1. C. Goodman, Doral Academy Preparatory High School, 2011-2013
Ch. 5: Periodic Table
C. Goodman, Doral Academy Preparatory High School,

2. Essential Question: Section 5.1
What is the history of the development of the Periodic Table?
What is the periodic law, and how can it be used to predict physical and chemical properties of elements?
What is the overall organization of the modern Periodic Table?
Three types of elements
Named groups
Other families

3. Section 5.1 Vocabulary Mendeleev Moseley Periodic Law Period Group
Main group elements

4. Periodic Table - Definition
Periodic Table -- an arrangement of the elements in order of their atomic numbers so that elements with the same chemical properties are in the same group (family). Examples: halogens, noble gases, alkali metals.

5. Why is it cool?

6. History of the Periodic Table I
Mendeleev: 1869
Atomic mass
Repeating (periodic) patterns of reactivity
In his favor: predicted the discovery of Gallium, which was isolated in his lifetime
Certain characteristic properties of elements can be foretold from their atomic weights
Problem: Iodine and Tellurium

7. History of the Periodic Table II
Moseley: 1914
Atomic number each element has a unique atomic number; resequenced the table by electronic charge (=atomic #) rather than atomic weight.
Periodic Law
In his favor: solved the “Iodine and Tellurium” problem

8. Moseley’s Periodic Law - Definition
Periodic Law The physical and chemical properties of the elements are periodic functions of their atomic numbers.

10. How to use the periodic table…
Atomic number:
# of protons in the nucleus of an atom
Symbol
Basically the abbreviation for the element
Average atomic mass
# of Protons + # of Neutrons (amu)
Weighted average mass of isotopes of the element
Remember nuclear notation for isotopes?
Notice that the atomic mass is a whole number – it’s not an average
Also notice the different locations of the atomic mass and atomic number.

11. Groups (families)
The Columns
Elements in groups have similar chemical properties
Periods
The Rows
Elements properties vary across periods
The length of each period is determined by the number of electrons that can occupy the sublevels being filled in that period

12. Important terms Main Group Elements s-block + p-block elements
Transition metals d-block elements
Lanthanides and actinides f-block elements
Metalloids, metals, non-metals (see below)

13. 2 Main Sections in Periodic Table
Metals
Nonmetals
- Majority of elements
Good Electrical & Heat Conductors
- Room temperature = most solids
Contain properties
Malleability
Ductility
High tensile strength
Poor Electrical Conductors
Poor Heat Conductors
Room temperature = most gases
One is a liquid at r.t. = Bromine
- Solid nonmetals generally brittle

14. Names of groups Group 1a – Alkali metals
Group 2a – Alkali earth metals
Group 7a (17) – halogens
Group 8a (18) – Noble gases

15. Names of families– Transition metals

16. Names of families– Semiconductors

19. Section 2: Electron Configuration
What is the relationship between the location of atoms in a group, their electron configuration, and their chemical and physical properties?
What are the s-, p-, d-, and f-blocks, and how can their electron configurations of their elements be determined?

20. Section 5.2 vocabulary
Ion
Valence
Valence electrons
s-, p-, d- and f-block elements

21. Valence Electrons
Valence = outermost energy level in which contains electrons (in unexcited state).
Valence electrons are the electrons on the outermost energy level of the element.
The number of valence electrons determines the type of chemical reactions available to the element!

22. What does this have to do with groups?
All main group elements in a particular group have the same number of valence electrons.
Prove it?
Heh heh heh – that’s your job!
Write electron configurations, noble gas notation, of the s and p block elements in the first 5 rows. Write valence electrons (s/b only) in contrasting color
Elements hydrogen – xenon, for columns #1a, 2a, 3a, 4a, 5a, 6a, 7a, 8a
Huhhh? See next slide

23. Valence electrons &energy levels
Purpose: to determine the relationship between the group number and number of valence electrons.
Procedures
For each of the elements in the first 5 periods of the following groups… group (1a, 2a, 3a, 4a, 5a, 6a, 7a, and 8a)
Write the name of the element
Write the type of element
Write the electron configuration, using noble gas notation
Conclusion (Answer the following question)s:
1. How does the group # relate to the number of valence electrons?
2. How do you think the chemical reactivity of the elements in a particular group, relates to this number of valence electrons?

24. Valence electrons & Energy Levels
Conclusion (Answer the following questions):
1. How does the group # relate to the number of valence electrons?
2. How do you think the chemical reactivity of the elements in a particular group, relates to this number of valence electrons?

25. Group 14 Carbon [He]2s22p2 4 Silicon [Ne]3s23p2 Germanium
Element
Electron configuration – noble gas notation
Number of valence electrons
Carbon
[He]2s22p2 4
Silicon
[Ne]3s23p2
Germanium
[Ar]4s23d104p2
Tin
[Kr]5s24d105p2

26. Valence Electrons
Main group elements have characteristic numbers of valence electrons.
Group 1 – 1 valence electron
Group 2 – 2 valence electrons
Groups 13-18
# valence electrons = Group # - 10
Example: Group 13 elements have = 3 valence electrons

27. Valence Electrons
Main group elements have characteristic numbers of valence electrons.
S block
Group 1 – 1 valence electron
Group 2 – 2 valence electrons
P block
Groups 13-18
Group # - 10
Example: Group 13 elements have = 3 valence electrons

28. Relationship Between Periodicity and Electron Configurations

29. Sample problem A
Without looking at the periodic table, identify the group, period, and block in which the element that has the electron configuration [Xe]6s2 is located.
Without looking at the periodic table, write the electron configuration for the Group 1 element in the third period. Is this element likely to be more reactive or less reactive than the element described in (a)?

30. Sample problem B
An element has the electron configuration [Kr] 5s24d5. Without looking at the periodic table, identify the period, block, and group in which this element is located. Then, consult the periodic table to identify this element and the others in its group.

31. Sample problem C
Without looking at the periodic table, write the outer electron configuration for the Group 14 element in the second period. Then, use your periodic table to name the element, and identify it as a metal, nonmetal, or metalloid.

32. In book, p. 135
1: Identify period, block, group, element
[Kr]5s2
2. write configuration of…
a. Group 2 elements
b. the group 2 element in the fourth period.
c. the element in the 3rd period, group 15

33. Sample Problem C Solution
p-block (group # >12
14-10 = 4 electrons in s, p
2 e- in s, 2e- in p
The outer electron configuration is 2s22p2.
The element is carbon, C, which is a nonmetal.

34. More practice problems!
Name the block and group in which each of the following elements is located in the periodic table. Use the periodic table to name each element. Identify each element as a metal, nonmetal, or metalloid. Finally, describe whether each element has high reactivity or low reactivity.
[Xe] 6s24f145d8
[Ne]3s23p2
[Ne]3s23p5
[Xe]4f66s1

35. How do I know which groups are more reactive than others?
For main group elements, look at the number of valence (s/p) electrons
8 valence electrons (noble gases) = not reactive
Less reactive 4<3<2<1 More reactive
Less reactive 4<5<6<7 More reactive

36. Sample Problem D Solution
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem D Solution
The 4f sublevel is filled with 14 electrons. The 5d sublevel is partially filled with nine electrons. Therefore, this element is in the d block.
The element is the transition metal platinum, Pt, which is in Group 10 and has a low reactivity.
b. The incompletely filled p sublevel shows that this element is in the p block.
A total of seven electrons are in the ns and np sublevels, so this element is in Group 17, the halogens.
The element is chlorine, Cl, and is highly reactive.

37. Sample Problem D Solution, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem D Solution, continued
c. This element has a noble-gas configuration and thus is in Group 18 in the p block.
The element is argon, Ar, which is an unreactive nonmetal and a noble gas.
d. The incomplete 4f sublevel shows that the element is in the f block and is a lanthanide.
Group numbers are not assigned to the f block.
The element is samarium, Sm. All of the lanthanides are reactive metals.

38. Section 3: Periodic Trends Essential Questions
Compare the periodic trends of atomic radii, ionization energy, electronegativity, and state the reasons for these variations.
What are valence electrons, and how many are present in atoms of each main-group element?
Compare the atomic radii, ionization energies, and electronegativities of the d-block elements with those of the main-group elements.

40. Section 5.3 Vocabulary Atomic radius Ion Ionization energy Cation
Anion
Electron affinity
Electronegativity

41. Atomic Radii: ½ the distance between the nuclei of identical atoms bonded

42. Periodic Trends: Atomic Radii
DECREASES across periods
Because of increasing positive charge of the nucleus
Holds electrons more tightly
INCREASES down groups
Higher principle quantum number
Valence electrons in higher main energy levels
Located farther from the nucleus

43. Watch it kid, I’ve got my ion you.
Ions

44. Ion – Definition
An Ion is… An atom or group of bonded atoms, which has a positive (+) or negative (-) charge.

45. There are two kinds of ions…
Cation
CRUNCH (subtract) an electron
This results in a positive charge
When an electron is removed, the atom loses bulk (like a muscle which shrinks when it atrophies)
So, the radius of a cation is smaller than the atomic radius

46. Anions ADD an electron This results in a NEGATIVE charge
When an electron is added, the atom gains bulk (like a muscle which grows when you work out)
So, the radius of a anion is larger than the atomic radius

47. Two sodium atoms bumped into each other.
One said: "Why do you look so sad?“
The other responded: "I lost an electron.“
The first one asked "Are you sure?“
The other replied "I'm positive."

48. Which elements form which type of ion?
Metal on left tend to form cations (+)
Nonmetals at the upper right tend to form anions (-)
Hydrogen is a non-metal, but it forms a cation (+)

49. Periodic Trends: Ionic Radii – very similar to trends for atomic radii

50. Ionization Energy (IE) Definition
Ionization energy is… The energy required to remove one electron from a neutral atom of an element, forming a cation. Note: does not apply to formation of anions! Unit of measure: kJ/mol

51. It takes energy to “steal” electrons
Ionization energy is (almost always) positive (J)

52. Low IE - lose electrons easily
Look p. 145, Figure 3.4
Low IE - lose electrons easily
High IE – it’s harder to lose electrons

53. Ionization Energy:
Atom (neutral) + energy  Cation (A+)+ (e-) (removed)

54. Ionization Energy (IE)
IE increases across periods because of increasing nuclear charge. Why?
Higher charges attracts e- in same energy level more strongly
Trend down groups decrease because of more electrons between the nucleus and furthest out electrons (lower nuclear charge)

55. Watch out!!!
Ionization energy only applies to the formation of cations.
If you want to talk anion formation, you need electron affinity.

56. Electron Affinity
The energy change that occurs when an electron is acquired by a neutral atom
Most release energy when they acquire an electron: Take a neutral atom (A) add an electron (e-)  get an anion (A-) + energy is released
Some must be “forced” to gain an electron by adding energy: A + e- + energy  A-
Common unit used is kJ/mol
Some positive affinities are difficult to determine with any accuracy

57. Here is a mnemonic for electron affinity:
If the ion is negatively charged (anion), the electron affinity is more strongly negative.

58. Why do some elements give up their electrons more easily?
Q: Which group of elements do you think hold on to their electrons the most strongly?
A: The noble gases! Their valence is full, so it’s very hard to remove their electrons!

59. Electron Affinity
Period Trends: Halogens (VII A) gain electrons the most readily
Group Trends: Electrons add with greater difficulty down a group

60. Electron Affinity
Question Based on this trend what elements will most likely form cations?
Which will most likely form anions?
Which will most likely have small ionic radii?

61. Adding Electrons to Negative Ions
Difficult to add a second electron to an already negatively charged ion
Second electron affinities are therefore all positive
Halogens become negative ions by adding one electron (i.e. Cl-)

62. Electronegativity
A measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound
Most electronegative element is Fluorine

63. Electronegativity Trend across periods increase (some exceptions)
More or less, trend down groups decrease
Similar to IE

64. Additional Help Website