Energy vs. Atomic Number

43. Ionization Energy

44. Ionization Energy 
As you move across a period ionization energy increases…
Elements on left of table want to lose electrons to have full energy level (requires low energy to remove electron)
As you move down a group ionization energy decreases…
The electrons that are further away from the nucleus are easier to remove and thus require less energy to remove

45. Successive Ionization Energies
more than one electron can be removed from atoms
Second Ionization energy: when a second electron is removed from an atom that has already lost one electron
Third Ionization energy: when a third electron is removed from an atom that has already lost two electrons

46. Sucessive Ionization Energies
Symbol First Second Third
HHeLiBeBCNO F Ne

47. Relationship Between Common Charge and I.E.
Consider Beryllium:
Electron config: [He] 2s2
Low energy to remove 1st and 2nd electrons
MUCH higher energy to remove 3rd electron because it would be removed form a noble gas configuration

48. Study Buddy Review-I.E. What is ionization energy?
Describe the pattern for ionization energy as you
Move down a family
Move across a row
What does “first” ionization energy mean?

49. Electron Affinity Electron affinity is:
the energy change associated with the adding an electron to a gaseous atom
the more attraction for an electron the energy is released when the atom gains the electron
Released energy is negative (-350 kJ)

50. Electron Affinity

51. Electron Affinity General Trend:
Halogens (s2p5 configurations) are most negative electron affinities. They are most likely to want to gain electrons to obtain noble gas configuration
As you go down a family, electron affinity is less negative (harder to gain electrons with increasing atomic size)

52. Study Buddy Review-E.A. What does it mean when an energy is negative?
Which elements generally have a very negative electron affinity?

53. Electronegativity
Electronegativity: the tendency for an atom to attract electrons to itself when it is chemically combined with another element.
As you move down a group, electronegativity decreases
As you move across a period, electronegativity increases
Which element is the MOST electronegative?

54. Electronegativity

55. Study Buddy Review-Electroneg
Define electronegativity.
Describe the pattern for electronegativity as you
Move down a group
Move across a period
Which element is the most electronegative?

56. Resources http://www.chemguide.co.uk/atoms/properties/atradius.html
Jeanette Boles
Tina Lula
Dr. Stephen L. Cotton, Charles Page High School