1. The Periodic Table
Chapter 6

2. Introduction
The periodic table represents an organizing principle which allows the prediction of the properties of each element based on their position in the periodic table.
The elements of the periodic table can be classified into different categories.
There are trends that appear in the periodic table that allow us to make predictions about an atoms size, ionization energy, and electronegativity.

3. Organizing Elements (Section 6.1)
Searching for an Organizing Principle
Mendeleev’s Periodic Table
The Periodic Law
Metals, Nonmetals, and Metalloids

4. Defining the Periodic Table
Periodic Table: An arrangement of elements in which the elements are separated into groups based on a set of repeating properties.
· A periodic table allows you to easily
compare the properties of one element
(or group of elements) to another
element (or group of elements).

5. I.) Searching for an Organizing Principle
Copper, gold, and silver have been known for thousands of years.
Only 13 elements were known of by the year 1700.
Between five more elements were discovered.
Problem was how to ID new elements and how many new ones there actually were.

6. The Dobereiner System Organized known elements into triades.
Recognized a relationship between atomic weights and chemical properties
Triade: A set of three elements with similar properties (ex. Cl, Br, I)
Not all the known elements could be grouped into triades
J.W. Dobereiner
German Chemist

7. The Dobereiner Triades
1 element in each trade
tended to have properties
with values that fell midway between the other two.
Here we see atomic
and mass numbers

8. II.) Mendeleev’s Periodic Table
From many other systems were proposed but none gained wide acceptance.
Mendeleev created his table while working on a text book for his students
Beat a competitor because he was better able to explain the table’s usefulness.
Dmitri Mendeleev
Russian Chemist/Teacher

9. Arranged the elements
in his table in order of
increasing atomic mass.
There was a close
match between the
predicted properties of
unknown elements and
the actual properties of
the elements.
This organizational
method had its
problems and does not
account for all elements
(ex. atomic masses for
Te and I).

10. III.) The Periodic Law
There was a problem with organizing elements by atomic masses.
Organizing elements by increasing atomic number was more useful.
Moseley determined the atomic number for the known elements at the time
Henry Moseley
British Physicist

11. Expression of the Periodic Law
When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

12. The Modern Periodic Table
Horizontal Rows = Periods
Vertical columns = Groups

13. Periods in the Periodic Table
These are 7 rows extending horizontally across the periodic table.
Period 1 = 2 elements
Period 2/3 = 8 elements
Period 4/5 = 18 elements
Period 6/7 = 32 elements
Each period corresponds to a principle energy level.
More elements in the higher periods because there are more orbitals in the higher energy levels
Properties of elements within elements within a period change as you move from left to right.

14. Groups in a Periodic Table
These are the 18 columns that run up and down the periodic table.
There are 3 different ways that the groups are numbered.
These groups also possess names.
Elements within a group in the periodic table have similar physical and chemical properties.

15. IV.) Metals, Nonmetals, & Metalloids
We saw how the periodic table can be divided into 7 periods and 18 groups.
We can also divide the table into three broad classes based on the general properties of the elements.
Metals
Nonmetals
Metalloids
Across the periods, the properties of elements become less metallic and more nonmetallic.

16. The Three Classes of Elements

17. Metals This is the most numerous class. Characteristics of metals:
Good conductors of heat and electric current
Possess luster and sheen
Solids at room temperature (except Hg)
Ductile (i.e. can be drawn into wires)
Most are malleable (i.e. can be hammered into thin sheets)

18. Examples of Metals Copper is ductile. Aluminum is malleable and has
luster and sheen.

19. Nonmetals Less numerous than the metals
There is greater variation in the characteristic of nonmetals.
One general characteristic: They are not metals.
Poor conductors of heat and electric current (carbon is an exception)
Solids tend to be brittle.
Most nonmetals are gases at room temperature, a few are solids, and 1 is a liquid (Br).

20. Examples of Nonmetals Chlorine is a gas Bromine is a liquid
Carbon is a solid and is
a good conductor of
electricity.

21. Metalloids Least numerous elements
Have properties that are similar to those of metals and nonmetals.
The behavior is often controlled by changing the conditions.

22. Examples of Metalloids
Silicon is not a good
conductor of electricity
until mixed with boron.
Arsenic has luster and
sheen like metals.

23. Classifying Elements (Section 6.2)
Squares in the Periodic Table
Electron Configuration in Groups
Transition Elements

24. I.) Squares in the Periodic Table
All periodic tables display at least the symbol, the atomic number, and the mass number of the elements.
Some periodic tables provide more information such as physical state, electron configuration, and classification of each element.

25. Noble Gases
Alkali
Metals
Alkaline
Earth
Metals
Halogens

26. II.) Electron Configuration Groups
Elements can be sorted into separate groups based on their electron configuration
Noble Gases
Representative Elements
Transition Metals
Inner Transition Metals

27. The Noble Gases Let’s show this for helium, neon, argon, and
These are the elements located in Group 18 (the farthest column to the right)
These are the “inert” gases because they rarely participate in reactions.
The “s” and “p” orbitals of the highest occupied energy levels are filled for all noble gases.
Let’s show this for helium, neon, argon, and
krypton by writing out the electron configuration
for each.

28. Helium
Neon
Argon
Krypton

29. The Representative Elements
These elements display a wide range of physical and chemical properties.
The atoms of the representative elements have “s” and “p” orbitals of the highest occupied energy levels that are not full.
For any element of this group, the group number (the American and European numbering system) equals the number of electrons in the highest occupied energy level.

30. Let’s Look at the Electron Configuration of Some Representative Elements
Lithium
Sodium
Carbon
Silicon

31. Transition Elements
Two kinds of transition metals: transition and inner transition metals – classification is based on the electron configuration of an element.
Transition Metals: Atoms have highest occupied sublevels that have electrons in the “s” and “d” orbitals.
Inner Transition Metals: Atoms of these metals have the highest occupied “s” orbital and nearby “f” orbitals that contain electrons.

32. Let’s Look at the Electron Configuration of Some Transition Elements
Iron
Silver
Nickel
Chromium

33. Let’s Look at the Electron Configuration of Some Inner Transition Elements
Cerium
Uranium

34. The Divisions Based on Electron Configuration
The electron configuration and
the position of an element in the periodic table gives a particular pattern. This pattern are the blocks we see here.

35. Using the periodic table to write electron configurations.
Based upon the blocks that were described in the previous slide, we can write the electron configuration for any element based on its location in the periodic table.
The steps:
Find the element on the periodic table.
Start counting from hydrogen.
Move towards the right
At the end of each row drop down one row
Begin counting towards the right again
Each row represents an energy level.
Each square represents an electron.

37. Noble Gas Configuration
This is a shorter way to write the shorthand notation for electron configurations.
There are four easy steps:
1.) Locate the element in the periodic table.
2.) Find the noble gas that precedes it.
3.) Place the symbol for this gas in brackets ([ ]).
4.) Write the remaining electron configuration.
Write the noble gas configuration for rubidium (Rb).

38. Let’s try this. Write the extended and Noble gas configuration for the following elements.
1.) chlorine
2.) lead

39. Section 6.3 Periodic Trends
Trends in Atomic Size
Ions
Trends in Ionization Energy
Trends in Ionic Size
Trends in Electronegativity

40. Atomic Size
Atomic radius: One half the distance between the nuclei of two atoms of the same element when the atoms are joined.
Distance between Nuclei
Atomic Radius

41. Ionization Energy Ions: An atom or a group of atoms that has a
The energy required to remove an electron from an atom.
Ions: An atom or a group of atoms that has a
positive or negative charge resulting from a
loss or gain of electrons, respectively.
Cation: An ion with a positive charge
Anion: An ion with a negative charge

42. Electronegativity This property is related to chemical bonding
The ability of an atom of an element to attract electrons when the atom is in a compound.
This property is related to chemical bonding
and is best understood by examining chemical
bonds.

43. I.) Periodic Trend in Atomic Radius
In general,
atomic size
increases
from top to
bottom
within a
group and
from right to left across
a period.

44. These trends depend upon the number of protons and electrons being added as we move through the periodic table.

45. Question: Based on the data for alkali metals and halogens, how does the atomic size change within a group? Why?

46. Within a group:
Increasing # of
occupied orbitals.
Shielding of outer
electrons increases
the atomic radius.
Within a period:
Electrons are being added
to the same energy level.
Shielding is constant
Increasing nuclear charge
pulls all electrons closer

48. II.) Ions An atom or a group of atoms that has a
positive or negative charge resulting from a
loss or gain of electrons, respectively.

49. Cations
These are positively charged ions, resulting from a loss of an electron.
Metals tends to lose electrons from their highest occupied energy levels to become cations.
The charge for a cation is written as a number followed by a plus sign.
Na1+
Ca2+

50. Representing the Formation of Cations
Atoms tend to lose their electrons from the outer most energy levels to become cations.
Na → Na+ + e-
Ca → Ca2+ + 2e-
Why would the atoms lose their electrons to become cations?

51. Anions
These are negatively charged ions resulting from a gain of electrons.
Nonmetals tend to add electrons into their highest occupied energy levels to become anions.
The charge for an anion is written with a number followed by a negative sign.
Cl1-
F1-

52. Representing the Formation of Anions
Atoms tend to accept electrons into their highest occupied energy levels to become anions.
e- + Cl → Cl-
e- + F → F-
Why would these atoms accept electrons to become anions?

53. Ionic Compounds
Fe2(SO4)3
FeS2
Fe2O3

54. III.) Trends in Ionization Energy
The energy required to remove an electron from an atom.
This energy is measured when an element is in the gaseous state.
There are ionization energies for every electron in an atom
The 1st ionization energy is the energy needed to remove the first electron from an atom.
Each successive ionization energy increases dramatically.

55. Periodic Trend in 1st Ionization Energy

56. Group Trend:
As size of the atom increases
the nuclear charge has a
smaller effect on the outer most electrons.
Period Trend:
Nuclear charge increases as we move across a period, thus nuclear attraction increases.

57. III.) Trends in Ionic Size

58. IV.) Electronegativity
The ability of an atom of an element to attract electrons when the atom is in a compound.
Electronegativity values are calculated from ionization energies.

59. In general, electronegativity values increase from bottom to top within a group. For representative elements, the values tend to increase from left to right across a period.

61. The Periodic Table
Chapter 6
The
End