1. Periodic properties of the elements
AP Chemistry Chapter 7

2. Development of the periodic table
1869: Dmitri Mendeleev
Arranged elements by increasing atomic mass
Grouped elements with similar properties in columns
Left blank spaces for undiscovered elements
Predicted the properties of Ga & Ge before their discoveries
1913: Henry Moseley
Developed concept of atomic numbers (after Rutherford’s Nuclear Model)
Identified this as the # of protons in the nucleus of the atom
Arranged the periodic table in based on atomic #’s
All elements grouped with similar properties (Ex. Ar & K)
Allowed for “holes” for undiscovered elements

3. Effective nuclear charge
Force of attraction between the nucleus and the electrons
Coulomb’s law: the strength of the interaction between 2 electrical charges (attraction/repulsion) depends on the magnitudes of the charges and the distance between them
Farther away = less force Greater magnitude of charge = more force
Each electron is simultaneously attracted by the nucleus and repelled by other electrons
We can only estimate the net attraction of each electron by the nucleus by considering how it interacts with the average environment created by the nucleus and other electrons in the atom.
We view this net electric field as if it results from a single positive charge located at the nucleus…..We call this the effective nuclear charge (Zeff)

4. Effective nuclear charge
Zeff is smaller than the actual charge (Z) because it takes into account the repulsion between the electrons
A valence electron is attracted to the nucleus and repelled by other electrons in the atom
The inner electrons partially cancel (shield) the attraction between the nucleus and the valence electrons
𝑍 𝑒𝑓𝑓 =𝑍 −𝑆
Screening constant
# of protons
Effective nuclear charge
Screening constant is usually the #
of core electrons in the atom

5. Effective nuclear charge
Example: Na
This is only a rough estimate…the situation is more complicated because of electron distributions of atomic orbitals (uncertainty)

6. Periodic trend for zeff
Zeff tends to increases for valence electrons from left to right across a period on the periodic table
Ex: Li & Be
As you move down a group (column), the Zeff changes far less than it does as you move across a row. (It does increase slightly – but it is less important than the increase in moving across a period)

7. Sizes of atoms & ions
According to the quantum mechanical model, atoms and ions do not have sharply defined boundaries where electron distribution becomes zero….leads to different definitions of atomic size
Nonbonding radius of atoms: The closest distance between colliding atoms without forming bonds
Bonding atomic radius: The attractive interaction between two atoms brings them closer together than in the non-bonding situation
Bonding atomic radius is shorter than the nonbonding radius

9. Sizes of atoms & ions
Space filling models are based on nonbonding radius (van der Waals radius)
Space filling models of molecules uses bonding atomic radii (covalent radii) to determine the amount of overlap of the atoms
Bonding atomic radii can be assigned based on measurements of distances between nuclei in molecules
Ex. I distance between I – I is 2.66 angstroms
assign a bond radius of 1.33 angstroms for I
Cl2
C – C
C – Cl

11. Sample exercise 7.1
Natural gas used in home heating and cooking is odorless. Because natural gas leaks pose the danger of explosion or suffocation, various smelly substances are added to the gas to allow detection of a leak. One such substance is methyl mercaptan, CH3SH, whose structure is shown. Use Figure 7.6 to predict the lengths of the C – S, C – H, and S – H bonds in this molecule.

12. Practice exercise page 267
Using Figure 7.6, predict which will be greater, the P – Br bond length in PBr3 or the As – Cl bond length in AsCl3

13. Periodic trends in atomic radii
Increases within a group (top to bottom)
Decreases left to right across a period
These trends apply to representative elements
The increase down a group has a larger effect than the decrease across a period

14. Sample exercise 7.2
Referring to a periodic table, arrange (as much as possible) the following atoms in order of increasing size: P, S, As, Se

15. Practice exercise page 268
Arrange the following atoms in order of increasing atomic radius: Na, Be, Mg

16. Periodic trends in ionic radii
Radii of ions are based on the distances between nuclei in ionic compounds
Formation of a cation vacates the outer most occupied orbital of the atom and decreases electron – electron repulsion
Results is that cations are smaller than their parent atoms
When electrons are added in the formation of anions
Electron – electron repulsions cause the electrons to spread out more in space
Result is that anions are larger than their parent atoms

17. Periodic trends in atomic radii
For ions with the same charge, size increases as we move down a column in the periodic table

18. Sample Exercise 7.3
Arrange these atoms and ions in order of decreasing size: Mg2+ , Ca2+, and Ca.

19. Practice Exercise (page 268)
Which of the following atoms and ions is largest: S2- , S, O2-

20. Isoelectronic series
A group of ions all containing the same number of electrons is an isoelectronic series
When isoelectronic series is arranged by increasing atomic # (increasing nuclear charge)….radius decreases (larger nuclear charge on the same # of electrons)
Example arrange: Na+, O2-, Al3+, F-, Mg2+

21. Sample Exercise 7.4
Arrange the ions K+, Cl-, Ca2+, and S2- in order of decreasing size.

22. Practice Exercise (page 270)
Which of the following is largest, Rb+, Sr2+ , or Y3+ ?

23. Ionization Energy
Ionization energy is the minimum amount of energy required to remove an electron from an atom or ion
First ionization energy is energy required to remove an electron from ground state of the gaseous atom
Na(g)  Na+ (g) + e- (1st ionization energy)
Na+ (g)  Na2+ (g) + e- (2nd ionization energy)
Greater ionization energy = more difficulty in removing an electron

24. Successive Ionization Energies
Table 7.2
* note sharp increase in ionization energy to remove inner shell electrons

26. Sample Exercise 7.5
Three elements are indicated in the periodic table in the margin. Based on their locations, predict the one with the largest second ionization energy. (Elements are sulfur, sodium, and calcium)

27. Practice Exercise
Which will have the greater third ionization energy, Ca or S?

28. Periodic Trends in First Ionization Energy
Figure 7.10
Within a period, I1 increases left  right (alkali metals lowest  noble gases highest)
Within a column (group) I1 decreases top  bottom
Representative elements show a larger variation in I1 than transition metals do

29. Some exceptions B has lower ionization energy than Be
O has lower ionization energy than N

30. Sample Exercise 7.6
Referring to a periodic table, arrange the following atoms in order of increasing first ionization energy: Ne, Na, P, Ar, K

31. Practice Exercise (page 274)
Which has the lowest first ionization energy, B, Al, C, or Si? Which has the highest first ionization energy?

32. Electron Configurations of Ions
When electrons are removed from atoms to form cations, they are always removed from the orbitals with the largest principal quantum numbers first (4s before 3d)
Li 
Fe2+
Fe3+

33. Sample Exercise 7.7
Write the electron configuration for (a) Ca2+ (b) Co3+ (c) S2-

34. Practice Exercise (page 275)
Write the electron configuration for: (a) Ga3+ (b) Cr3+ (c) Br-

35. Electron Affinities
Electron affinity is the energy change that occurs when an electron is added to a gaseous atom
Ionization energy:
Cl(g)  Cl- (g) + e- ∆E = 1251 kJ/mol
requires input of energy
Cl(g) + e-  Cl-(g) ∆E = -349 kJ/mol
releases energy
More negative electron affinity = a greater attraction between the atom and the added electron

36. Electron Affinity
If ∆E > 0 (like in noble gases)….means that the ion is unstable and will not form

37. Trends in Electron Affinity
Halogens:
Have most negative electron affinities
Gaining one electron gives them a noble gas configuration
Noble Gases:
Have highly positive electron affinity
Group 5A (15):
Have ½ filled p subshells
An additional electron would go into an already occupied orbital resulting in greater electron – electron repulsion
Electron affinity is less negative than elements to the left
Not a large difference in electron affinity as you move down in a group due to the larger orbital (more spread out results in less electron – electron repulsion) and the electron – nucleus attraction also decreases

38. Metals Middle & left of periodic table
Metallic character generally decreases left to right and increases down a group
Conduct heat & electricity
Malleable
Ductile
Shiny luster
Solids at room temp (except Hg)
Low ionization energy
Most metal oxides are basic:
Na2O(s) + H2O(l)  2NaOH(aq)

39. Sample Exercise 7.8
(a) Would you expect aluminum oxide to be a solid, liquid, or gas at room temperature? (b) Write a balanced chemical equation for the reaction of aluminum oxide with nitric acid

40. Practice Exercise (page 279)
Write the balanced chemical equation for the reaction of copper (II) oxide and sulfuric acid.

41. Nonmetals Vary greatly in appearance
Have electron affinities that cause them to gain electrons when bonding with metals
Most nonmetal oxides are acidic
Ex:
CO2(g) + H2O(l)  H2CO3(aq)
P4O10 (s) + 6H2O (l)  4H3PO4 (aq)

42. Sample Exercise 7.9
Write the balanced chemical equations for the reactions of solid selenium dioxide with (a) water (b) aqueous sodium hydroxide

43. Practice Exercise (page 281)
Write the balanced chemical equation for the reaction of solid tetraphosphorus hexoxide with water.

44. Metalloids Have properties between those of metals and nonmetals
Several (Si for example) are semiconductors and are used in computer chips

45. Group Trends: Alkali Metals
Soft, metallic solids
Silvery, metallic luster
High thermal and electrical conductivity
Low density and melting points
Very reactive (low first ionization energy)
Exist in nature only in compounds
Bonds with hydrogen to form hydrides
Can bond with oxygen to form metal peroxides (except lithium)
Some form superoxides (O2-)
Most cases form metal oxides
Each has characteristic flame color

46. Sample Exercise 7.10
Write a balanced equation that predicts the reaction of cesium metal with (a) Cl2(g) (b) H2O (l) (c) H2(g)

47. Practice Exercise (page 283)
Write a balanced equation for the reaction between potassium metal and elemental sulfur.

48. Alkaline Earth Metals Solids at room temperature
Harder, more dense, and higher melting points than alkali metals
Less reactive than alkali metals
Be & Mg are least reactive
Heavier alkaline earth metals give off a characteristic color when heated in a flame

49. Hydrogen Nonmetal Colorless, diatomic gas
Can be metallic under extreme pressure
High ionization energy
Less tendency to lose electrons than alkali metals
Share electrons when bonding with nometals to form molecular compounds
Reactions can be very exothermic
Reacts with active metals to from metal hydrides
Can lose electrons to form cation (H+)

50. Group 6A: Oxygen group Nonmetallic  metallic Oxygen is colorless gas
All other elements are solids
Allotropes (O2) and (O3)
Different forms of the same element in the same state
Sulfur…most common allotrope is S8
Most sulfur in nature exists as metal sufide

51. Group 7A: Halogens “salt formers”
Melting and boiling points increase with atomic number
Fluorine and Chlorine are gases
Bromine is liquid
Iodine is solid
Highly negative electron affinities
Fluorine is most reactive
Chlorine is most important industrial halogen
React with metals to from metal halides…will dissolve in water to form hydrohalic acids

52. Group 8A: Noble Gases Nonmetals Gases at room temperature
Large first ionization energies
Stable electron configurations

53. End of Chapter 7 Assignments (pages 292 – 297)
7.9
7.11
7.13
7.17
7.19
7.21
7.23
7.27
7.29
7.31
7.33
7.37
7.39
7.41
7.43
7.49
7.57
7.59
7.63
7.65
7.69
7.81
7.94
7.105