1. The Periodic Table
Trends in Properties

2. Arrangement of the elements
The elements were assigned to certain groups or families in the Periodic Table because they had certain properties.
They DO NOT have certain properties because of their location in the Periodic Table.

3. Periodic Properties
Metallic Character – the degree to which an element acts like a metal or non-metal.
Chemical Properties
Metals
Non-metals
Tend to lose electrons in chemical reactions (form positive ions)
Tend to gain electrons in chemical reactions (form negative ions)

4. Formation of Ions - Metals
Metals tend to lose electrons in chemical reactions, because of this the ions (a charged atom) that they form are smaller than the atoms that they are formed from:

5. Formation of Ions – Non-metals
Non-metals tend to gain electrons in chemical reactions, because of this the ions (a charged atom) that they form are larger than the atoms that they are formed from:

6. Physical Properties of Metals & Non-metals
Malleable
Brittle
Good conductors of heat and electricity
Insulators (Non-conductors
Shiny (have luster)
Dull (lack luster)
Ductile (capable of being drawn into wire)

7. Metallic Character in the Periodic Table

8. The Metalloids or Semi-Metals
Elements that have some properties of metals and non-metals. Metalloids are very important today because of their ability to act as semi-conductors (under some conditions they are good conductors and other conditions they are insulators). All of our modern electronic devices utilize semi-conductors.
Chemically, they can either lose or gain electrons.

9. The Location of the Metalloids
Adjacent to the “Mason-Dixon Line” of Chemistry

10. Ionization Energy
Ionization energy is the energy required to remove the outermost electron from an atom.
Ionization energy is a measure of how tightly held the electrons are in an atom. The higher the ionization energy the more difficult it is to remove an electron from an atom.

11. Let’s Do Some Graphing
To help you understand the trends in ionization energy, obtain a piece of graph paper and graph the ionization energies of the first 18 elements (found in Table S) vs. their atomic numbers. Paste this graph in your chemistry notebook.

12. Trends across a Period
In general, ionization energy increases as you go across a Period.

13. Trends within a Group
In general, ionization energy decreases as you go down a Group.

14. Now let’s graph atomic radii
To help you understand the trends in atomic radii, obtain a piece of graph paper and graph the radii of the first 18 elements (found in Table S) vs. their atomic numbers. Paste this graph in your chemistry notebook.

15. Atomic Radii Trends Across a Period
In general, atomic radii decreases as you go across a Period.

16. Atomic Radii Trends Down a Group
In general, atomic radii increases as you go down a Group.

17. Electronegativity
Electronegativity is a measure of an atom’s ability to attract electrons in a bond (to gain electrons).
Non-Metals have high electronegativities while metals have low electronegativities .
Fluorine is the most electronegative atom.

18. Review of Trends Across a Period Down a Group Metallic Character
Decreases
Increases
Ionization Energy
Atomic Radii
Electronegativity

19. But Why?
Trends in the Periodic Table are a result of the degree to which the electrons are attracted to the nucleus by the electric force.
The two factors that effect the amount of electric force are:
Amount of charge
Distance between the charged particles

20. Charge and the electric force
The bigger the charge on the particles, the stronger the attraction.

21. Distance & the electric force
The farther apart charged particles are, the weaker the attraction between them.

22. Explaining the trends in the Periodic Table
These two characteristics of the electric force are all that are needed to explain the three trends we have seen in the Periodic Table.

23. Nuclear charge within a Period
As we go across Period 2 (from Li to Ne) for example, the atomic number (# of protons in the nucleus) increases.
The increasing charge in the nucleus, exerts a stronger pull on the outermost electrons. This explains the trends we saw earlier:
The electrons are pulled closer to the nucleus by the stronger attraction making the radii smaller. (Radii decrease)
Since the attraction is increasing, the electrons are more tightly held and harder to remove (Ionization energy increases and metallic character decreases)
Larger charge in the nucleus has a greater attraction for bonding electrons (electronegativity increases)

24. Down a group
As we go down a group the nuclear charge will also increase, but another factor is apparently more important, the distance of the outermost electrons.
Since the outermost electrons are going into higher energy levels (at greater distances from the nucleus), the atomic radius of the atoms will become larger (radii increase).

25. Down a group (continued)
And since the outer most electrons are further from the nucleus, the attraction for those electrons will decrease making them easier to remove from the atom. (ionization energy decreases and metallic character increases)
And electrons gained in a bond would not be able to get as close to the nucleus decreasing their attraction to the nucleus. (electronegativity decreases)

26. Let’s recap! Across a Period Down a Group Metallic Character
Decreases
Increases
Ionization Energy
Atomic Radii
Electronegativity
Explanation
Increasing nuclear charge increases the attraction felt by the outermost electrons
Electrons going into higher energy levels at greater distances from the nucleus decreases the attraction felt by the outermost electrons