1. Chapter 7 – Periodic Properties of the Elements
Homework: 9, 10, 12, 13, 14, 15, 17, 19, 21, 25, 27, 28, 30, 33, 35, 37, 39, 41, 43, 47, 51, 54, 55, 57, 59, 63, 65, 67, 69, 71, 75, 78, 105

2. 7.2 – Effective Nuclear Charge
Because electrons are – charged, they are attracted to nuclei (which are + charged)
Many properties of atoms depend not only on their electron configurations, but also on how strongly their outer electrons are attracted to the nucleus.
Coulomb’s law tells us the strength of this interaction depends on the signs and magnitude of their charges, as well as the distance between them.
Thus, the force of the attraction between electron and nucleus depends on the:
Net nuclear charge acting on the electron
Average distance between the nucleus and the electron
This force of attraction increases as the nuclear charge increase and decreases as the electron moves farther from the nucleus.

3. In many-electron atoms, each electron is simultaneously attracted to the nucleus and repelled by other electrons.
Generally, there are so many electron-electron repulsions that we cannot analyze the situation exactly.
We CAN estimate the net attraction of each electron to the nucleus by looking at how the electron interacts with its average environment
Environment created by the nucleus and the other electrons in the atom.

4. Effective Nuclear Charge
By looking at the average situation, we treat each electron individually as if it were moving in a net electric field created by the nucleus and the electron density of other electrons.
We view this net electric field as if it results from a single positive charge located at the nucleus.
This charge is called the effective nuclear charge, Zeff

5. The Zeff for an electron is less than the actual nuclear charge, because it takes the electron-electron repulsions into account.
So if Z = actual nuclear charge
Zeff < Z

6. Finding Zeff
A valence electron is attracted to the nucleus, but repelled by other electrons in the atom.
In particular, the electron density that is due to the inner electrons are very effective at partially canceling the attraction of the valence electron to the nucleus.
We say the inner electrons partially shield or screen the outer electrons from the attraction in the nucleus.
We can write a relationship between the effective nuclear charge and the number of protons in the nucleus (Z)
Zeff = Z – S

7. Where S is a number called the screening constant
S represents how much the inner electrons shield the valence electron from the nucleus.
The value of S is usually close to the number of core electrons in an atom.
Valence electrons don’t shield each other, only core electrons shield.

8. Expected Zeff Values
What would we expect the Zeff value for Sodium to be?
[Ne]3s1
So Z would be 11 (11 protons)
We would expect S to be 10 (10 inner, or core, electrons)
So the expected Zeff value would be +1.
Is this what the Zeff actually is? No
Because the 3s orbital has a very small change of being close to the nucleus (figure 6.18(c)), there is a change of the 3s electron will have a greater attraction than normal.

9. Implications?
This effective nuclear charge explains energy amounts for the different l values within the same energy level (same n)
2s sublevel has greater probability to be closer to nucleus than the 2p.
This leads to slightly greater attractive force between nucleus and 2s
Which leads to lower energy for 2s than for 2p.
This rule applies to all of the energy differences between sublevels.

10. Trends in Zeff for Valence Electrons
The effective nuclear charge as we move across a row (period) increases.
This is because the core electrons remain the same, but the total nuclear charge increases.
In other words, Z increases, but S remains the same.
Going down a column (family), the effective nuclear charge changes much less than moving across.
The expected nuclear charge should stay the same
But we actually experience a slight increase in effective nuclear charge as we move down a family
But this slight increase is MUCH less than the increase in moving across a period.

11. 7.3 – Sizes of Atoms and Ions
An important property of atoms or ions is their size
But don’t think of an atom or ion as a hard, spherical object
Remember, there are no sharply defined boundaries at which electrons are, then aren’t
But, we can still define atom size in several ways, based upon the distance between individual atoms in different situations.

12. Ways to Determine Radius
Imagine a bunch of argon atoms at room temperature
When they collide, the ricochet apart like pool balls.
This happens because the electron clouds of each atom can’t penetrate the other atom
Electrons repel
The closest distance the nuclei get during such a collision determines the apparent radii of the argon atoms.
We call this radius the nonbinding atomic radius

13. Now imagine a chlorine molecule (Cl2)
A chemical bond holds each chlorine atom to the other
This attractive interaction brings the atoms closer than they would be in a nonbinding collision.
This distance is called the bonding atomic radius, and is shorter than the nonbinding atomic radius.

14. Space-Filling Models
Space-filling atomic models, like the one below are based on the non-binding radii (also called the van der Waals radii) and binding radii
Non-binding radii used to determine size of each sphere
Binding atomic radius used to determine how deep each sphere penetrates the other

15. Chemists use different methods to measure the distance separating nuclei in molecules.
In the I2 molecule
The distance separating the nuclei is found to be 2.66Å (1 Å = m)
We find the bonding atomic radius to be ½ the bond distance, so 1.33Å.
Generally, the bonding atomic radius is ½ the nonbinding radius

16. Where Can We Find Nonbinding Radii?
pg. 266
Figure 7.6
Where Can We Find Nonbinding Radii?

17. Periodic Trends in Atomic Radii
Within each column (family) atomic radius tends to increase as you go down the column.
This is primarily because of the increase in the principal quantum number
Electrons have a greater probability of being farther from the nucleus.
Within each row (period), atomic radius tends to decrease from left to right
This is primarily because of the increase in the Zeff, which draws valence electrons closer to the nucleus
Causing atomic radius to decrease

18. Periodic Trends in Ionic Radii
The radii of ions are based upon the distances between ions in ionic compounds.
Like the size of an atom, the size of an ion depends on its nuclear charge, on the # of electrons it has, and on the orbitals in which the valence electrons are found.

19. Cation
The formation of a cation empties the larger occupied orbitals in an atom, and also decrease the number of electron-electron repulsions.
Cations will be smaller than their parent atoms

20. Anions
The formation of an anion begins to fill the larger occupied orbitals in an atom, and also increases the number of electron-electron repulsions.
Anions will be larger than their parent atoms

21. Similar Charges
For ions with the same charge size increases as we go down in a column of the periodic table

23. Isoelectric Ions
An isoelectric series is a group of ions that have the same number of electrons
O2-, F-, Na+, Mg2+, and Al3+ all have 10 electrons
In any isoelectric series, we can list the members in order of increasing atomic number (nuclear charge)
Because the number of electrons remains the same, the radius of the ions decrease with increasing atomic number, as the electrons are more strongly attracted to the nucleus

24. 7.4 – Ionization Energy
The ease with which electrons are removed from an atom or ion has a major effect on chemical behavior.
The ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground state of the gaseous atom or ion.

25. The first ionization energy, or I1, is the energy needed to remove the first electron from a neutral atom
Na(g)  Na+(g) + e-
The second ionization energy, or I2, is the minimum energy needed to remove the second electron from an ion that has already lost one electron
Na+(g)  Na2+(g) + e-
The greater the ionization energy, the more difficult it is to remove an electron.

26. Variations in Successive Ionization Energies
As successive electrons are removed from an atom, the ionization energy increases for that ion.
I1 < I2 < I3 and so on
This is because it is harder to pull an electron away from a more positive and positive ion.

27. There is a sudden increase in ionization energy when an inner shell electron is removed.
Silicon
[Ne]3s23p2
I1 = 786 kJ kJ/mol
I4 = 4360 kJ/mol
I5 = 16,100 kJ/mol
This is because the 2p electrons are closer to the nucleus than the four 3rd shell electrons, and thus the 2p electrons experience a greater Zeff.

28. Every element shows a large increase in ionization energy when electrons are removed from its noble-gas core.
This lends support to the idea that only the outermost electrons (beyond the noble-gas core) are involved in bonding.
The inner electrons are too tightly bound to the nucleus to be lost or even shared with another atom.

29. Periodic Trends in the First Ionization Energies
Within each row (period), I1 will generally increase with increasing atomic number.
So alkali metals show the lowest, and noble gases show the highest.
Within each column (group, the ionization energy generally decreases with increasing atomic energy.
Helium has higher ionization energy than xenon.

30. The representative elements show a larger range of values of I1 than transition-metals do.
Generally, transition metals increase slowly as we proceed from left to right in a period.
The f-block metals only show a very small increase in ionization energies.

32. In general Smaller atoms have higher ionization energies.
Because the same factors that effect atomic size also influence ionization energies.
The energies needed to remove an electron from the outermost shell depends on both the effective nuclear charge and the distance from the nucleus.
As Zeff increases, ionization energies increase.

33. There are occasional irregularities within a given row.
Explained by energy levels
Higher energy level generally = lower ionization energies.

34. Electron Configuration of Ions
When electrons are removed from an atom to form a cation, they always come from the orbitals of the highest energy level first.
Iron example
Fe  Fe2+
[Ar]3d64s2  [Ar]3d6
If we removed a 3rd electron, it would now be removed from the 3rd energy level, now that the 4th is empty.

35. 7.5 – Electron Affinities
The first ionization energy of an atom measures the energy change associated with the removal of an electron from the atom to form a positively charged ion.
Cl(g)  Cl+(g) + e- ΔE = 1251 kJ/mol
The positive ΔE tells us energy must be put into the atom to remove the electron.

36. Most atoms can gain electrons to form negatively charged ions.
This energy change that occurs when an electron is added to a gaseous atom is called the electron affinity
Measures the attraction, or affinity, of the atom for the added electron.
For most atoms, energy is released when an electron is added.

37. Example
Cl  Cl-
Cl(g) + e-  Cl-(g) ΔE = -349 kJ/mol

38. Ionization Energy ≠ Electron Affinity
Ionization energy measures the ease with which an atom LOSES an electron
Electron affinity measures the ease with which an atom GAINS an electron

39. The greater the attraction between a given atom and an added electron, the more negative the atom’s electron affinity.
For some elements (like the noble gases), the electron affinity has a positive value.
Means the anion is higher in energy than the separated atom and electron
Means we must add energy to it to add an electron

41. Trends in Electron Affinity
The halogens will have the highest electron affinity
Gaining an electron gives them a noble-gas configuration
The electron affinities of the noble-gases, Be, and Mg are positive
Any more electrons and we’d have to start filling up in a currently-empty higher-energy subshell
Interesting notes
(N, P, As, Sb)
The added electron gets put into an orbital that is already occupied (p subshell), leading to increased electron-electron repulsion

42. As we move down a group electron affinity doesn’t change greatly.
Although the electrons would be more spaced out (less electron-electron repulsion) the distance to the nucleus is greater (lower electron-nucleus attraction).

43. 7.6 – Metals, Nonmetals and Metalloids
The elements can be broadly grouped into the categories of metals, nonmetals and metalloids.

45. Most elements are metals, found in the left and middle portion of the periodic table.
Nonmetals are found at the top right corner
Metalloids are between the metals and nonmetals

46. Metals Most have a shiny luster Conduct heat and electricity Malleable
All (except mercury) are solid at room temperature
Low ionization energies
Tend to form cations relatively easily
Characterized as an “oxidizer” because they tend to lose electrons during chemical reactions.

47. Metals and Others
Compounds of metals with nonmetals tend to be ionic substances
Oftentimes produce oxides, a group metal ions where a metal bonds with an oxygen
NiO, Na2O
Most metal oxides are basic
When they dissolve in water, they react to form metal hydroxides
The metal hydroxides are basic
The basicity of metal oxides are due to the oxide ion.

48. Nonmetals Nonmetals vary greatly in appearance Not lustrous
Generally poor conductors of heat and electricity
Melting points tend to be lower than those of metals
Seven nonmetals exist as diatomic molecules
H2, N2, O2, F2, Cl2, Br2, I2
Nonmetals can be solids, liquids or gases at room temperature

49. Chemical Reactions
Because of their electron affinities, nonmetals tend to gain electrons when they react with metals.
Compounds made entirely of nonmetals are molecular substances.
Most nonmetal oxides are acidic
When they dissolve in water they will react to form acids

50. Metalloids B, Si, Ge, As, Sb, Te, At are the metalloids
Have properties somewhere between metals and nonmetals
Silicon looks like a metal, but is brittle and is a poor conductor of heat and electricity than most metals.
Compounds of metalloids have characteristics of the compounds of metals or nonmetals, depending on the compound.

51. 7.7 – Group Trends for the Active Metals
We’re going to look at the chemistry of the alkali metals (1A)and the alkaline earth metals (2A)

52. Group 1A: The Alkali Metals
Soft metallic solids
Silvery metallic luster and high thermal and electrical conduciveness
Named from the Arabic word meaning “ashes”
Originally isolated from wood ashes by early chemists
Low densities and melting points
Melting points decrease as we move down the group
Densities increase as we move down the group

53. Since they have the lowest I1 values, they can have electrons very easily removed
Very reactive
So reactive, they exist in nature only as compounds.
Combine directly with most nonmetals

54. React vigorously with water, producing H2 gas and a solution of an alkali metal hydroxide
M stands for any alkali metal
2M(s)  2H2O(l)  2MOH(aq) + H2(g)
This reaction is very exothermic
Often enough energy is released to ignite the H2
This reaction is more violet with lower alkali metals, since they more easily lose their electron, and are more reactive

55. Other reactions
When oxygen reacts with metals, metal oxides (containing O2- ion) are formed
4Li(s) + O2  2Li2O
When dissolved in water, Li2O and other soluble metal oxides react with water to form OH- ions from the reaction of O2- with H2O

56. Other alkali metals all react with oxygen to form metal peroxides, which contain the O22- ion
2Na(s) + O2(g)  Na2O2(s)
K, Rb, Cs also form compounds that contain the O2- ion, called superoxide ion.
K(s) + O2(g)  KO2(s)

57. Generally Generally we only worry about the oxide formation
Don’t worry about the peroxides or superoxides

58. Flame Tests
Although alkali metal ions are colorless, each emits a color when put into a flame.
The ions are reduced to gaseous metal atoms, then the flame exists the valence electrons.
As the valence electrons fall back, they emit light.

59. Reaction Example
Write a balanced equation for the reaction of cesium metal with water.

60. Group 2A: The Alkaline Earth Metals
Like alkali metals, all solid at room temperature and have typical metallic properties
Tend to be harder, more dense and have higher melting points than the alkali metals
I1 is low, but higher than the alkali metals, so less reactive.

61. Reactions Be doesn’t react with water
Mg won’t react with liquid water, but will react with steam
Mg(s) + H2O(g)  MgO(s) + H2(g)
Ca, Sr, Ba and Ra will react with water (but slower than the alkali metals)
Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
In general, tends to lose their two outer s electrons to form 2+ cations.

62. Burning!
Like the alkali metals, the heavier alkaline earth metals give off colors when strongly heated

63. 7.8 – Group Trends for Selected Nonmetals
Hydrogen
Because it has a 1s1 electron, usually positioned with the alkali metals
But H really doesn’t belong in any group
Not an alkali metal, because it’s usually a colorless diatomic gas
Though under extreme pressure, hydrogen can be metallic
Like found at the core of Jupiter or Saturn

64. Reactions between hydrogen and nonmetals tend to be VERY exothermic
Because there is no nuclear shielding of its sole electron, the ionization energy for H is quite high
Actually closer to the ionization energies for nonmetals such as O or Cl
Less likely to lose an electron than the alkali metals
Tends to share its electrons and form molecular compounds
Reactions between hydrogen and nonmetals tend to be VERY exothermic
Reacts with metals to form solid metal hydrides
Contains the hydride ion, H-
This property of hydrogen gaining electrons makes it more like a halogen than an alkali metal
But hydrogen can lose its electron to form a cation
H+(aq) ion is relatively common, but we’ll study it later

65. Group 6A: The Oxygen Group
Going down 6A, we change from nonmetallic to metallic elements
Oxygen, sulfur and selenium are nonmetals
Tellurium is a metalloid
Polonium is a metal
All except oxygen are solids at room temperature

66. Oxygen 2 common forms of oxygen O2 and O3
O2 commonly referred to as just “oxygen”
O3 referred to as “ozone.”
O2 is more stable
These two forms are examples of allotropes
An allotrope are different forms of the same element in the same state

67. Sulfur Most stable of the sulfur allotropes is S8
Even though solid sulfur contains S8, we usually write it as S(s) to simplify stoichiometric coefficients.
Like oxygen, tends to gain electrons to form sulfides, which form the S2- ion
Most sulfur in nature is actually in the form of metallic sulfides.
Though less likely to form sulfides than oxygen is to form oxides.

68. Group 7A: The Halogens
Astatine is usually left out of this group, because it’s unstable and many of its properties not yet well known.
Typical nonmetals
Melting and boiling points increase with increasing atomic number
F and Cl are gasses, Br is a liquid, and I is a solid
All are diatomic

69. Chemical Properties Highly negative electron affinities
Tend to gain electrons from other atoms to form halide ions, X-
X being used to represent any of the halogens
F and Cl more reactive than Br and I