1. 5.3 NOTES
Periodic Trends

2. 5.3 NOTES
There are many trends in the periodic table. These trends occur in the main group elements (s block and p block)

3. 5.3 NOTES
1. Atomic radii (estimated size of an atom) Found by measuring the distance nucleus to nucleus of two identical atoms. (p. 150) Typically measured in pm. Why does atomic radii increase as you go down the group. (p. 151)

4. 5.3 NOTES
Why does it decrease as you go across the period? (p. 152)

5. 5.3 NOTES
2. Ionization Energy (energy required to remove electrons from an atom) When an atom loses or gains electrons, it is no longer an atom. Now it is an ion and has a charge. Positive ions (cations) have lost electrons. cation comes from the Greek word meaning “down” Negative ions (anions) have gained electrons. anion comes from the Greek word meaning “up”

6. 5.3 NOTES
Why does ionization energy decrease as you go down a group? (p. 153)

7. 5.3 NOTES
Why does ionization energy increase as you go across a period? (p. 154)

8. 5.3 NOTES
The shielding effect is where inner electrons partially block the pull of the nucleus from the outer electrons. This makes outer electrons easier to pull away from the atom. The more electrons, the more the shielding effect.

9. 5.3 NOTES
What is IE1, IE2, or IE3?

10. 5.3 NOTES
IE1 is the energy required to remove the first electron from an atom. IE2 is for the second electron. Etc.

11. 5.3 NOTES
Why is IE2 always bigger than IE1? IE3 › IE2? (p. 155)

12. 5.3 NOTES
As each electron is removed, the pull of the protons in the nucleus compared to the electrons increases, making it harder to remove an additional electron. Noble gases actually have a higher IE than halogens.

13. 5.3 NOTES
3. Electron affinity (energy change when electrons are added to atoms) A loss of energy (-kJ/mol) indicates an electron is easily added . A gain in energy (+kJ/mol) just the opposite.

14. 5.3 NOTES
Which groups of elements would tend to have very negative electron affinities? Why? (p. 157)

15. 5.3 NOTES
The elements that tend to add electrons to fill outer energy levels tend to have high negative electron affinities because they become more stable as an electron is added.

16. 5.3 NOTES
There is not an easy pattern with electron affinity. Why?

17. 5.3 NOTES
Because atoms become stable with partially filled energy levels, it is difficult to form patterns as you go down a group or across a period. (p. 158)

18. 5.3 NOTES
4. Ionic radii (estimated size of an ion) Typically measured in pm.

19. 5.3 NOTES
Why does ionic radii increase as you go down a group? (p. 159)

20. 5.3 NOTES
Why does ionic radii decrease as you go across a period until group 14? Then what happens? Why? (p. 159)

21. 5.3 NOTES
5. Electronegativity (ability of an atom to attract electrons in a chemical bond) Has no unit of measurement, it is based on a relative scale. The higher the number, the higher the pull for electrons.

22. 5.3 NOTES
Why does this decrease or stay the same as you go down a group? (p. 161)

23. 5.3 NOTES
Why does electronegativity increase as you go across a period? (p. 162)

24. 5.3 NOTES
IONIZATION ENERGY
ELECTRONEGATIVITY
ATOMIC RADII

25. CHAPTER 5 TEST
26 multiple choice (2 points each)
24 Periodic Table fill in (2 points each)
1 extra credit (2 points)

26. History of the periodic table Group names and properties –
CHAPTER 5 TEST
History of the periodic table
Group names and properties –
alkali metals
alkaline earth metals
transition metals
halogens
noble gases
lanthanides
actinides
s, p, d, f blocks
electron configurations of atoms and ions

27. CHAPTER 5 TEST
periodic trends – define and use
atomic radius
ionization energy (IE1 IE2 IE3)
electron affinity (define only)
ionic radius
electronegativity