1. Periodic Trends
Ch 6

2. Trends in Atomic Size (sec 6.3, pg 170)
Atomic size increases from top to bottom within a group and decreases from left to right across a period.

4. Question
1. If a halogen and an alkali metal are in the same period, which one will have the larger radius?
2. How does atomic size change within a group (group trends)?
3. Is an atom of barium, smaller or larger than an atom of cesium?

5. Why -1?
As the atomic no increases within a group, no of occupied energy levels increases.
The shielding effect - (no of occupied orbitals shields electrons in the highest occupied energy level from the attraction of protons in the nucleus) also increases.

6. Why -2?
The shielding effect is constant for all the elements in a period.
The increasing nuclear charge pulls the electrons closer to the nucleus and the atomic size decreases.

7. Question What is the difference between an atom and an ion?
What is the difference between a cation and an anion?

8. Ions
In normal condition no. of proton in an atom is equal to the no. of electron. So the net charge is zero.
If an atom loose or gain electron/s, it becomes an ion.
An ion is an atom has either a net positive or net negative charge.

9. Anion & Cation
If the atom has more electrons than protons, it is a negative ion, or ANION (Cl-). If it has more protons than electrons, it is a positive ion or CATION (Na+).

11. Concept Check
1. When a sodium atom loses an electron, does it become a positively charged or negatively charged ion?
2. When a chlorine atom gains an electron, what it becomes?

12. Trends of Metal & Non-metal
Atoms of metallic elements tend to lose one or more electrons from their highest occupied energy levels.
Atoms of nonmetallic elements, tend to form ions by gaining one or more electrons.

13. Concept Check
What happens to the protons and neutrons during movement of electron from one atom to other?

14. Question 1. Write the electron configuration of these elements:
A. The noble gas in period 3
B. the metalloid in period 3
C. the alkali earth metal in period 3
2. Where are the alkali metals, the alkaline earth metals, the halogens, and the noble gases located in the periodic table?

15. Answer 1. a. Ar, 1s22s22p63s23p6 b. Si, 1s22s22p63s23p2
c. Mg, 1s22s22p63s2
2. Gr. 1A, Gr. 2A, Gr. 7A, and Gr. 8A

16. Question
1. How does atomic size change within groups and across periods?
2. How does an ion form?
3. Compare the size of ions (cations and anions) to the size of the atoms from which they are formed?

17. Answer
1. The size increases from top to bottom in a group and decreases from left to right across the period.
2. When an atom lose or gain electron/s
3. The size of cation is smaller and the size of anion is larger than its atom.

18. Periodic Trends in Ionic Radii (p175)
Analysis and Conclusion
1. Describe how the size changes when an atom forms a cation and when an atom forms an anion.
Cations are smaller than their atoms; anions are larger than their atoms.

19. 2. How do the ionic radii vary within a group of metals
2. How do the ionic radii vary within a group of metals? How do they vary within a group of nonmetals?
Ionic radius increase from top to bottom within a group of metals or within a group of nonmetals.

20. 3. Describe the shape of a portion of the graph that corresponds to one period.
There will be two portions of the curve that slope down from left to right.
4. Is the trend across a period similar or different for periods 2, 3, 4, and 5?
The trend is similar for the periods.

21. 5. Propose explanations for the trends you have described for ionic radii within groups and across periods.
The radii increase within a group because the number of occupied energy levels increases.
The radii of cations decrease across a period because the nuclear charge increases, the shielding effect is constant.

22. Electron Configuration
Arsenic – 1s22s22p63s23p64s23d104p3
Zinc –
1s22s22p63s23p64s23d10

23. Electronegativity
Electronegativity is the ability of an atom to attract electrons when the atom is in a compound.
In NaCl, the chlorine atom takes an electron from the sodium atom and converting the atoms into ions (Na+) and (Cl-).

24. More example
In water (H2O), oxygen atom pulls the electron of the hydrogen atom and make water molecule.

25. Periodic and Group Trends
Electronegativity increases as you move from left to right across a period and decreases as you move down a group (excluding noble gas).

26. Ionization Energy
The energy required to remove an electron from an atom is called ionization energy.
The energy required to remove the first electron from an atom is called the first ionization energy.

27. Example Removing one electron from Li is written as Li →Li + + e-
That means when a lithium atom loses an electron, it becomes a lithium ion with a 1+ charge.

28. 2nd Ionization Energy
The second ionization energy is the energy required to remove an electron from an ion with a 1+ charge.
So the ion produce has a 2+ charge.

29. 3rd Ionization Energy
The 3rd ionization energy is the energy required to remove an electron from an ion with a 2+ charge.
What is the charge of this ion?
3+ charge.

30. Li → Li + + e- 1st ionization energy = 520 KJ/mol
Li + →Li 2+ + e- 2nd ionization energy = 7298 KJ/mlol
Li 2+ →Li 3+ + e- 3rd ionization energy = 11,815 KJ/mol

31. What ionization energy trend (both along the group and across the period) do you observe from video?

32. Group and Periodic Trends in Ionization Energy
Ionization energy decreases from top to bottom within a group.
Ionization energies increase from left to right across a period.

33. Ionization energy tells you how strongly an atom holds onto its outermost electron.
Atoms with high ionization energies hold onto their electrons very tightly than the atoms with low ionization energies .

34. Atoms with low ionization energies are more likely to lose one or more of their outermost electrons and gain a positive charge.

35. Concept Check 1 Why does ionization energy decrease along the group?
Less energy is required to remove an electron from highest occupied energy level.

36. Concept Check 2 Why does ionization energy increase across the period?
electrons are tightly bound to the nucleus as the atoms gets smaller. So it takes more energy to remove an electron from an atom.

37. Concept Check 3
Why is the first ionization energy of a non-metal, such as chlorine, much higher than that of an alkali metal, such as potassium?

38. Answer
The nuclear charge increases from left to right across a period and the shielding effect stays the same, so it becomes harder to remove an electron.

39. Pg 174, fig 6.17
1. Which element in period 2 has the lowest first ionization energy? In period 3?
2. What is the group trend for first ionization energy for noble gases and alkali metals?
3. If you drew a graph for second ionization energy, which element would you have to omit? Explain.

40. Question
1. Which element in each pair has a greater first ionization energy?
a) lithium, boron
b) magnesium, strontium
c) cesium, aluminum
2. Explain the difference between the first and second ionization energy of an element.

41. Answer 1. a) boron b) magnesium c) aluminum
2. The first ionization energy is the energy needed to remove a first electron from an atom .
The 2nd ionization energy is the energy needed to remove a second electron.