1. 11/13/2008
Anomalous Electron Configurations
• A few exceptions to the Aufbau
principles exist. Stable configuration:
– half-filled d shell:
• Cr has [Ar]4s13d5;
• Mo has [Kr] 5s14d5
– filled d subshell:
• Cu has [Ar]4s13d10
• Ag has [Kr]5s14d10.
• Au has [Xe]6s14f145d10
• Exceptions occur with larger elements
where orbital energies are similar.
Anomalous electron configuration
of some elements!!!
Element Atomic Expected Actual
number configuration Configuration
Cr 24 3d4 4s2 3d5 4s1
Cu 29 3d9 4s2 3d10 4s1
Mo 42 4d4 5s2 4d5 5s1
*Pd 46 4d8 5s2 4d10 5s0
Ag 47 4d9 4s2 4d10 5s1
*Pt 78 5d8 6s2 5d9 6s1
Au 79 5d9 6s2 5d10 6s1
The explanation for this deviation lies in the superior
stability of completely filled or all half-filled orbitals than
half-
nearly filled of nearly half-filled orbitals.
half- orbitals.
****They are exception to this rule also. So, remember.
Some other of this kinds are Nb(41), Ru(44), W(74),
Nb(41), Ru(44),
Sg(106)****
Sg(106)****
1

2. 11/13/2008
An application of Electronic Configuration of
Atom
Locating elements in The Periodic Table
• In The Periodic Table elements are organized on the basis
of their electronic configuration
• The highest ‘n’ value in electronic configuration
determines the period of that element.
• No of outer shell e determines the Group
§ If at s & p orbital of the outer shell have 1 to 7 e (s1 to s2
p5) then they are elements of Group I to VII
Ex: Na (11) – 1s22s22p63s1 : Group I
Cl (17) – 1s22s22p63s23p5 : Group VII
§ If at s & p orbital of outer shell have s2 p6 then the
element is in group O. Ex: Ne (10)
§ Again at d & s orbital of outer shell determines the Group
Ex: Sc (21) - 1s22s22p63s23p63d14s2 : Group III
Cont’d
§ But at d & s orbital of outer shell, if there are 8/9/10
e then it is in the Group VIII
Ex: Fe (26) - 1s22s22p63s23p63d64s2 : Group VIII
§ If you have more then 10 e at d & s of outer shell
then only ‘s’ orbital’s e will give the Group
Ex: Cu (29) - 1s22s22p63s23p63d104s1 : Group I
• ‘A’ Sub Group: if at the outer shell you do not have
‘d’ orbital or ‘d’ is filled only then it in ‘A’ Sub
Ex: Cl (17) – 1s22s22p63s23p5 : Group VIIA
Ga (31) - 1s22s22p63s23p63d104s24p1 : Group IIIA
• ‘B’ Sub Group: if you have e in ‘d’ orbital at outer
shell (d1–d10) then it is in ‘B’ Sub Group
Ex: Sc (21) - 1s22s22p63s23p63d14s2 : Group IIIB
Zn (30) - 1s22s22p63s23p63d104s2 : Group IIB
2

3. 11/13/2008
Electron Configurations can be
Determined From the Position in
the Periodic Table:
• Elements in group 1(1A) end in ns1
• Elements in group 2 (2A) end in ns2
• Elements in group 13 (3A) end in ns2np1
• Elements in group 14 (4A) end in ns2np2
• Elements in group 15 (5A) end in ns2np3
• Elements in group 16 (6A) end in ns2np4
• Elements in group 17 (7A) end in ns2np5
• Elements in group 18 (8A) end in ns2np6
Valence Electrons
• The electrons in all the sub-shells with the
highest principal energy shell are called the
valence electrons
• e in lower energy shells are called core
electrons
• Starting with one valence electron for the first
element in a period, the number of electrons
increases as you move from left to right across
a period.
• chemists have observed that one of the most
important factors in the way an atom behaves,
both chemically and physically, is the number
of valence electrons
3

4. 11/13/2008
Down the Periodic Table
•Family: Are arranged vertically down the periodic table
(columns or group, 1- 18 or 1-8 A,B)
1- 1-
•These elements have the same number electrons in the outer most
shells, the valence shell.
1 18
IA VIIIA
2
Alkali Family: 13 14 15 16 17
1 IIA 1 e- in the valence shell IIIA IVA VA VIA VIIA
2
3 4 5 6 7 8 9 10 11 12
3 IIIB IVB VB VIB VIIB VIIIB IB IIB
4
Halogen Family:
7 e- in the valence shell
5
6
7
When an atom or molecule gain or loses an e- it becomes
an ion
• A cation has lost e- and therefore has positive charge
• An anion has gained an e- and therefore has a negative
charge.
Electron Configuration of Anions
• Anions are formed when atoms gain enough
electrons to have 8 valence electrons
– filling the s and p sublevels of the valence shell
• The sulfur atom has 6 valence electrons
S atom = 1s22s22p63s23p4
• In order to have 8 valence electrons, it must
gain 2 more
S2- anion = 1s22s22p63s23p6
4

5. 11/13/2008
Electron Configuration of
Cations
• cations are formed when an atom loses all its
valence electrons
– resulting in a new lower energy level valence
shell
– however the process is always endothermic
• the magnesium atom has 2 valence electrons
Mg atom = 1s22s22p63s2
• when it forms a cation, it loses its valence
electrons
Mg2+ cation = 1s22s22p6
Electron Configuration of
Cations
• for transition metals electrons, may be
removed from the sublevel closest to the
valence shell
Al atom = 1s22s22p63s23p1
Al+3 ion = 1s22s22p6
Fe atom = 1s22s22p63s23p64s23d6
Fe+2 ion = 1s22s22p63s23p63d6
Fe+3 ion = 1s22s22p63s23p63d5
Cu atom = 1s22s22p63s23p64s13d10
Cu+1 ion = 1s22s22p63s23p63d10
5

6. 11/13/2008
The Periodic Table of
Elements
Week -7
Lecture 13 &14
Development of the Periodic Table
• There were 114 elements known by 1999.
• The majority of the elements were discovered
between 1735 and 1843.
• How do we organize 114 different elements in
a meaningful way that will allow us to make
predictions about undiscovered elements?
• Arrange elements to reflect the trends in
chemical and physical properties.
• First attempt (Mendeleev and Meyer) arranged
the elements in order of increasing atomic
weight.
6

7. 11/13/2008
Arranging the Elements
The elements were first arranged in this way by Dmitri
Mendeleev, a professor at St. Petersburg University, in 1869.
His arrangement was based on atomic mass.
When Mendeleev was setting out the table,
only 63 elements had been discovered. His
big idea was to leave gaps for yet to be
discovered elements. He was able to predict
the properties of some of these elements,
including silicon and boron. When his
predictions were shown to be accurate his
table became accepted, and it is the basis of
the one we use today.
The Father of the Periodic Table —
Dimitri Mendeleev
• Mendeleev was the first scientist to notice the
relationship between the elements
– Arranged his periodic table by atomic mass
• Moseley later discovered that the periodic nature of
the elements was associated with atomic number, not
atomic mass
• The Periodic Law
– When elements are arranged in order of increasing
atomic number, there is a periodic pattern in their
physical and chemical properties.
7

8. 11/13/2008
Introduction of The Periodic Table
– The periodic table is made up of rows of elements
and columns.
– An element is identified by its chemical symbol.
– The number above the symbol is the atomic
number
– The number below the symbol is the atomic
weight of the element.
– A row is called a period
– A column is called a family or group
– Elements are arranged left to right and top to
bottom in order of increasing atomic number
– This order usually coincides with increasing
atomic mass
8

9. 11/13/2008
Atoms & Elements
Key Concepts:
Elements
are
consist of
Metals or Non metals
Atoms
that have that have
Subatomic particles
Chemical symbols
Neutrons Electrons
arranged in the Protons
are in
Periodic Table
determine Make up the
by Energy levels
Nucleus with
Groups Atomic
Periods Number
Outer shell electrons
has a
determine
Mass Number Periodic
law Group number
• Periodic Patterns
– The chemical behavior of elements is determined by its
electron configuration
– The first three periods contain just A families. Each
period begins with a single electron in a new outer
electron shell.
– Each period ends with a completely filled outer shell
that has the maximum number of electrons for that
shell.
– The outer shell electrons are responsible for chemical
reactions. Elements in the same family have the same
number of outer shell electrons; so they will have
similar chemical properties.
– Group A elements are called representative elements
– Group B elements are called transition elements.
9

10. 11/13/2008
• Chemical Families
– IA are called alkali metals because the react with
water to from an alkaline solution, very soft metals.
(except H2)
– Group IIA are called the alkali earth metals
because mostly we found them in soils as
salts/minerals & they are also reactive, but not as
reactive as Group IA.
• They are also soft metals, though not as soft as alkali
metals
– Group VIIA are the halogens
• These need only one electron to fill their outer shell
• They are very reactive.( disinfectants, bleach)
– Group VIIIA are the noble gases as they have
completely filled outer shells
• They are almost non reactive.
Alkali Metals
10

11. 11/13/2008
Alkaline Earth Metals
Transition Metals
11

12. 11/13/2008
These elements are also
called the rare-earth
elements.
Inner-Transition Metals
Halogens
12

13. 11/13/2008
Noble Gases
Reading the Periodic Table: Classification
• Nonmetals, Metals, Metalloids, Noble gases
13

14. 11/13/2008
Metals, Nonmetals, and Metalloids
Metals
•Metallic character refers to the properties of metals
– Shiny or lustrous
– Malleable (can be hammered into shape)
– Ductile (can be drawn out into wires)
– All except mercury are solids at room temperature
– They are sonorous (make a ringing sound when hit)
– In solution lose electrons in reactions - oxidized
– Most oxides are basic and ionic
Ex: Metal oxide + water → metal hydroxide
Na2O(s) + H2O(l) → 2NaOH(aq)
– Tends to form cation in aqueous solution
• Metallic character increases down a group.
• Metallic character decreases across a period.
Only a few metals are magnetic.
Magnetism is not a property of most metals!
Metals, Nonmetals, and Metalloids
Metals
• When metals are oxidized they tend to form
characteristics cations.
• All group 1A metals form M+ ions.
• All group 2A metals form M2+ ions.
• Most transition metals have variable charges.
14

15. 11/13/2008
Properties of Non-Metals
§ They are poor conductors of electrical
energy
Both a diamond and a pencil ‘lead’ are
§ They are poor conductors of thermal
made of the same element – carbon.
energy
§ Many of them are gases
§ They are brittle if they are solid
§ Form anions
§ Most oxides are acidic
Ex: nonmetal oxide + water → acid
P4O10(s) + H2O(l) → 4H3PO4(aq)
§ Gain electrons in reactions – reduced
§ When nonmetals react with metals,
nonmetals tend to gain electrons:
metal + nonmetal → salt
2Al(s) + 3Br2(l) → 2AlBr3(s)
Metals, Nonmetals, and Metalloids
Metalloids
Metalloids have properties that are
intermediate between metals and nonmetals.
Example: Si (shown here) has a metallic luster
but it is brittle.
Metalloids have found fame in the semiconductor
industry.
15

16. 11/13/2008
Tro, Chemistry: A Molecular
31
Approach
The Groups of the Periodic Table
• Group 1: The Alkali Metals
– Most reactive metals on the PT
– Rarely found free in nature
– Charge of +1 = 1 valence electron
• Group 2: The Alkaline Earth Metals
– Still quite reactive
– Charge of +2 = 2 valence electrons
• Groups 3-12: Transition Metals
– Found freely and in compounds in nature
– Charge is usually +2 but can vary = usually 2
valence electrons
• Group 13: Boron Family
– Charge is +3 = 3 valence electrons
16

17. 11/13/2008
The Groups of the Periodic Table
• Group 14: The Carbon Family
– Contains elements that can form unusual bonds
(carbon and silicon)
– Charge is +4 or -4 = contains 4 valence
electrons
• Group 15: The Nitrogen Family
– Charge is -3 = contains 5 valence electrons
• Group 16: The Oxygen Family
– Also known as the chalcogens
– Charge is -2 = 6 valence electrons
• Group 17: The Halogens
– Most reactive nonmetals
– charge is -1 = 7 valence electrons
• Group 18: The Noble Gases (The Inert Gases)
– Nonreactive
– Charge is 0 = 2 or 8 valence electrons
Periodic Properties
• Periodic law = elements arranged by
atomic number gives physical and
chemical properties varying periodically.
• Various Elemental Properties change fairly
smoothly going across a period or down a
group.
• We will study the following periodic
trends:
– Atomic radii
– Ionization energy
– Electron affinity
– Melting Points and Boiling Points
– Density
17

18. 11/13/2008
Electron Shells and the Sizes of
Atoms
Atomic Sizes
• As a consequence of the ordering in the
periodic table, properties of elements vary
periodically.
• Atomic size varies consistently through the
periodic table.
• As we move down a group, the atoms become
larger.
• As we move across a period, atoms become
smaller.
There are two factors at work:
•principal quantum number, n, and
•the effective nuclear charge, Zeff.
Electron Shells and the Sizes of
Atoms
Atomic Sizes
• As the principle quantum number increases
(i.e., we move down a group), the distance of
the outermost electron from the nucleus
becomes larger. Hence, the atomic radius
increases.
• As we move across the periodic table, the
number of core electrons remains constant.
However, the nuclear charge increases.
Therefore, there is an increased attraction
between the nucleus and the outermost
electrons. This attraction causes the atomic
radius to decrease.
18

19. 11/13/2008
Atomic Radius
Fig. 8.15 Atomic Radii for Main Group Elements
Trends in Atomic Radius
Transition Metals
• increase in size down the Group
• atomic radii of transition metals roughly
the same size across the d block
– valence shell ns2, not the d electrons
– effective nuclear charge on the ns2 electrons
approximately the same
19

20. 11/13/2008
See Figure 8.16
Trends in Ionic Radius
• Ions in same group have same charge
• Ion size increases down the group
– higher valence shell, larger
• Cations smaller than neutral atom; Anions
bigger than neutral atom
• Cations smaller than anions
– except Rb+1 & Cs+1 bigger or same size as F-1
and O-2
• Larger positive charge = smaller cation
– for isoelectronic species
– isoelectronic = same electron configuration
• Larger negative charge = larger anion
– for isoelectronic series
20

21. 11/13/2008
41
Tro, Chemistry: A Molecular
42
Approach
21

22. 11/13/2008
Ionization Energy
• minimum energy needed to remove an
electron from an atom
– gas state
– endothermic process
– valence electron easiest to remove
– M(g) + IE1 → M1+(g) + 1 e-
– M+1(g) + IE2 → M2+(g) + 1 e-
• first ionization energy = energy to remove electron from
neutral atom; 2nd IE = energy to remove from +1 ion; etc.
• IE increases (irregularly) as you move from left
to right across a period.
• IE decreases (irregularly) as you move down a
group.
44
22

23. 11/13/2008
Ionization Energy
• The first ionization energy, I1, is the amount of
energy required to remove an electron from a
gaseous atom:
Na(g) → Na+(g) + e-.
•The second ionization energy, I2, is the energy
required to remove an electron from a gaseous
ion:
Na+(g) → Na2+(g) + e-.
The larger ionization energy, the more difficult it
is to remove the electron.
There is a sharp increase in ionization energy
when a core electron is removed.
Chapter 7 45
General Trends in 1st Ionization
Energy
• larger the effective nuclear charge on the
electron, the more energy it takes to
remove it
• the farther the most probable distance the
electron is from the nucleus, the less
energy it takes to remove it
• 1st IE decreases down the group
– valence electron farther from nucleus
• 1st IE generally increases across the period
– effective nuclear charge increases
46
23

24. 11/13/2008
Tro, Chemistry: A Molecular
47
Approach
Ionization Energy: Periodic table
Fig. 8.18 Ionization Energy vs atomic #
Chapter 8-48
24

25. 11/13/2008
49
Example – Choose the Atom in Each Pair
with the Higher First Ionization Energy
1) Al or S
2) As or Sb
3) N or Si
4) O or Cl? opposing trends
Tro, Chemistry: A Molecular
50
Approach
25

26. 11/13/2008
Irregularities in the Trend
• Ionization Energy generally increases from
left to right across a Period
• except from 2A to 3A, 5A to 6A
↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑
Be N
1s 2s 2p 1s 2s 2p
↑↓ ↑↓ ↑ ↑↓ ↑↓ ↑↓ ↑ ↑
B O
1s 2s 2p 1s 2s 2p
Which is easier to remove an electron
from N or O? Why?
from B or Be? Why?
Tro, Chemistry: A Molecular
51
Approach
Irregularities in the
First Ionization Energy Trends
↑↓ ↑↓ ↑↓ ↑
Be Be+
1s 2s 2p 1s 2s 2p
To ionize Be you must break up a full sublevel, cost extra energy
↑↓ ↑↓ ↑ ↑↓ ↑↓
B B+
1s 2s 2p 1s 2s 2p
When you ionize B you get a full sublevel, costs less energy
B, Al, Ga, etc.: their ionization energies are slightly less than the ionization
energy of the element preceding them in their period.
• Before ionization ns2np1.
• After ionization is ns2. Higher energy ⇒ smaller radius.
Tro, Chemistry: A Molecular
52
Approach
26

27. 11/13/2008
Irregularities in the
First Ionization Energy Trends
↑↓ ↑↓ ↑ ↑ ↑ ↑↓ ↑↓ ↑ ↑
N N+
1s 2s 2p 1s 2s 2p
To ionize N you must break up a half-full sublevel, cost extra energy
↑↓ ↑↓ ↑↓ ↑ ↑ ↑↓ ↑↓ ↑ ↑ ↑
O O+
1s 2s 2p 1s 2s 2p
When you ionize O you get a half-full sublevel, costs less energy
Group 6A elements.
• Before ionization ns2np4.
• After ionization ns2np3 where each p electron in different orbital (Hund’s rule).
Tro, Chemistry: A Molecular
53
Approach
Trends in Successive
Ionization Energies
• removal of each successive
electron costs more energy
– shrinkage in size due to
having more protons than
electrons
– outer electrons closer to the
nucleus, therefore harder to
remove
• regular increase in energy
for each successive valence
electron
• large increase in energy
when start removing core
electrons
Tro, Chemistry: A Molecular
54
Approach
27

28. 11/13/2008
HIGHER IONIZATION ENERGIES
Chapter 8-55
56
28

29. 11/13/2008
Electron Affinities
• Electron affinity is the opposite of ionization
energy.
• Electron affinity is the energy change when a
gaseous atom gains an electron to form a gaseous
ion:
Cl(g) + e- → Cl-(g)
• Electron affinity can either be exothermic (as
the above example) or endothermic:
• more energy released (more -); the larger the EA
generally increases across period
becomes more negative from left to right
not absolute
lowest EA in period = alkali earth metal or noble gas
highest EA in period = halogen
57
The added electron in Cl is placed in the 3p orbital to form the stable 3p6
electron configuration.
29

30. 11/13/2008
Summary of Trend
1. Ionization Energy: Largest toward
3. Electron Affinity: Most favorable
1. Atomic Radius: Largest toward
Magnetic Properties of
Transition Metal Atoms & Ions
• electron configurations that result in unpaired
electrons mean that the atom or ion will have a
net magnetic field – this is called paramagnetism
– will be attracted to a magnetic field
• electron configurations that result in all paired
electrons mean that the atom or ion will have no
magnetic field – this is called diamagnetism
– slightly repelled by a magnetic field
• both Zn atoms and Zn2+ ions are diamagnetic,
showing that the two 4s electrons are lost before
the 3d
– Zn atoms [Ar]4s23d10
– Zn2+ ions [Ar]4s03d10
60
30

31. 11/13/2008
Example 8.6 – Write the Electron Configuration
and Determine whether the Fe atom and Fe3+
ion are Paramagnetic or Diamagnetic
• Fe Z = 26
• previous noble gas = Ar
– 18 electrons
• Fe atom = [Ar]4s23d6
• unpaired electrons 4s 3d
• paramagnetic
• Fe3+ ion = [Ar]4s03d5
• unpaired electrons
• paramagnetic
61
Melting Points and Boiling Points
•Trends in melting Points and boiling points
can be used as a measure of the attractive
forces between atoms or molecules.
•Within the halogens (group 17 or VIIA)
melting points and boiling points increase so
that at room temperature fluorine and
chlorine are gases, bromine is a liquid, and
iodine is a solid as you go down this periodic
group.
•This indicates that the intermolecular forces
become stronger going down a group.
31

32. 11/13/2008
Melting Point Trends in Period
•In the second period, melting points
increase, going from left to right across
the period for the first four elements.
• Melting points then decrease
drastically for nitrogen, oxygen, and
fluorine, which are all diatomic
molecules.
• The lowest melting point is for neon,
which is monatomic.
Melting Points of Elements
32

33. 11/13/2008
Trends in Density
•Densities of elements increase in a
group as atomic number increases.
•In periods, going from left to right,
densities increase, then decrease.
•Elements with the greatest densities
are at the center of period 6.
Densities of Elements
33

34. 11/13/2008
34