1. Summary: periodic trends
15. Explain Periodic Trends In general, how can the periodic trends exhibited by the elements be explained?

2. The Octet Rule All elements gain or lose electrons so that they end up with the same electron configuration as the nearest noble gas.
The Octet Rule is the driving force for chemical reactions and properties. When we say that an atom “wants to” do something, what we really means is that the atom is doing it so that it will become more stable.

3. The Shielding Effect
Electrons closer to the nucleus repel electrons farther away from the nucleus because they both have negative charge. As a result, outer electrons are less tightly bound to the nucleus than inner electrons.

4. Periodic Trends
Periodic Trends – properties that show patterns when examined across the periodic table.
Atomic Radius – one half the distance between the nuclei of identical atoms that are bonded together.
Ionization Energy – the energy required to remove one electron from a neutral atom of an element.
Ion – an atom or group of atoms that has a positive or negative charge
Ionization – the process of forming an ion. *** Change the electrons, NOT PROTONS!!!! ***
First Ionization Energy –the first electron…
Second Ionization Energy–the second electron…
Electronegativity – a measure of the ability of an atom in a compound to attract electrons. An uneven concentration of charge.

5. Atomic Radii – how big is the atom
Atomic Radii – how big is the atom? decreases across periods because the higher nuclear charge (positive) pulls the electrons closer to the nucleus increases down groups because energy levels are being added outside the nucleus

6. What are the trends among the elements for ionic radius?
During reactions between metals and nonmetals, metal atoms tend to lose electrons and nonmetal atoms tend to gain electrons. This transfer of electrons has a predictable effect on the size of the ions that form.
Cations are always smaller than the atoms from which they form.
Anions are always larger than the atoms from which they form.

7. Ionization Energy – the energy required to remove one electron from a neutral atom of an element. increases across periods because it takes more energy to remove electrons from an atom with a full outermost energy level. The nucleus “pulls” them in tighter. It becomes “clingy” towards the electrons. (octet rule) decreases down groups because it is easier to overcome the nuclear “pull” for the outermost electrons as the number of energy levels increases (shielding effect)

8. Electronegativity – a measure of the ability of an atom in a compound to attract electrons from other atoms increases across periods as a result of the increasing nuclear charge and ability of the nucleus to attract electrons. The nucleus gets “greedy” for electrons as you move across a period. decreases down groups because the nuclear charge is less able to attract electrons as additional energy levels are added

9. What are the trends among the elements for atomic radius?
The atomic radius of an element is one half of the distance between the nuclei of two atoms of the same element when the atoms are joined.
Atomic radii are often measured in picometers (pm).
1. Interpret Diagrams Which element in the figure above has the largest atomic radius?

10. Group trends in atomic radius
As atomic number increases within a group:
Charge on the nucleus increases—draws electrons closer to the nucleus
Number of occupied energy levels increases—shields electrons from the attraction of protons in the nucleus
The shielding effect is greater than the effect of the increase in nuclear charge—atomic radius increases from top to bottom within a group
3. Predict Is a barium atom larger or smaller than a strontium atom?

11. Period trends in atomic radius
Look again at the graph of atomic radius versus atomic number.
4. Identify Periodic Trends How does atomic radius change within a period on the periodic table?
(contd.)

12. Period trends in atomic radius
As atomic number increases across a period:
Each element has one more proton and one more electron than the preceding element
Electrons are added to the same principal energy level—shielding effect is constant for all the elements in a period
The increasing nuclear charge pulls the electrons in the highest occupied energy level closer to the nucleus—atomic radius decreases across a period
5. Predict Is a barium atom larger or smaller than a cesium atom?

13. Summary: trends in atomic radius
6. Compare If a halogen and an alkali metal are in the same period, which element will have the larger atomic radius?

14. What are the trends among the elements for ionization energy?
The energy required to remove an electron from an atom is called ionization energy.
This energy is measured when an element is in its gaseous state.
The energy required to remove the first electron from an atom is called the first ionization energy.
7. Define What is the ionization energy of an element?

15. Group trends in ionization energy
Looking at the data for alkali metals and noble gases, first ionization energy generally decreases from top to bottom within a group.
(contd.)

16. Group trends in ionization energy
As atomic number (and atomic radius) increases within a group:
Nuclear charge has a smaller effect on the electrons in the highest occupied energy level
Less energy is required to remove an electron from this energy level—first ionization energy is lower
8. Apply Concepts Would it be easier to remove an electron from a boron atom or from an aluminum atom?

17. Period trends in ionization energy
Look again at the graph of first ionization energy versus atomic number.
9. Identify Periodic Trends How does first ionization energy change within a period on the periodic table?
(contd.)

18. Period trends in ionization energy
In general, the first ionization energy of representative elements tends to increase from left to right across a period. This trend can be explained by the nuclear charge and the shielding effect.
The nuclear charge increases across the period, but the shielding effect remains constant.
As a result, there is an increase in the attraction of the nucleus for an electron.
Thus, it takes more energy to remove an electron from an atom.
10. Apply Concepts Would it be easier to remove an electron from a boron atom or from a carbon atom?

19. Summary: trends in ionization energy
11. Compare Which element would have a larger first ionization energy: an alkali metal in Period 2 or an alkali metal in Period 4?

20. What are the trends among the elements for ionic radius?
During reactions between metals and nonmetals, metal atoms tend to lose electrons and nonmetal atoms tend to gain electrons. This transfer of electrons has a predictable effect on the size of the ions that form.
Cations are always smaller than the atoms from which they form.
Anions are always larger than the atoms from which they form.
12. Compare Is a sodium atom (Na) larger or smaller than a sodium ion (Na+)?

21. Group trends in ionic radius
For each of the Group 1A metals below, the ion is much smaller than the atom.
When an atom loses an electron, the attraction between the remaining electrons and the nucleus is increased.
As a result, the electrons are drawn closer to the nucleus.
Metals that are representative elements tend to lose all their outermost electrons during ionization—the ion has one fewer occupied energy level.
(contd.)

22. Group trends in ionic radius
For each of the Group 7A nonmetals below, the ion is much larger than the atom.
As the number of electrons increases, the attraction of the nucleus for any one electron decreases.
13. Identify Periodic Trends In general, how does ionic radius vary from top to bottom within a group?

23. Period trends in ionic radius
From left to right across a period, two trends are visible:
A gradual decrease in the size of the positive ions (cations)
A decrease in the size of the negative ions (anions)

24. Summary: trends in ionic radius

25. What are the trends among the elements for electronegativity?
Electronegativity is the ability of an atom of an element to attract electrons when the atom is in a compound.
Scientists use factors such as ionization energy to calculate values for electronegativity.
The electronegativity data on the next slide are expressed in Pauling units.

26. Electronegativity Values for Selected Elements
H 2.1
Li 1.0
Be 1.5
B 2.0
C 2.5
N 3.0
O 3.5
F 4.0
Na 0.9
Mg 1.2
Al 1.5
Si 1.8
P 2.1
S 2.5
Cl 3.0
K 0.8
Ca 1.0
Ga 1.6
Ge 1.8
As 2.0
Se 2.4
Br 2.8
Rb 0.8
Sr 1.0
In 1.7
Sn 1.8
Sb 1.9
Te 2.1
I 2.5
Cs 0.7
Ba 0.9
Tl 1.8
Pb 1.9
Bi 1.9

27. Trends in electronegativity
In general, electronegativity values decrease from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period.
Metals at the far left of the periodic table have low values.
Nonmetals at the far right (excluding noble gases) have high values.
The values among the transition metals are not as regular.
14. Sequence Which element in the periodic table is the most electronegative? Which element is the least electronegative?

28. Summary: periodic trends
15. Explain Periodic Trends In general, how can the periodic trends exhibited by the elements be explained?