1. Unit 6 – The Periodic Table

2. Origins of the Periodic Table
By the year 1700, only 13 elements had been identified
Scientific discovery led to a higher rate of element discovery
A logical organization of elements was needed for all the new elements

3. Early Organization J.W. Dobereiner (1829) organized elements in triads
Triad – three elements with similar properties (ex: Cl, Br, I)
J.R. Newlands (1864) organized elements in octaves
Octave – repeating group of 8 elements

4. Mendeleev
Dmitri Mendeleev (1869) arranged elements according to their properties
Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, there was a repeating pattern to their properties
This is known as Periodicity
Mendeleev left some spaces on his table blank, but was able to predict the properties of the unknown elements

5. Mendeleev’s Periodic Table

6. Moseley Mendeleev’s table was imperfect – Te and I had to be reversed
Henry Moseley (1913) arranged elements according to atomic number
The periodic repetition of chemical and physical properties when elements are arranged by atomic number is known as the Periodic Law

7. Modern Periodic Table
The modern periodic table consists of Rows and Columns
Rows -
Horizontal
Also known as Periods
Numbered 1-7
Columns -
Vertical
Also known as Groups and Families
Numbered 1-18

8. Classifying Elements
The elements on the periodic table can be simply classified by groups
Groups 1,2,13-18 (1A-8A) are known as the Representative Elements

9. Classifying Elements
Groups of representative elements have the same valence electrons and Oxidation State
Oxidation State is how many electrons are gained or lost by an atom in a chemical reaction
Lost Electrons = Positive Oxidation State
Gained Electrons = Negative Oxidation State
Think of Oxidation State as the charge of the ion

10. Driving Force
Full Energy Levels require lots of energy to remove their electrons.
Noble Gases have full orbitals.
Atoms behave in ways to achieve noble gas configuration

11. Classifying Elements
Groups 3-12 (3B-2B) , as well as the lanthanide and actinide series are known as Transition Metals

12. Metals The most common class of elements is Metals
Metals become cations
What is a cation? How are they formed?
Positively charged atom/positive oxidation state - Lose electrons
Metals are generally solid (except Hg), conductive of heat and electricity, malleable, ductile, and shiny

13. Alkali Metals Group 1 elements are known as Alkali Metals
Alkali metals include Li, Na, K, Rb, Cs, Fr
Alkali metals are generally dull, soft, and reactive – rarely found as free elements

14. Alkali Metals How many valence electrons do all Alkali Metals have?
Write the noble gas configuration for each Alkali Metal
[He]2s1
[Ne]3s1
[Ar]4s1
[Kr]5s1
[Xe]6s1
[Rn]7s1
How many valence electrons do all Alkali Metals have?
One
What is the oxidation state of Alkali Metals? +1

15. Alkaline Earth Metals
Group 2 elements are known as Alkaline Earth Metals
Alkaline earth metals include Be, Mg, Ca, Sr, Ba, and Ra
Alkaline earth metals are harder, denser, and stronger than alkali metals
Less reactive than alkali metals, but still rarely found as free elements

16. Alkaline Earth Metals
Write the noble gas configuration for each Alkaline Earth Metal
[He]2s2
[Ne]3s2
[Ar]4s2
[Kr]5s2
[Xe]6s2
[Rn]7s2
How many valence electrons do all Alkaline Earth Metals have?
Two
What is the oxidation state of Alkaline Earth Metals? +2

17. Transition Metals
Elements in groups 3-12 (3B-2B) are known as Transition Metals
Transition metals include Mn, Fe, Ag, Au, Mo, etc.
Transition metals fill in the d orbital and often have multiple oxidation states
Lanthanide and Actinide Series elements fill in the f orbitals – known as inner transition elements

18. Metalloids
Elements that border the staircase on the periodic table are known as Metalloids
Metalloids include: B, Si, Ge, As, Sb, Te, Po, At
Metalloids have properties of both metals and nonmetals

19. Nonmetals
Nonmetals are found to the right of the staircase on the periodic table
Nonmetals generally become anions
What is an Anion? How are they formed?
Negatively charged atom/oxidation state - Gain electrons
Nonmetals are often gases or dull, brittle solids
Nonmetals generally show poor conductivity, ductility, and malleability

20. Halogens Group 17 elements are known as Halogens
Halogens include F, Cl, Br, and I
Halogens are the most reactive nonmetals – often found in compounds

21. Halogens Write the noble gas configuration for each Halogen [He]2s22p5
[Ne] 3s23p5
[Ar] 4s23d104p5
[Kr] 5s24d105p5
How many valence electrons do all Halogens have? Oxidation State?
Seven / -1
Why are the Halogens the most reactive non-metals?
They are 1 electron short of having an octet.

22. Noble Gases Elements in group 18 are known as Noble Gases
Noble Gases include He, Ne, Ar, Kr, Xe, Rn
Noble gases are extremely unreactive

23. Noble Gases How many valence electrons do all Noble Gases have? Eight
Write the electron configuration for each Noble Gas
1s2
[He]2s22p6
[Ne]3s23p6
[Ar]4s23d104p6
[Kr]5s24d105p6
[Xe]6s25d106p6
How many valence electrons do all Noble Gases have?
Eight
Why are Noble Gases so unreactive?
They contain a full octet – atoms gain/lose electrons to achieve noble gas notation

24. Other Groups
All other groups can be identified by the top most element in that group.
Ex: Group 15 can be called the Nitrogen Group
Oxidation State: -3
Q: What is another name for Group 16?
A: Oxygen group
Q: Oxidation State
A: -2

25. Periodic Trends
The elements on the periodic table show repeating trends related to electron configuration

26. What is the trend for Oxidation State?

27. Atomic Radius
The Atomic Radius is ½ the distance between nuclei of bonded atoms from the same element
Atomic radius decreases from left to right across a period
Atomic radius increases from top to bottom in a period

28. Why?
Not changing energy level, but increasing nuclear force (more positive charge in nucleus)

29. Ionization Energy
If an atom is becoming an ion, it is gaining or losing electrons in an effort to have an octet (8 valence electrons)

30. Ionization Energy
The energy required to remove an electron from an atom is called Ionization Energy

31. Ionization Energy
1st Ionization Energy- energy required to remove 1st electron from an atom
2nd Ionization Energy- energy required to remove 2nd electron from an atom
2nd Ionization Energy is ALWAYS higher than the 1st
3rd Ionization Energy- energy required to remove 3rd electron from an atom
3rd Ionization Energy is ALWAYS higher than the 1st or 2nd

32. Ionization Energy IE Decreases as you move down a group Why?
Electron is further away

33. Ionization Energy IE Increases as you move across a period Why?
You are in the same energy level but have more nuclear charge

34. Ionization Energy
Full Energy Levels require lots of energy to remove their electrons.
Noble Gases have full orbitals.
Atoms behave in ways to achieve noble gas configuration.

35. Ionization Energy Write the electron configuration for Be 1s22s2
How many valence electrons does Be have? 2
Why is the ionization energy low?
It is easier for Be to lose those 2 valence electrons than it is to gain 6. Therefore, it has a low ionization energy.

36. Ionization Energy
Move across the period. Write the electron configuration for F.
1s22s22p5
How many valence electrons does F have? 7
Why is the ionization energy high?
It is easier for F to gain 1 valence electron than is it for it to lose 7. Therefore, its’ ionization energy (energy to lose an electron) is high

37. Ionization Energy

38. Electron Affinity
Electron affinity is the energy change associated with adding an electron to a gaseous atom.
Easiest to add to group 7A (halogens).
Why?
Gets them to full octet.
Increase from left to right: atoms become smaller, with greater nuclear charge.
Decrease as we go down a group.

39. Ionic Size
Cations are smaller than the atoms from which they form (less electrons)
Anions are larger than the atoms from which they form (more electrons)

40. Ionic Size
Across the period, nuclear charge increases so they get smaller.
Energy level changes between anions and cations.
N3-
O2-
F1-
B3+
Li1+
C4+
Be2+

41. Electronegativity
Electronegativity is the ability for an atom to attract electrons in a compound
Electronegativity increases from left to right in a period
Electronegativity decreases from top to bottom in a group

42. Electronegativity
We do not consider noble gases when talking about electronegativity because they do not bond.
What is the most electronegative element?
Fluorine

43. Electronegativity Write the electron configuration for Li 1s22s1
How many valence electrons does Li have? 1
Why is the electronegativity low??
It is easier for Li to lose 1 valence electrons than it is to gain 7. It has a low electronegativity because it would be difficult for Li to attract 7 electrons

44. Electronegativity
Move across the period. Write the electron configuration for O.
1s22s22p4
How many valence electrons does O have? 6
Why is the electronegativity high?
It is easier for O to gain 2 valence electrons than is it for it to lose 6. Electronegativity is high because it can gain electrons more easily than it can lose them.

45. Electronegativity

46. Ionization energy, Electronegativity, and Electron Affinity INCREASE

47. Atomic size increases,
Ionic size increases