1. Topic: Periodic Trends
Students will identify the different trends in the periodic table in terms of atomic radius, electronegativity, electron affinity
Warm-up: Post lab discussion
Classwork: Periodic Trends Worksheet
Test on Chapter 6

2. Elemental Properties and Patterns
Periodic Trends
Elemental Properties and Patterns

3. The Periodic Law
Dimitri Mendeleev was the first scientist to publish an organized periodic table of the known elements.
He was perpetually in trouble with the Russian government and the Russian Orthodox Church, but he was brilliant never-the-less.

4. The Periodic Law
Mendeleev even went out on a limb and predicted the properties of 2 at the time undiscovered elements.
He was very accurate in his predictions, which led the world to accept his ideas about periodicity and a logical periodic table.

5. The Periodic Law Mendeleev understood the ‘Periodic Law’ which states:
When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.

6. The Periodic Law
Atoms with similar properties appear in groups or families (vertical columns) on the periodic table.
They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior.

7. Valence Electrons
Do you remember how to tell the number of valence electrons for elements in the s- and p-blocks?
How many valence electrons will the atoms in the d-block (transition metals) and the f-block (inner transition metals) have?
Most have 2 valence e-, some only have 1.

8. A Different Type of Grouping
Besides the 4 blocks of the table, there is another way of classifying element:
Metals
Nonmetals
Metalloids or Semi-metals.
The following slide shows where each group is found.

9. Metals, Nonmetals, Metalloids

10. Metals, Nonmetals, Metalloids
There is a zig-zag or staircase line that divides the table.
Metals are on the left of the line, in blue.
Nonmetals are on the right of the line, in orange.

11. Metals, Nonmetals, Metalloids
Elements that border the stair case, shown in purple are the metalloids or semi-metals.
There is one important exception.
Aluminum is more metallic than not.

12. Metals, Nonmetals, Metalloids
How can you identify a metal?
What are its properties?
What about the less common nonmetals?
What are their properties?
And what the heck is a metalloid?

13. Metals
Metals are lustrous (shiny), malleable, ductile, and are good conductors of heat and electricity.
They are mostly solids at room temp.
What is one exception?

14. Nonmetals Nonmetals are the opposite.
They are dull, brittle, nonconductors (insulators).
Some are solid, but many are gases, and Bromine is a liquid.

15. Metalloids Metalloids, aka semi-metals are just that.
They have characteristics of both metals and nonmetals.
They are shiny but brittle.
And they are semiconductors.
What is our most important semiconductor?

16. Periodic Trends
There are several important atomic characteristics that show predictable trends that you should know.
The first and most important is atomic radius.
Radius is the distance from the center of the nucleus to the “edge” of the electron cloud.

17. Atomic Radius
Since a cloud’s edge is difficult to define, scientists use define covalent radius, or half the distance between the nuclei of 2 bonded atoms.
Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is 1 x m.

18. Covalent Radius
Two Br atoms bonded together are 2.86 angstroms apart. So, the radius of each atom is 1.43 Å.
2.86 Å
1.43 Å

19. Atomic Radius
The trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom of the family.
Why?
With each step down the family, we add an entirely new PEL to the electron cloud, making the atoms larger with each step.

20. Atomic Radius The trend across a horizontal period is less obvious.
What happens to atomic structure as we step from left to right?
Each step adds a proton and an electron (and 1 or 2 neutrons).
Electrons are added to existing PELs or sublevels.

21. Atomic Radius
The effect is that the more positive nucleus has a greater pull on the electron cloud.
The nucleus is more positive and the electron cloud is more negative.
The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period.

22. Effective Nuclear Charge
What keeps electrons from simply flying off into space?
Effective nuclear charge is the pull that an electron “feels” from the nucleus.
The closer an electron is to the nucleus, the more pull it feels.
As effective nuclear charge increases, the electron cloud is pulled in tighter.

23. Atomic Radius
The overall trend in atomic radius looks like this.

24. Atomic Radius Here is an animation to explain the trend.
On your help sheet, draw arrows like this:

25. Shielding
As more PELs are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus.
The effective nuclear charge (enc) on those outer electrons is less, and so the outer electrons are less tightly held.

26. Ionization Energy This is the second important periodic trend.
If an electron is given enough energy (in the form of a photon) to overcome the effective nuclear charge holding the electron in the cloud, it can leave the atom completely.
The atom has been “ionized” or charged.
The number of protons and electrons is no longer equal.

27. Ionization Energy
The energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ)
The larger the atom is, the easier its electrons are to remove.
Ionization energy and atomic radius are inversely proportional.
Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.

28. Ionization Energy

29. Ionization Energy (Potential)
Draw arrows on your help sheet like this:

30. Electron Affinity What does the word ‘affinity’ mean?
Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ).
Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.

31. Electron Affinity
Electron affinity is exothermic if there is an empty or partially empty orbital for an electron to occupy.
If there are no empty spaces, a new orbital or PEL must be created, making the process endothermic.
This is true for the alkaline earth metals and the noble gases.

32. Electron Affinity
Your help sheet should look like this: + +

33. Metallic Character
This is simple a relative measure of how easily atoms lose or give up electrons.
Your help sheet should look like this:

34. Electronegativity
Electronegativity is a measure of an atom’s attraction for another atom’s electrons.
It is an arbitrary scale that ranges from 0 to 4.
The units of electronegativity are Paulings.
Generally, metals are electron givers and have low electronegativities.
Nonmetals are are electron takers and have high electronegativities.
What about the noble gases?

35. Electronegativity
Your help sheet should look like this:

36. Overall Reactivity
This ties all the previous trends together in one package.
However, we must treat metals and nonmetals separately.
The most reactive metals are the largest since they are the best electron givers.
The most reactive nonmetals are the smallest ones, the best electron takers.

37. Overall Reactivity
Your help sheet will look like this:

38. The Octet Rule
The “goal” of most atoms (except H, Li and Be) is to have an octet or group of 8 electrons in their valence energy level.
They may accomplish this by either giving electrons away or taking them.
Metals generally give electrons, nonmetals take them from other atoms.
Atoms that have gained or lost electrons are called ions.

39. Ions
When an atom gains an electron, it becomes negatively charged (more electrons than protons ) and is called an anion.
In the same way that nonmetal atoms can gain electrons, metal atoms can lose electrons.
They become positively charged cations.

40. She’s unhappy and negative.
Ions
Here is a simple way to remember which is the cation and which the anion:
This is Ann Ion.
This is a cat-ion.
She’s unhappy and negative.
He’s a “plussy” cat!

41. Ionic Radius Cations are always smaller than the original atom.
The entire outer PEL is removed during ionization.
Conversely, anions are always larger than the original atom.
Electrons are added to the outer PEL.

42. Cation Formation
Effective nuclear charge on remaining electrons increases.
Na atom
1 valence electron
Remaining e- are pulled in closer to the nucleus. Ionic size decreases.
11p+
Valence e- lost in ion formation
Result: a smaller sodium cation, Na+

43. Anion Formation
A chloride ion is produced. It is larger than the original atom.
Chlorine atom with 7 valence e-
17p+
One e- is added to the outer shell.
Effective nuclear charge is reduced and the e- cloud expands.

44. Periodic Table & Trends

45. History of the Periodic Table
1871 – Mendeleev arranged the elements according to: 1. Increasing atomic mass Elements w/ similar properties were put in the same row
1913 – Moseley arranged the elements according to: 1. Increasing atomic number Elements w/ similar properties were put in the same column

46. Group Names +3 -3 -2 H 1 He 2 Li 3 Be 4 B 5 C 6 N 7 O 8 F 9 Ne 10 Na
Alkali +1
Alkaline Earth
Metals +2 +3 -3 -2
Halogen -1
Noble Gases H 1 He 2 Li 3 Be 4 B 5 C 6 N 7 O 8 F 9 Ne 10 Na 11 Mg 12 Al 13 Si 14 P 15 S 16 Cl 17 Ar 18

47. S & P block – Representative Elements
METALS
NONMETALS
TRANSITION METALS
S & P block – Representative Elements
Metalloids (Semimetals, Semiconductors) – B,Si, Ge, As, Sb, Te (properties of both metals & nonmetals)
Columns – groups or families Rows - periods

48. 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)

49. Periodic Trends
Periodic Trends – patterns (don’t always hold true) can be seen with our current arrangement of the elements (Moseley)
Trends we’ll be looking at:
Atomic Radius
Ionization Energy
3. Electronegativity

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

51. 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)

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

53. 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

54. Ionic Radius Trend
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+)

55. Ionic Radius

56. 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?

57. Ionic Radius
How do I remember this??? The more electrons that are gained, the greater the increase in size. P-3 S-2 protons 15 protons 16 electrons 18 electrons 18 Which ion is smaller?

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

59. 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)

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

61. 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-

62. 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

63. 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

64. Part I – Atomic Size

65. Atomic Radius
Measures as distance from nucleus to nucleus and divided by 2.
Unit commonly used is pm
picometer= 10-12m
Example: iodine atomic radius 140pm

66. How does atomic radius change across a period?
It is smaller to the right.
Why?
More protons in the nucleus higher electrical force pulls electrons closer to nucleus.

67. How does atomic radius change down a group?
It is larger down the group.
Why?
Valence electrons are at higher energy levels and are not bound as tightly to the nucleus because they are screened or shielded ( pushed away) by other electrons in inner levels.

68. Note
There are some exceptions. Example column 13.

69. The Periodic Table and Atomic Radius

70. Example: Which is larger: a lithium atom or a fluorine atom?

71. Example: Which is larger: an arsenic atom or a sulfur atom?

72. Part II – Ionization Energy

73. Ionization energy
Ionization energy is the amount of energy needed to remove an electron from a gaseous atom.
First ionization energy –
Second ionization energy –

74. Ion Positive ion ---removal of electron
Negative ion--- addition of electron

75. How does ionization energy change down a group?
The first ionization energy decreases as you move down a group.
Why?
The size of the atom increases.
Electron is further from the nucleus.

76. How does ionization energy change across a period?
The first ionization energy increases as you move from left to right across a period.
Why?
Nuclear charge increases while shielding is constant.
Attraction of the electron to the nucleus increases.

78. Ionic size Metallic elements easily lose electrons.
Non-metals more readily gain electrons.
How does losing or gaining an electron effect the size of the atom (ion) ?

79. Positive ions
Positive ions are always smaller that the neutral atom. Loss of outer shell electrons.

80. Negative Ions
Negative ions are always larger than the neutral atom. Gaining electrons.

81. Ion size trends in periods.
Going from left to right there is a decrease in size of positive ions.
Starting with group 5, there is sharp increase followed by a decrease in the size of the anion as you move from left to right.

82. Ion size trends in columns.
Ion size increases as you move down a column for both positive and negative ions

83. Electronegativity: the ability of an atom in a bond to pull on the electron. (Linus Pauling)

84. Electronegativity
When electrons are shared by two atoms a covalent bond is formed.
When the atoms are the same they pull on the electrons equally. Example, H-H.
When the atoms are different, the atoms pull on the electrons unevenly. Example, HCl

85. Trends in Electronegativity
Electronegativity generally decreases as you move down a group.
Electronegativity of the representative elements (Group A elements) increases as you move across a period.

86. Electronegativities of Some Elements
Element Pauling scale F Cl O N S C H Na
Cs 0.7

87. Note Most electronegative element is F (EN 4.0)
Least electronegative stable element is Cs
(EN 0.7)

88. Summary Shielding increases Ionic size increases
Ionization energy decreases
Electronegativity decreases
Nuclear charge increases
Atomic radius increases
Shielding is constant
Atomic Radius decreases
Ionization energy increases
Electronegativity increases
Nuclear charge increases

89. Classwork Rank the following in terms of increasing atomic radius.
Carbon, Aluminum, Oxygen, Potassium

90. Rank the following elements by increasing electronegativity.
sulfur, oxygen, neon, aluminum

91. Why does flourine have a higher ionization energy?
Why do elements in the same family generally have similar properties?

92. Indicate whether the following properties increase or decrease from left to right across the periodic table.
A. atomic radius (excluding noble gases)
B. First ionization energy
C. Electronegativity

93. What trend in atomic radius occurs down a group on the periodic table
What trend in atomic radius occurs down a group on the periodic table? What causes this trend?
What trend in ionization energy occurs across a period on the periodic table? What causes this trend?

94. Circle the atom in each pair that has the largest atomic radius.
Al or B
Na or Al
S or O
O or F
Br or Cl
Mg or Ca

95. Circle the atom in each pair that has the greatest ionization energy.
Li or Be
Ca or Ba
Na or K
P or Ar
Cl or Si
Li or K

96. Define electronegativity.

97. Circle the atom in each pair that has the greatest electronegativity.
Ca or Ga
Br or As
Li or O
Ba or Sr
Cl m or S
O or S