1. Trends on the Periodic Table
Chemical Periodicity
Trends on the Periodic Table

2. Development of the Periodic Table
70 elements discovered by the mid-1800s
No system developed to organize them
Dimitri Mendeleev put the elements on cards
Grouped in columns by increasing atomic mass
Arranged the columns by similarities in the elements

3. Published in 1869
Predicted the missing elements

4. Mendeleev Left blanks on the table
Predicted the elements that would be there based on the trends
Eventually the elements were discovered
Found to have the properties predicted

6. The Modern Periodic Table
Henry Moseley, British physicist (1900’s)
Arranged the elements by atomic number instead of mass
Properties change as you move left to right, from element to element
Properties repeat in the next row.
Similar chemical properties in groups

7. Periodic Law
When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties
Electron configurations in groups are similar in their valence electron configurations

8. Group A Representative Elements
All have electrons in the s, or s and p orbitals
Groups number indicates number of valence electrons
Except Helium (group 8 with 2 electrons)

9. Group IA (group 1) Alkali Metals Soft, silvery color
Too reactive to be found pure or free form in nature
Hydrogen in NOT an alkali metal s1

10. Group IIA (group 2) Alkaline Earth Metals
Harder, denser and stronger than alkali metals
Higher melting points
Less reactive than alkali metals
Still too reactive to be found pure or free form in nature s2

11. Group VIII (group 18) Noble Gases 8 valence electrons (He =2)
Inert (nonreactive), all gases
Do not form ions
Rarely in compounds
Argon is most abundant (1% of air)
s2p6

12. Group VIIA (group 17) Halogens
Very reactive; react with metals to form salts
Non-metals
7 valence electrons
s2p5

13. Hydrogen Class by itself
Because it contains just one proton and one electron
Behaves unlike any other element
Reacts with many other elements
Most common element in the universe
1s1

14. Group B: Transition Metals
Usually only two valence electrons
Groups 3-12
Called the d block
Can lose one, two, or three valence electrons depending on the element that it react with
Form compounds and solutions that are brightly colored
Some are relatively inactive; gold, silver, platinum

15. Inner Transition Elements
Called the Rare Earth Metals
F-block
Lanthanide series; soft malleable metals with high lustre and conductivity

17. Metals

18. Metals
Inner Transition Metals

19. Nonmetals

20. Metalloids Elements touching the stair step Boron is a metaloid
Aluminum (acts more like a metal)
Have properties of metals and non-metals
Silicon and Germanium used for computer chips and solar cells

21. Metalloids

23. Classifying Elements by Electron Configuration
Periodic table useful for identifying properties of elements
Electron plays the most significant role in chemical reactions and properties of elements
There is a relationship between electron configuration, placement on periodic table and chemical and physical properties

24. Four Classes Noble Gases Representative Elements (main group)
Transition Elements
Inner transition elements

25. Periodic Table Blocks

27. Periodic Trends Atomic Radius Electronegativity (Electron Affinity)
Ionization Energy
Density
Melting Point
Boiling Point

28. Concepts & Terms Electron Shielding Effective Nuclear Charge
Increasing Principal Energy Level
Lowest Energy Level (closest to the nucleus)
Increasing mass (increasing atomic number)
Increasing/decreasing volume (electron cloud)

29. Atomic Radius: Group Trends
Increases as you move down a group
Why?
Increase in protons, but increase in principal energy levels (further away from the nucleus)
Electron shielding by the electrons in the energy levels closed to the nucleus = reduced effective nuclear charge

30. Atomic Radius: Periodic Trends
Atomic Radius DECREASES across a period
Why?
Principal Energy level remains the same (same electron shielding)
Nucleus is becoming more positive thus increasing the effective nuclear charge

31. Atomic Radius

32. Trends in Ionization Energy
When an atom gains or loses an electron, it becomes an ion.
The energy required to overcome the attraction of the nuclear charge and remove an electron from a gaseous atom is called the ionization energy.
Removing one electron results in the formation of a positive ion with a 1+ charge.
Na(g) Na+(g) + e-

33. The energy required to remove this first outermost electron is called the first ionization energy.
To remove the outermost electron from the gaseous 1+ ion requires an amount of energy called the second ionization energy, and so forth.

34. Trends in Ionization Energy

35. Ionization energies
Large increase in ionization energy between the first and second ionization energies in sodium
Easy to remove the first electron
Much harder to remove the second electron
Second electron is in the next lowest energy level
Group 2A large increase for the third electron (change in energy level)

36. Ionization Energy: Group Trends
Ionization Energy decreases as you move down a group
Why?
The size of the atom increases, energy level increases
Outermost electron is farther from the nucleus
Outermost electron in most easily removed.

37. Ionization Energy: Periodic Trends
First Ionization Energy increases as you move across the rows from left to right
Why?
The nuclear charge increases with increasing numbers of protons
The shielding effect is constant as the electrons are in the same principal energy level

39. Trends in First Ionization Energy

40. Cation vs Anion

41. Trends in Electronegativity
Electronegativity is the tendency for an atom to attract shared electrons to itself
The atom becomes more negative
Electronegativity decreases as you move down a group
Electronegativity increases as you move across a row from left to right

42. The metallic elements at the far left of the periodic table have low electronegativities. By contrast, the nonmetallic elements at the far right (excluding the noble gases), have high electronegativities.
The electronegativity of cesium, a metal, the least electronegative element, is 0.7; the electronegativity of fluorine, a nonmetal, the most electronegative element, is 4.0.
Because fluorine has such a strong tendency to attract electrons, when it is chemically combined to any other element it either attracts the shared electrons or forms a negative ion.
In contrast, cesium has the least tendency to attract electrons.

45. Electron Affinity

46. Periodic Trends

47. Periodic Table Puzzle Activity Key
Hair = Valence Electrons
Fingers = Protons/ Atomic Number
Arms = Period/energy levels
Pattern = group
Size = Radius (in groups only, does not work for a period)
Smile = electronegativity, electron affinity