1. Periodic Table It is a systematic catalog of the elements.
Elements are arranged in order of atomic number.

2. Periodicity
When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

3. Periodic Table The rows on the periodic chart are periods.
Columns are groups.
Elements in the same group have similar chemical properties.
Chemists who knew nothing about electrons developed the table- meant to correlate behaviors of elements and to help us remember many facts

4. Groups
These five groups are known by their names.

5. Periodic Table
Nonmetals are on the right side of the periodic table (with the exception of H).
Gases – Group 18 and F, O , N , H, Cl
Liquid – Br
Rest are solids

6. Periodic Table
Metalloids border the stair-step line (with the exception of Al, Po, and At).

7. Periodic Table Metals are on the left side of the chart.
Properties – luster and high electrical and heat conductivity malleable
All except Hg are solids at room temperature

8. Development of Periodic Table
Elements in the same group generally have similar chemical properties.
Physical properties are not necessarily similar, however.

9. Development of Periodic Table
Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

10. Development of Periodic Table
Mendeleev, for instance, predicted the discovery of germanium (which he called eka-silicon meaning under silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.
He left spaces in the table which allowed him to boldly predicted existence and properties

11. Development of Periodic Table
1913 – Henry Mosley developed concept of atomic number
Bombarded different elements with high-energy electrons, found each element produced X-rays of a unique frequency and that the frequency generally increased with the atomic mass
He arranged the X-ray frequencies in order by assigning them a whole number and called it the atomic number
He correctly identifies the atomic number as the number of protons in a nucleus
He left spaces in the table which allowed him to boldly predicted existence and properties

12. Development of Periodic Table
Concept of atomic number clarified some problems in the periodic table
Ex) atomic weight of Ar (18) is greater than that of K (19) yet the chemical properties of Ar are much more like those of Ne and Kr than those of Na and Rb
When elements are arranged by atomic number, rather than increasing atomic weight, Ar and K appear in the correct places
This made is possible to discover previously unknown elements
He left spaces in the table which allowed him to boldly predicted existence and properties

13. Periodic Trends
In this chapter, we will rationalize observed trends in
Sizes of atoms and ions.
Ionization energy.
Electron affinity.

14. Effective Nuclear Charge
Coulombs Law states that the strength of the interaction between two electrical charges depends on the magnitude of the charges and the distances between them .

15. Effective Nuclear Charge
The attractive force between an electron and the nucleus depends on the nuclear charge and on the average distance between the nucleus and the electron
The force increases as the nuclear charge increases and decreases as the electron moves further away

16. Effective Nuclear Charge
In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons.
The nuclear charge that an electron experiences depends on both factors.
In general so many electron-electron repulsions that we cannot analyze the situation exactly

17. Effective Nuclear Charge
Can treat each electron as though it were moving in the net electric field created by the nucleus and the electron density of the other electrons
We view this net electric field as if it results from a single positive charge in the nuclues

18. Effective Nuclear Charge
In many electron atoms the inner electrons partially screen outer electrons from the attraction of the nucleus

19. Effective Nuclear Charge
The effective nuclear charge, Zeff, is found this way:
Zeff = Z − S
where Z is the atomic number and S (positive number) is a screening constant, usually close to the number of inner electrons.
The effective nuclear charge is smaller than the actual nuclear charge because the effective nuclear charge includes the effect of the other electrons
Screening constant represents the portion of the nuclear charge that is screened from the valence electrons by other electrons in the atom
Electrons in the same valence shell do not screen one another very effectively, but they do affect the value of S slightly

20. Effective Nuclear Charge
Na nuclear charge is Z= 11+ there are 10 core electrons 1s22s22p6
Expect S = 10 and the 3s electron to experience an effective nuclear charge of Zeff = = 1+
Situation is more complicated – 3s electron has small probability of being closer to the nucleus in the region occupied by the core electrons – thus there is a possibility that this electron experiences an attraction greater than our simple S = 10 model suggests
This greater attraction turns out to increase the value Zeff for the 3s electron in Na from our expected Zeff = 1+ to Zeff = 2.5+
The fact that the 3s electron spends some small amount of time close to the nucleus changes the value of S from 10 – 8.5

21. Effective Nuclear Charge
For many-electron atom the energies of orbitals with the same n value increases with increasing l value
In C 1s22s22p6, the energy of the 2p orbital (l=1) is higher than that of 2s (l=0) even though both orbitals are in the n=2 shell
Difference in energies is due to the radial probability functions for the orbitals
The greater attraction between 2s electron and the nucleus leads to a lower energy for the 2s orbital than for the 2 p orbital

22. Effective Nuclear Charge
Trends in valence electron Zeff values: The effective nuclear charge increases from left to right across any period of the periodic table
Number of core electrons stays the same across the period, protons increase
Down a column the effective nuclear charge experienced by the valence electrons changes far less than it does across a period

23. Effective Nuclear Charge
n is larger than the value of n for the electron of interest contributes 0
n is equal to the value of n for the electron of interest contribute 0.35
n is 1 less than the n value for the electron of interest contribute .85
n is even smaller contribute 1.00
Ex) F
1s22s22p5

24. Effective Nuclear Charge – Slater
Ex) F
1s22s22p5
Electron of interest n = 2
n = 2 (7 electrons will look at the effect of 6 on 1 )
S = (6*.35) + (2*.85) = 3.8
Zeff = Z- S = 9 – 3.8 = 5.2+

25. What Is the Size of an Atom?
The shortest distance separating the two nuclei during collisions is twice the radii of the atoms – non bonding atomic radius or van der Waals radius

26. What Is the Size of an Atom?
The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. – shorter than the nonbonding atomic radius
The attractive force between two adjacent atoms in the molecule leading to a chemical bond causes this

27. What Is the Size of an Atom?
Example: I2 molecule, the distance separating the nuclei is observed to be 2.66Å, which means the bonding atomic radius of an iodine atom is (2.66Å)/2 = 1.33Å
Helium and neon, the bonding atomic radius must be estimated because there are no known compounds of these elements

28. What Is the Size of an Atom?
Knowing atomic radii allows us to estimate bond lengths
Ex) Cl-Cl bond length in Cl2= 1.99Å so bonding atomic radius is 0.99Ǻ
In CCl4 bond length 1.77Ǻ bonding atomic radius
Close to the sum ( )
1 Ǻ = 10-10m

29. Sizes of Atoms Bonding atomic radius tends to…
…decrease from left to right across a row
(due to increasing Zeff).
…increase from top to bottom of a column
(due to increasing value of n).

30. Sizes of Ions Ionic size depends upon: The nuclear charge.
The number of electrons.
The orbitals in which electrons reside.

31. Sizes of Ions Cations are smaller than their parent atoms.
The outermost electron is removed and repulsions between electrons are reduced.
MELPS Helps!!

32. Sizes of Ions Anions are larger than their parent atoms.
Electrons are added and repulsions between electrons are increased.

33. Sizes of Ions
Ions carrying the same charge ionic radius will increase in size as you go down a column.
This is due to increasing value of n.

34. Sizes of Ions
In an isoelectronic series, ions have the same number of electrons.
Ionic size decreases with an increasing nuclear charge.

35. Ionization Energy
The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.
The first ionization energy is that energy required to remove first electron.
The second ionization energy is that energy required to remove second electron, etc.

36. Ionization Energy
It requires more energy to remove each successive electron.
When all valence electrons have been removed, the ionization energy takes a quantum leap.
Supports the idea that only the outermost electrons are involved in the sharing and transfer of electrons that give rise to chemical bonding and reactions

37. Trends in First Ionization Energies
As one goes down a column, less energy is required to remove the first electron.
For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

38. Trends in First Ionization Energies
Generally, as one goes across a row, it gets harder to remove an electron.
As you go from left to right, Zeff increases.

39. Trends in First Ionization Energies
However, there are two apparent discontinuities in this trend.

40. Trends in First Ionization Energies
The first occurs between Groups IIA and IIIA.
In this case the electron is removed from a p-orbital rather than an s-orbital.
The electron removed is farther from nucleus.
There is also a small amount of repulsion by the s electrons.

41. Trends in First Ionization Energies
The second occurs between Groups VA and VIA.
The electron removed comes from doubly occupied orbital.
Repulsion from the other electron in the orbital aids in its removal.

42. Electron Configurations of Ions
When electrons are removed from an atom to form a cation, they are always removed first from the occupied orbitals having the largest n value
Electrons added to form an anion are added to the empty or partially filled orbital having the lowest value of n

43. Electron Affinity
Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom:
Measures the attraction or affinity of the atom fro the added electron – negative sign would indicate that energy is released during the process
Cl + e−  Cl−

44. Trends in Electron Affinity
In general, electron affinity becomes more exothermic as you go from left to right across a row.

45. Trends in Electron Affinity
There are again, however, two discontinuities in this trend.

46. Trends in Electron Affinity
The first occurs between Groups IA and IIA.
The added electron must go in a p-orbital, not an s-orbital.
The electron is farther from nucleus and feels repulsion from the s-electrons.

47. Trends in Electron Affinity
The second occurs between Groups IVA and VA.
Group VA has no empty orbitals.
The extra electron must go into an already occupied orbital, creating repulsion.

48. Properties of Metal, Nonmetals, and Metalloids
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49. Metals versus Nonmetals
Differences between metals and nonmetals tend to revolve around these properties.

50. Metals versus Nonmetals
Metals tend to form cations.
Nonmetals tend to form anions.
MELPS

51. Metals
They tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

52. Metals Compounds formed between metals and nonmetals tend to be ionic.
Metal oxides tend to be basic.

53. Nonmetals
These are dull, brittle substances that are poor conductors of heat and electricity.
They tend to gain electrons in reactions with metals to acquire a noble gas configuration.

54. Nonmetals
Substances containing only nonmetals are molecular compounds.
Most nonmetal oxides are acidic.

55. Metalloids
These have some characteristics of metals and some of nonmetals.
For instance, silicon looks shiny, but is brittle and fairly poor conductor.

56. Metals Compounds formed between metals and nonmetals tend to be ionic.
Metal oxides tend to be basic.

57. Metallic Character
Metallic character decreases from left to right across a period
Metallic character increases from top to bottom in a group

58. Metal Ions
Metals lose electrons and become positive ions (cations) with a smaller radius than the parent ion
They are easily oxidized
Have low ionization energies

59. Metal Ions Alkali metals have 1+ and alkaline earth metal 2+
Transition metals do not follow a pattern
Compounds between metal and nonmetal are ionic

60. Alkali Metals Alkali metals are soft, metallic solids.
The name comes from the Arabic word for ashes.

61. Alkali Metals
They are found only in compounds in nature, not in their elemental forms.
They have low densities and melting points.
They also have low ionization energies.

62. Alkali Metals
Their reactions with water are famously exothermic.

63. Alkali Metals
Alkali metals (except Li) react with oxygen to form peroxides.
K, Rb, and Cs also form superoxides:
K + O2  KO2
They produce bright colors when placed in a flame.
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64. Alkaline Earth Metals
Alkaline earth metals have higher densities and melting points than alkali metals.
Their ionization energies are low, but not as low as those of alkali metals.

65. Alkaline Earth Metals
Beryllium does not react with water and magnesium reacts only with steam, but the others react readily with water.
Reactivity tends to increase as you go down the group.

66. Alkaline Earth Metals
The heavier alkaline earth ions give off characteristic colors when heated in a hot flame.

67. Nonmetals
Substances containing only nonmetals are molecular compounds.
Most nonmetal oxides are acidic.

68. Nonmetals
These are dull, brittle substances that are poor conductors of heat and electricity.
They tend to gain electrons in reactions with metals to acquire a noble gas configuration.

69. Group 6A Oxygen, sulfur, and selenium are nonmetals.
Tellurium is a metalloid.
The radioactive polonium is a metal.

70. HYDROGEN
Located above be alkali metals because of its electron configuration
Does not truly belong to any group - it is a nonmetal
High ionization energy

71. Oxygen There are two allotropes of oxygen: There can be three anions:
O3, ozone
There can be three anions:
O2−, oxide
O22−, peroxide
O21−, superoxide
It tends to take electrons from other elements (oxidation).

72. Sulfur Sulfur is a weaker oxidizer than oxygen.
The most stable allotrope is S8, a ringed molecule.

73. Group VIIA: Halogens The halogens are prototypical nonmetals.
The name comes from the Greek words halos and gennao: “salt formers”.

74. Group VIIA: Halogens They have large, negative electron affinities.
Therefore, they tend to oxidize other elements easily.
They react directly with metals to form metal halides.
Chlorine is added to water supplies to serve as a disinfectant

75. Group VIIIA: Noble Gases
The noble gases have astronomical ionization energies.
Their electron affinities are positive.
Therefore, they are relatively unreactive.
They are found as monatomic gases.

76. Group VIIIA: Noble Gases
Xe forms three compounds:
XeF2
XeF4 (at right)
XeF6
Kr forms only one stable compound:
KrF2
The unstable HArF was synthesized in 2000.

77. Metals Valence electrons are in a partially-filled band.
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78. Structures of Metallic Solids
Crystal structures of metals are simple enough we can generate the structure by placing a single atom on each lattice point
Primitive cubic structure are rare – radioactive polonium
Body-centered cubic metals – iron, chromium, sodium, and tungsten
Face-centered cubic metals – aluminum, lead, copper, silver, and gold

79. CLOSE PACKING
Shortage of valence electrons and the fact that they are collectively shared make it favorable for the atoms in a metal to pack together closely
Atoms are spherical – understand structure of metals by looking at how spheres pack together
Efficient way – pack one layer of equal sized spheres to surround each sphere by six neighbors
For 3-D structure keeps stacking layer
Second layer will sit in the depressions made by the first spheres

80. CLOSE PACKING Third layer – we have 2 choices
First – third layer in the depressions that lie directly over the spheres in the first layer
Fourth layer would lie directly over the spheres in the second layer (FIG 12-13)
HEXAGONAL CLOSE PACKING (HCP)

81. CLOSE PACKING
Second – third layer does not sit directly above the spheres in either of the first two layers, subsequent layers repeat this sequence giving an ABCABC pattern
CUBIC CLOSE PACKING (CCP)

82. CLOSE PACKING
In both each sphere has 12 equidistant nearest neighbors - 6 in the same layer 3 from the layer above and 3 from the layer below
Each sphere has a coordinate number of 12 – the number of atoms immediately surrounding a given atom in a crystal structure

83. Ceramics These are inorganic solids, usually hard and brittle.
They are highly resistant to heat, corrosion and wear.
Ceramics do not deform under stress.
They are much less dense than metals, and so are used in place of metals in many high-temperature applications.
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84. Alloys
Alloys contain more than one element and have the characteristic properties of metals.
Solid Solution alloys are homogeneous mixtures.
Pure metals and alloys have different physical properties.
An alloy of gold and copper is used in jewelry (the alloy is harder than the relatively soft pure 24 karat gold).
14 karat gold is an alloy containing 58% gold.

85. Metal Alloys-Solid Solutions Substance has mixture of element and metallic properties.
1. Substitutional Alloy: some metal atoms replaced by others of similar size. Electronegativities usually are similar. The atoms must have similar atomic radii.
The elements must have similar bonding characteristics.
brass = Cu/Zn

86. Metal Alloys (continued)
2. Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. Solute atoms occupy interstices “small holes” between solvent atoms. One element (usually a nonmetal) must have a significantly smaller radius than the other (in order to fit into the interstitial site).
steel = iron + carbon
3. Both types: Alloy steels contain a mix of substitutional (Cr, Mo) and interstitial (Carbon) alloys.

87. Alloys vs. Pure Metal
The alloy is much harder, stronger, and less ductile than the pure metal (increased bonding between nonmetal and metal).
An example is steel (contains up to 3% carbon).
mild steels (<0.2% carbon)
useful for chains, nails, etc.
medium steels ( % carbon)
useful for girders, rails, etc.
high-carbon steels ( % carbon)
used in cutlery, tools, springs.
Other elements may also be added to make alloy steels.
Addition of V and Cr increases the strength of the steel and improves its resistance to stress and corrosion.
The most important iron alloy is stainless steel. It contains C, Cr (from ferrochrome, FeCr2), and Ni.
Heterogeneous alloys: The components are not dispersed uniformly (e.g., pearlite steel has two phases: almost pure Fe and cementite, Fe3C).

88. Substitutional
Alloy
Interstitial Alloy

89. Ordered intermetallic structures.
Keywords:
(a) The face-centered cubic unit cell of a cubic close-packed metal. (b) The ordered structure of Ni3Al with nickel atoms shown in gray and aluminum atoms in blue. (c) The ordered structure of TiAl with titanium atoms shown in red and aluminum atoms in blue.

90. Figure 23.17abcd Figure 23-17 Title: Ordered intermetallic structures.
Caption:
(a) The face-centered cubic unit cell of a cubic close-packed metal. (b) The ordered structure of Ni3Al with nickel atoms shown in gray and aluminum atoms in blue. (c) The ordered structure of TiAl with titanium atoms shown in red and aluminum atoms in blue. (d) The ordered structure of ZnCu with copper atoms shown in copper and zinc atoms in gray.
Notes:
Keywords:

91. Which two substances are most likely to form an interstitial alloy?
Nickel and titanium
Silver and tin
Tin and lead
Copper and zinc
Tungsten and carbon
Solute atoms need to have much smaller bonding atomic radius than the solvent atoms typically a nonmetal that makes a covalent bond to the neighboring metal atoms – presence of the extra bonds provided by the interstitial component causes the metal lattice to become stronger and less ductile