1. Ch 5.3 Electron Configuration and Periodic Properties

2. Science Content Standards for California Public Schools
Atomic and Molecular Structure
1. The periodic table displays the elements in increasing atomic number and shows how periodicity of the physical and chemical properties of the elements relates to atomic structure. As a basis for understanding this concept:
b. Students know how to use the periodic table to identify metals, semimetals, non-metals, and halogens.
c. Students know how to use the periodic table to identify alkali metals, alkaline earth metals and transition metals, trends in ionization energy, electronegativity, and the relative sizes of ions and atoms.
f.* Students know how to use the periodic table to identify the lanthanide, actinide, and transactinide elements and know that the transuranium elements were synthesized and identified in laboratory experiments through the use of nuclear accelerators.
g.* Students know how to relate the position of an element in the periodic table to its quantum electron configuration and to its reactivity with other elements in the table.

3. Atomic Size }
Radius
Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.

4. #1. Atomic Size - Period Trends
Going from left to right across a period, the size gets smaller.
Electrons are in the same energy level.
Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

5. #1. Atomic Size - Group trendsH
As we increase the atomic number (or go down a group), each atom has another energy level.
Valence electrons get further from the nucleus.
So the atoms get bigger. Li Na K Rb

6. Trend in Atomic Radius

7. Ions
An ion is an atom (or group of atoms) that has a positive or negative charge
Atoms are neutral because the number of protons equals electrons
Positive and negative ions are formed when electrons are lost or gained between atoms

8. #2: Ionic Group trends Li1+ Na1+ K1+ Rb1+ Cs1+
Ions therefore get bigger as you go down, because of the additional energy level.
Na1+
K1+
Rb1+
Cs1+

9. Ionic Period Trends
Across the period from left to right, they get smaller.
N3-
O2-
F1-
B3+
Li1+
Be2+
C4+

10. Trends in the Periodic Table
Atomic Size vs. Ion Size

11. #3. Trends in Ionization Energy
Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom).
The energy required to remove 1 electron is called the first ionization energy.

12. Ionization Energy
The second ionization energy is the energy required to remove the second electron.
Always greater than first IE.
The third IE is the energy required to remove a third electron.
Greater than 1st or 2nd IE.

13. 3. Trend in Ionization Potential
The energy required to remove the valence electron.

14. Ionization Energy - Group trends
As you go down a group, the first IE decreases because...
The electron is further away from the attraction of the nucleus, and
There is more shielding.

15. Ionization Energy - Period trends
All the atoms in the same period have the same energy level.
So IE generally increases from left to right.

16. He has a greater IE than H.
Both elements have the same shielding since electrons are only in the first level
But He has a greater nuclear charge H
First Ionization energy
Atomic number

17. These outweigh the greater nuclear charge
Li has lower IE than H
more shielding
further away
These outweigh the greater nuclear charge H
First Ionization energy Li
Atomic number

18. greater nuclear chargeHe
Be has higher IE than Li
same shielding
greater nuclear charge
First Ionization energy H Be Li
Atomic number

19. greater nuclear charge He
B has lower IE than Be
same shielding
greater nuclear charge
By removing an electron we make s orbital half-filled
First Ionization energy H Be B Li
Atomic number

20. First Ionization energyHe
First Ionization energy H C Be B Li
Atomic number

21. First Ionization energyHe N
First Ionization energy H C Be B Li
Atomic number

22. He
Oxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbital N
First Ionization energy H C O Be B Li
Atomic number

23. First Ionization energyHe F N
First Ionization energy H C O Be B Li
Atomic number

24. Ne has a lower IE than He Both are full, Ne has more shielding
Greater distance F N
First Ionization energy H C O Be B Li
Atomic number

25. Na has a lower IE than Li Both are s1 Na has more shieldingHe Ne
Na has a lower IE than Li
Both are s1
Na has more shielding
Greater distance F N
First Ionization energy H C O Be B Li Na
Atomic number

26. First Ionization energy
Atomic number

27. Electron Affinity
Is the energy change that occurs when a neutral atom acquires an electron.
Most atoms release energy when they gain an electron. (negative value)
However, some must be forced to gain an electron which requires energy. (positive value)
Halogens gain electrons most readily.

28. Metals tend to LOSE electrons, from their outer energy level
Ions
Metals tend to LOSE electrons, from their outer energy level
Sodium loses one electron.
There are now more protons (11) than electrons (10), and thus a positively charged particle is formed = “cation”
The charge is written as a number followed by a plus sign: Na1+
Now named a “sodium ion”
Lost an electron, so a decrease in size from atom to ion.

29. Nonmetals tend to GAIN one or more electrons
Ions
Nonmetals tend to GAIN one or more electrons
Chlorine will gain one electron
Protons (17) no longer equals the electrons (18), so a charge of -1
Cl1- is re-named a “chloride ion”
Negative ions are called “anions”
Gained an electron so increase in size from atom to ion.

30. Valence Electrons
The electrons available to be lost, gained, or shared in the formation of chemical compounds.
Main group elements: valence electrons are in the outermost s and p orbitals.

31. Electronegativity is the tendency for an atom to attract electrons.
It decreases as it goes down a group because electrons get farther away from the nucleus.

32. Electronegativity Period Trend
Metals (left side)
They let their electrons go easily
Low electronegativity
Nonmetals (right side).
They want more electrons.
Try to take them away from others
High electronegativity.

33. Summary of Trends Ionization Energy and Electronegativity
Atomic and Ionic Radius

34. Additional Assessment Questions
Topic 5
Question 1
For each of the following pairs, predict which atom is larger.
a. Mg, Sr
d. Ge, Br
b. Sr, Sn
e. Cr, W
c. Ge, Sn

35. Answers a. Mg, Sr Sr b. Sr, Sn Sr c. Ge, Sn Sn d. Ge, Br Ge e. Cr, W W
Additional Assessment Questions
Topic 5
Answers
a. Mg, Sr Sr
b. Sr, Sn Sr
c. Ge, Sn Sn
d. Ge, Br Ge
e. Cr, W W

36. Additional Assessment Questions
Topic 5
Question 2
For each of the following pairs, predict which atom or ion is larger.
a. Mg, Mg2+
d. Cl–, I–
b. S, S2–
e. Na+, Al3+
c. Ca2+, Ba2+

37. Answers a. Mg, Mg2+ Mg b. S, S2– S2– c. Ca2+, Ba2+ Ba2+ d. Cl–, I– I–
Additional Assessment Questions
Topic 5
Answers
a. Mg, Mg2+ Mg
b. S, S2–
S2–
c. Ca2+, Ba2+
Ba2+
d. Cl–, I– I–
e. Na+, Al3+
Na+

38. Additional Assessment Questions
Topic 5
Question 3
For each of the following pairs, predict which atom has the higher first ionization energy.
a. Mg, Na
d. Cl, I
b. S, O
e. Na, Al
c. Ca, Ba
f. Se, Br

39. Answers a. Mg, Na Mg b. S, O O c. Ca, Ba Ca d. Cl, I Cl e. Na, Al Al
Additional Assessment Questions
Topic 5
Answers
a. Mg, Na Mg
b. S, O O
c. Ca, Ba Ca
d. Cl, I Cl
e. Na, Al Al
f. Se, Br Br