1. Tour of the Periodic Table (Chapter 6)
Dr. Walker

2. Objectives Describe the historical development of the periodic table
Name the groups of the periodic table and describe their properties
List elements that are diatomic
Differentiate elements based on periodic trends

3. History First periodic table created by Dmitri Mendeleev
Arranged elements by increasing atomic mass, which caused mistakes
Was incomplete, but pattern predicted existence of elements found later
Current periodic table created by Henry Moseley
Arranged by increasing atomic number

4. Mendeleev’s Periodic Table

5. The columns are called groups or families
The columns are called groups or families. Groups have similar physical and chemical properties.

6. Periodic Law
When elements are arranged in order of increasing atomic numbers, their physical and chemical properties show a periodic pattern.
Elements in the same groups have the same general physical and chemical properties because of their similar number of valence electrons.
Periodicity is regularly repeating patterns or trends in the chemical and physical properties of the elements arranged in the periodic table.

7. More About Periodic Law
Valence Electrons
Electrons in the outer shell of an atom
Determines chemical properties

8. Determining Valence Electrons
How do we determine valence electrons?
Valence electrons = last number of group number
Examples
Sodium – Group 1 = 1 valence electron
Boron – Group 13 = 3 valence electrons
Chlorine – Group 17 = 7 valence electrons
Nitrogen – Group 15 = 5 valence electrons

9. Determining Valence Electrons
We can also get valence electrons from the electron configuration
Superscripts for highest energy level are valence electrons
Includes s AND p electrons
Examples
Sodium – 1s22s22p63s1 = 1 valence electron
There is one electron in the third energy level
Boron – 1s22s22p1 = 3 valence electrons
There are three electrons in the second energy level
Chlorine – 1s22s22p63s23p5 = 7 valence electrons
There are seven electrons in the third energy level

10. Determining Groups Each group has a distinctive electron configuration
ns2np2 = 4 valence electrons = Group 14
1s22s22p2 = carbon
1s22s22p6 3s23p2 = silicon
1s22s22p63s23p6 4s23d104p2 = germanium

11. Determining Valence Electrons
What about the transition metals (groups 3-12), actinides, and lathanides?
We don’t worry about these elements
These elements can use electrons from d and f orbitals, which we won’t deal with

13. Groups of the Periodic Table
Group 1-Alkali Metals
Physical Properties
Soft
Low melting points
Low densities
Chemical Properties
Explodes in water
Tarnish rapidly in air

14. Groups of the Periodic Table
Group 2 - Alkaline Earth Metals
Physical Properties
Soft
Chemical Properties
React well with hot water
Strong reducing agents

15. Groups of the Periodic Table
Groups Transition Metals
Physical Properties
High density
High melting point
Magnetic

16. Groups of the Periodic Table
Metalloids – Diagonal Elements beginning with Boron
B, Si, Ge, As, Sb, and Te
Physical Properties
Semi-conductors
Chemical Properties
Act like metals when they react with non-metals
Act like non-metals when they react with metals

17. Groups of the Periodic Table
Group 17 – Halogens
Physical Properties
increase in density as you go down the column
colored (yellow-green to brown to black)
Chemical Properties
Form salts with elements from alkaline metals
Exist as diatomic molecules
Form acids with hydrogen

18. Groups of the Periodic Table
Group 18 - Noble Gases
Physical Properties
Colorless
Odorless
Tasteless
Chemical Properties
chemically inert (do not react) – full valence shell

19. Hydrogen Why is this separate?
Since it only has one electron, it is placed with group 1
Doesn’t possess any chemical similarities with alkali metals
Physical Properties
Colorless, Tasteless, Odorless
Chemical Properties
Inflammable
Highly reactive
Exists as diatomic molecule

20. Diatomic Molecules Some elements exist as diatomic molecules
Exists as TWO atoms covalently bonded to each other
Mostly top right of periodic table
H2, N2, F2, O2, I2, Cl2, Br2
ClIF H BrON

21. Inner Transition Metals
Not in guided notes!
Inner Transition Metals
Otherwise known as lanthanide and actinide series or “f-block” elements
Some not found in nature (primarily actinides)

22. The rows are called periods
The rows are called periods. The period number matches the principle energy level of the element.

23. Periodic Trends
Atomic Radius: Measure of the distance between radii of two identical atoms of an element.
Ionization Energy: The energy required to remove an electron from an atom in its gaseous state
Electronegativity: Measure of the attraction of an atom for electrons in a bond.

24. Atomic Radius Across periods: Radius size decreases due to increased
nuclear charge
Down a group: Radius size increases to due higher number
of occupied shells

25. Ionization Energy
The energy required to remove an electron from an atom in its gaseous state

26. Electronegativity
Electronegativity: The tendency of an atom to attract electrons to itself when chemically combined with another element.
The halogen group has the highest electronegativity of the families. The first period has the highest electronegativity.

27. Question
If electronegativity increases to the right, why don’t the noble gases have the highest electronegativity?

28. Question
If electronegativity increases to the right, why don’t the noble gases have the highest electronegativity?
Electronegativity involves atoms in a bond. Noble gases DO NOT BOND! Therefore, no electronegativity value

29. Nuclear Shielding
Outer shell electrons feel less effect of the positive nucleus because of the inner shell electrons.
The nuclear shielding effect is constant within a given period and increases within given groups from top to bottom.

30. Shielding Trends
Shielding is equal among elements in the same period

31. Chemical Reactivity
Reactivity refers to how likely or vigorously an atom is to react with other substances.
Metals and nonmetals have their own trends (can’t be easy….can it?)
Not in textbook, but on SOL (figures!)
From

32. Chemical Reactivity - Metals
Increases as you go down
As you go down valence electrons are further from the nucleus
Further from the nucleus, easier to get rid of
Decreases left to right
Further to the right, more electrons to get rid of
Takes more energy to do this
Not in textbook, but on SOL (figures!)

33. Chemical Reactivity - Nonmetals
Decreases as you go down
Higher elements have more electronegativity
The more they “hog” electrons, the faster they react
Increases left to right
The closer you are to filling an octet the more reactive it is
Noble gases don’t count – their octet is full, so they don’t react

34. Chemical Reactivity

35. Review What group is chemically unreactive? What is the Periodic Law?
Give three diatomic elements.
What group consists of elements that explode in water?

36. Review What group is chemically unreactive? Noble gases
What is the Periodic Law?
When elements are arranged in order of increasing atomic numbers, their physical and chemical properties show a periodic pattern.
Give three diatomic elements.
H2, N2, F2, O2, I2, Cl2, Br2
What group consists of elements that explode in water?
Alkali Metals

37. Review Of F, B, and O, which is the most electronegative element?
Of F, B, and O, which element has the largest radius?
Of Rb, Na, and K, which element has the largest shielding effect?

38. Review Of F, B, and O, which is the most electronegative element?
Fluorine (farthest right)
Of F, B, and O, which element has the largest radius?
Boron (farthest left)
Of Rb, Na, and K, which element has the largest shielding effect?
Rubidium (lowest on periodic table)

39. Terms to Know Mendeleev (name) Inner Transition Metals
Moseley (name) Diatomic elements
Valence electrons Atomic Radius
Alkali metals Electronegativity
Alkaline earth metals Ionization Energy
Transition metals Nuclear Shielding
Metalloids
Halogens
Noble gases

40. Skills To Master Determining diatomic elements
Determining valence electrons from their position on the periodic table, by group name, and by electron configuration
Classifying elements by their group name
Arranging elements according to their periodic trends
Electronegativity
Atomic Radius
Shielding
Reactivity