1. Trends found on the Periodic Table

2. Periodic Groups
Elements in the same column have similar chemical and physical properties
These similarities are observed because elements in a column have similar e- configurations (same amount of electrons in outermost shell)

3. Periodic Trends
Periodic Trends –can be seen with our current arrangement of the elements (Moseley)
Trends we’ll be looking at:
Electron affinity
Atomic Radius
Ionization Energy
Electronegativity

4. Atomic Size Size goes UP on going down a group.
Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.
Size goes DOWN on going across a period.

5. Atomic Size
Size decreases across a period owing to increase in the positive charge from the protons. Each added electron feels a greater and greater + charge because the protons are pulling in the same direction, where the electrons are scattered.
Large
Small

6. Which is Bigger?
Na or K ?
Na or Mg ?
Al or I ?

7. Ion Sizes Does the size go
up or down when losing an electron to form a cation?

8. Ion Sizes CATIONS are SMALLER than the atoms from which they come.Li +
, 78 pm
2e and 3 p
Forming a cation.
Li,152 pm
3e and 3p
CATIONS are SMALLER than the atoms from which they come.
The electron/proton attraction has gone UP and so size DECREASES.

9. Ion Sizes
Does the size go up or down when gaining an electron to form an anion?

10. Ion Sizes Forming an anion.-
, 133 pm
10 e and 9 p
F, 71 pm
9e and 9p
Forming an anion.
ANIONS are LARGER than the atoms from which they come.
The electron/proton attraction has gone DOWN and so size INCREASES.
Trends in ion sizes are the same as atom sizes.

11. Trends in Ion Sizes
Figure 8.13

12. Which is Bigger?
Cl or Cl- ?
K+ or K ?
Ca or Ca+2 ?
I- or Br- ?

13. . Trend in Electron Affinity:
The energy release when an electron is added to an atom. Most favorable toward NE corner of PT since these atoms have a great affinity for e-.
Period Trends: The halogens gain e- most easily, while elements of groups 2 & 18 are lest likely to gain e-
Group Trends: more difficult to explain

15. Atomic Radius Atomic Radius – size of an atom
(distance from nucleus to outermost e-)

16. Atomic Radius Trend
Group Trend – As you go down a column, atomic radius increases
As you go down, e- are filled into orbitals that are farther away from the nucleus (attraction not as strong)
Periodic Trend – As you go across a period (L to R), atomic radius decreases
As you go L to R, e- are put into the same orbital, but more p+ and e- total (more attraction = smaller size) 16

17. Ionic Radius
Ionic Radius –
size of an atom when it is an ion

18. Ionic Radius Trend
Metals – lose e-, which means more p+ than e- (more attraction) SO…
Cation Radius < Neutral Atomic Radius
Nonmetals – gain e-, which means more e- than p+ (not as much attraction) SO…
Anion Radius > Neutral Atomic Radius

19. Periodic Table: electron behavior
The periodic table can be classified by the behavior of their electrons

20. Ionic Radius Trend anion radius decreases, too.
Group Trend – As you go down a column, ionic radius increases
Periodic Trend – As you go across a period (L to R), cation radius decreases,
anion radius decreases, too.
As you go L to R, cations have more attraction (smaller size because more p+ than e-). The anions have a larger size than the cations, but also decrease L to R because of less attraction (more e- than p+)

21. Ionic Radius

22. Ionic Radius
How do I remember this????? The more electrons that are lost, the greater the reduction in size. Li+1 Be+2 protons 3 protons 4 electrons 2 electrons 2 Which ion is smaller?

23. Ionization Energy
Ionization Energy – energy needed to remove outermost e-

24. Ionization Energy
Group Trend – As you go down a column, ionization energy decreases
As you go down, atomic size is increasing (less attraction), so easier to remove an e-
Periodic Trend – As you go across a period (L to R), ionization energy increases
As you go L to R, atomic size is decreasing (more attraction), so more difficult to remove an e-
(also, metals want to lose e-, but nonmetals do not)

25. Ionization Energy
IE = energy required to remove an electron from an atom (in the gas phase).
Mg (g) kJ ---> Mg+ (g) + e-
This is called the FIRST ionization energy because we removed only the OUTERMOST electron
Mg+ (g) kJ ---> Mg2+ (g) + e-
This is the SECOND IE.

26. Trends in Ionization Energy
IE increases across a period because the positive charge increases.
Metals lose electrons more easily than nonmetals.
Nonmetals lose electrons with difficulty (they like to GAIN electrons).

27. Trends in Ionization Energy
IE increases UP a group
Because size increases (Shielding Effect)

28. Which has a higher 1st ionization energy?
Mg or Ca ?
Al or S ?
Cs or Ba ?

29. Electronegativity, 
 is a measure of the ability of an atom in a molecule to attract electrons to itself.
Concept proposed by
Linus Pauling

30. Periodic Trends: Electronegativity
In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity increases UP a group of elements.
In a period: More protons, while the energy levels are the same, means atoms can better attract electrons. So, electronegativity increases RIGHT in a period of elements.

31. Electronegativity

32. Which is more electronegative?
F or Cl ?
Na or K ?
Sn or I ?

33. Electronegativity
Electronegativity- tendency of an atom to attract e-

34. Electronegativity Trend
Group Trend – As you go down a column, electronegativity decreases
As you go down, atomic size is increasing, so less attraction to its own e- and other atom’s e-
Periodic Trend – As you go across a period (L to R), electronegativity increases
As you go L to R, atomic size is decreasing, so there is more attraction to its own e- and other atom’s e-

35. Reactivity Reactivity – tendency of an atom to react
Metals – lose e- when they react, so metals’ reactivity is based on lowest Ionization Energy (bottom/left corner) Low I.E = High Reactivity
Nonmetals – gain e- when they react, so nonmetals’ reactivity is based on high electronegativity (upper/right corner)
High electronegativity = High reactivity

36. Metallic Character Properties of a Metal – 1. Easy to shape
Conduct electricity 3. Shiny
Group Trend – As you go down a column, metallic character increases
Periodic Trend – As you go across a period (L to R), metallic character decreases (L to R, you are going from metals to non- metals

37. Summary of Trend Periodic Table and Periodic Trends
1. Electron Configuration
3. Ionization Energy: Largest toward NE of PT
4. Electron Affinity: Most favorable NE of PT
2. Atomic Radius: Largest toward SW corner of PT

38. Electron configuration
THe arrangement of electrons in atoms
There are distinct electron configurations for each element on the periodic table

39. Rules governing electron configuration
Aufbau principle ( means building up in german) States that as protons are individually added to the nucleus to build up the element, electrons are added to the atomic orbitals. ( large elements don’t always follow this rule)
Hund’s rule: orbitals of equal energy are each added to the nucleus to build up the elements

40. Paulie exclusion principle: no 2 electrons in the same atom can have the same set of 4 quantum numbers
Heisenberg uncertainty principle It is not possible to accurately measure both the velocity and position of an electron at the same time

41. Aufbau Principle -- “Bottom Up Rule”41

42. Pauli exclusion principle
An orbital can contain a maximum of 2 electrons,
and they must have the opposite “spin.”
Example:
Determine the electron configuration and orbital notation for the ground state neon atom. 42

43. Basic Principle: electrons occupy lowest energy levels available
Rules for Filling Orbitals
Bottom-up
(Aufbau’s principle)
Fill orbitals singly before doubling up
(Hund’s Rule)
Paired electrons have opposite spin
(Pauli exclusion principle) 43

44. Identify examples of the following principles:
1) Aufbau 2) Hund’s rule 3) Pauli exclusion

45. Atomic Spectra
One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit.

46. Atomic Spectra and Bohr
Bohr said classical view is wrong.
Need a new theory — now called QUANTUM or WAVE MECHANICS.
e- can only exist in certain discrete orbits
e- is restricted to QUANTIZED energy state (quanta = bundles of energy)

47. Quantum or Wave Mechanics
Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms.
He developed the WAVE EQUATION
Solution gives set of math expressions called WAVE FUNCTIONS, 
Each describes an allowed energy state of an e-
E. Schrodinger

48. Heisenberg Uncertainty Principle
Problem of defining nature of electrons in atoms solved by W. Heisenberg.
Cannot simultaneously define the position and momentum (= m•v) of an electron.
We define e- energy exactly but accept limitation that we do not know exact position.
W. Heisenberg

49. Arrangement of Electrons in Atoms
Electrons in atoms are arranged as
LEVELS (n)
SUBLEVELS (l)
ORBITALS (ml)

50. QUANTUM NUMBERS n (principal) ---> energy level
The shape, size, and energy of each orbital is a function of 3 quantum numbers which describe the location of an electron within an atom or ion
n (principal) ---> energy level
l (orbital) ---> shape of orbital
ml (magnetic) ---> designates a particular suborbital
The fourth quantum number is not derived from the wave function
s (spin) ---> spin of the electron (clockwise or counterclockwise: ½ or – ½)

51. QUANTUM NUMBERS
So… if two electrons are in the same place at the same time, they must be repelling, so at least the spin quantum number is different!
The Pauli Exclusion Principle says that no two electrons within an atom (or ion) can have the same four quantum numbers.
If two electrons are in the same energy level, the same sublevel, and the same orbital, they must repel.
Think of the 4 quantum numbers as the address of an electron… Country > State > City > Street

52. Energy Levels
Each energy level has a number called the PRINCIPAL QUANTUM NUMBER, n
Currently n can be 1 thru 7, because there are 7 periods on the periodic table

53. Energy Levels
n = 1
n = 2
n = 3
n = 4

54. Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

55. Types of Orbitals
The most probable area to find these electrons takes on a shape
So far, we have 4 shapes. They are named s, p, d, and f.
No more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwise

56. Types of Orbitals (l)
s orbital
p orbital
d orbital

57. p Orbitals
this is a p sublevel with 3 orbitals
These are called x, y, and z
There is a PLANAR NODE thru the nucleus, which is an area of zero probability of finding an electron
3py orbital

58. p Orbitals
The three p orbitals lie 90o apart in space

59. d Orbitals
d sublevel has 5 orbitals

60. The shapes and labels of the five 3d orbitals.

61. f Orbitals
For l = 3, > f sublevel with 7 orbitals

62. Diagonal Rule
Must be able to write it for the test! This will be question #1 ! Without it, you will not get correct answers !
The diagonal rule is a memory device that helps you remember the order of the filling of the orbitals from lowest energy to highest energy
_____________________ states that electrons fill from the lowest possible energy to the highest energy

63. Diagonal Rule 1 2 3 4 5 6 7 s s 2p s 3p 3d s 4p 4d 4f s 5p 5d 5f 5g?
Steps:
Write the energy levels top to bottom.
Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level.
Draw diagonal lines from the top right to the bottom left.
To get the correct order, follow the arrows! 1 2 3 4 5 6 7 s
s 2p
s 3p 3d
s 4p 4d 4f
By this point, we are past the current periodic table so we can stop.
s 5p 5d 5f 5g?
s 6p 6d 6f 6g? 6h?
s 7p 7d 7f 7g? 7h? 7i?

64. Why are d and f orbitals always in lower energy levels?
d and f orbitals require LARGE amounts of energy
It’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy
This is the reason for the diagonal rule! BE SURE TO FOLLOW THE ARROWS IN ORDER!

65. How many electrons can be in a sublevel?
Remember: A maximum of two electrons can be placed in an orbital.
s orbitals
p orbitals
d orbitals
f orbitals
Number of
orbitals
Number of electrons

66. Electron Configurations
A list of all the electrons in an atom (or ion)
Must go in order (Aufbau principle)
2 electrons per orbital, maximum
We need electron configurations so that we can determine the number of electrons in the outermost energy level. These are called valence electrons.
The number of valence electrons determines how many and what this atom (or ion) can bond to in order to make a molecule
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

67. Electron Configurations
2p4
Number of electrons in the sublevel
Energy Level
Sublevel
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

68. Let’s Try It! H Li N Ne K Zn Pb
Write the electron configuration for the following elements: H Li N Ne K Zn Pb

69. An excited lithium atom emitting a photon of red light to drop to a lower energy state.

70. An excited H atom returns to a lower energy level.

71. Orbitals and the Periodic Table
Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental)
s orbitals
d orbitals
p orbitals
f orbitals

72. Shorthand Notation A way of abbreviating long electron configurations
Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration

73. Shorthand Notation Step 1: It’s the Showcase Showdown!
Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ].
Step 2: Find where to resume by finding the next energy level.
Step 3: Resume the configuration until it’s finished.

74. Shorthand Notation [Ne] 3s2 3p5 Chlorine
Longhand is 1s2 2s2 2p6 3s2 3p5
You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s2 2s2 2p6
The next energy level after Neon is 3
So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17
[Ne] 3s2 3p5

75. Practice Shorthand Notation
Write the shorthand notation for each of the following atoms: Cl K Ca I Bi

76. Valence Electrons Br [Ar] 3d10 4s2 4p5
Electrons are divided between core and valence electrons
B 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1
Br [Ar] 3d10 4s2 4p5
Core = [Ar] 3d10 , valence = 4s2 4p5

77. Rules of the Game
No. of valence electrons of a main group atom = Group number (for A groups)
Atoms like to either empty or fill their outermost level. Since the outer level contains two s electrons and six p electrons (d & f are always in lower levels), the optimum number of electrons is eight. This is called the octet rule.

78. Keep an Eye On Those Ions!
Electrons are lost or gained like they always are with ions… negative ions have gained electrons, positive ions have lost electrons
The electrons that are lost or gained should be added/removed from the highest energy level (not the highest orbital in energy!)

79. Keep an Eye On Those Ions!
Tin
Atom: [Kr] 5s2 4d10 5p2
Sn+4 ion: [Kr] 4d10
Sn+2 ion: [Kr] 5s2 4d10
Note that the electrons came out of the highest energy level, not the highest energy orbital!

80. Keep an Eye On Those Ions!
Bromine
Atom: [Ar] 4s2 3d10 4p5
Br- ion: [Ar] 4s2 3d10 4p6
Note that the electrons went into the highest energy level, not the highest energy orbital!

81. Try Some Ions! Write the longhand notation for these: F- Li+ Mg+2
Write the shorthand notation for these:
Br-
Ba+2
Al+3

82. Exceptions to the Aufbau Principle
(HONORS only)
Exceptions to the Aufbau Principle
Remember d and f orbitals require LARGE amounts of energy
If we can’t fill these sublevels, then the next best thing is to be HALF full (one electron in each orbital in the sublevel)
There are many exceptions, but the most common ones are
d4 and d9
For the purposes of this class, we are going to assume that ALL atoms (or ions) that end in d4 or d9 are exceptions to the rule. This may or may not be true, it just depends on the atom.

83. Exceptions to the Aufbau Principle
(HONORS only)
Exceptions to the Aufbau Principle
d4 is one electron short of being HALF full
In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d5 instead of d4.
For example: Cr would be [Ar] 4s2 3d4, but since this ends exactly with a d4 it is an exception to the rule. Thus, Cr should be [Ar] 4s1 3d5.
Procedure: Find the closest s orbital. Steal one electron from it, and add it to the d.

84. Exceptions to the Aufbau Principle
(HONORS only)
Exceptions to the Aufbau Principle
OK, so this helps the d, but what about the poor s orbital that loses an electron?
Remember, half full is good… and when an s loses 1, it too becomes half full!
So… having the s half full and the d half full is usually lower in energy than having the s full and the d to have one empty orbital.

85. Exceptions to the Aufbau Principle
(HONORS only)
Exceptions to the Aufbau Principle
d9 is one electron short of being full
Just like d4, one of the closest s electrons will go into the d, this time making it d10 instead of d9.
For example: Au would be [Xe] 6s2 4f14 5d9, but since this ends exactly with a d9 it is an exception to the rule. Thus, Au should be [Xe] 6s1 4f14 5d10.
Procedure: Same as before! Find the closest s orbital. Steal one electron from it, and add it to the d.

86. Write the shorthand notation for: Cu W Au
(HONORS only)
Try These!
Write the shorthand notation for: Cu W Au

87. Representing electron configuration
There are 3 different types of notation
Orbital notation
Electron dot notation
Electron configuration notation

88. Ar Kr Xe Ra

89. Orbital notation
An unoccupied orbital is represented by a line________
An orbital containing:
1 electron is represented as an arrow going up
2 electrons is represented as one arrow up and one arrow down ( showing opposite spins of electrons)

90. Stern-Gerlach Experiment
Electron spin
How could an orbital hold two electrons
without electrostatic repulsion?  
Stern-Gerlach Experiment 90

91. Electron dot notation
shows only electrons in the highest or outermost main energy level ( with the highest principle quantum numbers)

93. Electron dot notation with elements leads to the use of lewis structure with compounds

94. Electron configuration notation
eliminates the lines and arrows of orbital notation
Instead the number of electrons in a sublevel is shown

95. 2 ways to write electron configurations
spdf Notation 1 s
value of energy level
sublevel
no. of
electrons
spdf NOTATION
for H, atomic number = 1
Orbital Box Notation
Arrows show electron spin
(+½ or -½)
ORBITAL BOX NOTATION
for He, atomic number = 2 1s 2  95

97. Periodic Table e- configuration from the periodic periodic table (To be covered in future chapters)Li
2s1 Be
2s2 B
2p1 C
2p2 B
2p1 N
2p3 O
2p4 F
2p5 Ne
2p6 Na
3s1 Mg
3s2 Al
3p1 Si
3p2 P
3p3 S
3p4 Cl
3p5 Ar
3p6 K
4s1 Ca
4s2 Sc
3d1 Ti
3d2 V
3d3 Cr
4s13d5 Mn
3d5 Fe
3d6 Co
3d7 Ni
3d8 Cu
4s13d10 Zn
3d10 Ga
4p1 Ge
4p2 As
4p3 Se
4p4 Be
4p5 Kr
4p6 Rb
5s1 Sr
5s2 Y
4d1 Zr
4d2 Nb
4d3 Mo
5s14d5 Tc
4d5 Ru
4d6 Rh
4d7 Ni
4d8 Ag
5s14d10 Cd
4d10 In
5p1 Sn
5p2 Sb
5p3 Te
5p4 I
5p5 Xe
5p6 Cs
6s1 Ba
6s2 La
5d1 Hf
5d2 Ta
5d3 W
6s15d5 Re
5d5 Os
5d6 Ir
5d7 Ni
5d8 Au
6s15d10 Hg
5d10 Tl
6p1 Pb
6p2 Bi
6p3 Po
6p4 At
6p5 Rn
6p6 Fr
7s1 Ra
7s2 Ac
6d1 Rf
6d2 Db
6d3 Sg
7s16d5 Bh
6d5 Hs
6d6 Mt
6d7

99. Shorthand notation practice
[Noble Gas Core] + higher energy electrons
Examples
● Aluminum: 1s22s22p63s23p [Ne]3s23p1
● Calcium: 1s22s22p63s23p64s2
[Ar]4s2
● Nickel: 1s22s22p63s23p64s23d8
[Ar]4s23d8 {or [Ar]3d84s2}
● Iodine: [Kr]5s24d105p5 {or [Kr]4d105s25p5}
● Astatine (At): [Xe]6s24f145d106p5
{or [Xe]4f145d106s26p5}

100. Outer electron configuration for the elements
100

101. Using the periodic table to know configurations1 2 3 4 5 6 7 Ne Ar Kr Xe

102. Valence e’s for “main group” elements

103. Electron configuration for As
103

104. Full Configuration: 1s22s22p63s23p3 Valence Configuration: 3s23p3
Phosphorus
Symbol: P
Atomic Number: 15
Full Configuration: 1s22s22p63s23p3
Valence Configuration: 3s23p3
Shorthand Configuration: [Ne]3s23p3     1s 2s 2p 3s 3p
Box Notation
104

105. electron’s energy depends principally on this
Quantum numbers and orbital energies Each electron in an atom has a unique set of quantum numbers to define it { n, l, ml, ms }
n = principal quantum number
electron’s energy depends principally on this
l = azimuthal quantum number
for orbitals of same n, l distinguishes different shapes (angular momentum)
ml = magnetic quantum number
for orbitals of same n & l, ml distinguishes different orientations in space
ms = spin quantum number
for orbitals of same n, l & ml, ms identifies the two possible spin orientations
Unit 6- Atomic Electon Configurations and Periodicity
105

106. Concept: Each electron in an atom has a unique set of quantum numbers to define it { n, l, ml, ms }
106

107. Electronic configuration of Br
1s2 2s22p6 3s23p63d10 4s24p5
[Ar] 3d104s24p5
[Ar] = “noble gas core”
[Ar]3d10 = “pseudo noble gas core”
(electrons that tend not to react)
Atom’s reactivity is determined by valence electrons
valence e’s in Br: 4s24p5
highest n electrons
107

108. transition metals v. main group elements
Valence e- shells for
transition metals v. main group elements
d orbitals sometimes
included in valence shell
d orbitals not included
in valence shell
(pseudo noble gas cores)

109. Rule-of-thumb for valence electrons
Identify all electrons at the highest principal quantum number (n)
Examples
● Sulfur: 1s22s22p63s23p4 or [Ne]3s23p4
valence electrons: 3s23p4
● Strontium: [Kr]5s2
valence electrons: 5s2
● Gallium: [Ar]4s23d104p1
valence electrons: 4s24p1
● Vanadium: [Ar]4s23d3
valence electrons: 4s2 or 3d34s2
Use on exams,
but recognize
limitations
Use Table 8.9
for online HW

110. Selenium’s valence electrons
Written for increasing energy:
Pseudo noble gas core includes:
 noble gas electron core
 d electrons (not very reactive)
110

111. Core and valence electrons in Germanium
Written for increasing energy:
Pseudo noble gas core includes:
 noble gas core
 d electrons
111