Chemistry unit 3

1. Unit 3: TheUnit 3: The
Periodic TablePeriodic Table
Chapter 6Chapter 6

2. Lesson1: Development of the Periodic Table
1) History of the Periodic Table – By the end
of the 1700’s scientists had identified only 30
elements
(ex. Cu, Ag, Au, H2, N2, O2, C). By the mid
1800’s by using spectroscopy additional
elements were identified by using their line
spectra and about 65 elements had been
identified.

3. A. J.W. Dobereiner: 1829
Organized the elements into
groupsgroups with similar propertiesproperties.
He called these groups triadstriads.
The middle element is often
the average of the other two.
Ex) Cl – 35.5
Br – 79.9
I – 126.9
Ca
Avg Sr
Ba
Cl + I
Avg.
2
=

4. Triads on the Periodic Table

5. B. J.A.R. Newlands: 1867
• Law of OctavesOctaves.
He said properties
repeated every 88thth
element. There
were 62 known
elements at the
time. He was also
a musician.

6. C. Dimitri Mendeleev: 1869
1. Organized the 1st
periodic table
according to
increasing atomicatomic
massmass and put
elements with
similar properties in
the same columncolumn.

7. 2. He arranged some
elements out of atomic
mass orderorder to keep
them together with
other elements with
similar propertiesproperties. He
also left three blanksblanks in
his table and correctly
predictedpredicted the properties
of these 3 unidentified
elements that were later
identified and matched
his predictions.

8. D. Moseley: 1915
Each element has a certain amount
of positive charge in the nucleus
which are called protonsprotons.
1. Moseley reorganized the periodic
table by Atomic NumberAtomic Number.
2. The Periodic Law: When elements
are arranged in order of increasing
atomic numberatomic number, their physical and
chemical propertiesproperties show a periodicperiodic
pattern.

9. Glenn Seaborg “Seaborgium”
Sg #106• Born in 1912 in Michigan,
Seaborg proposed
reorganizing the Periodic
Table one last time as a
young chemist working on
the Manhattan Atomic Bomb
Project during WWII by
pulling the “f-block” elements
out to the bottom of the table.
He was the principle or co-
discoverer of 10 transuranium
elements. He was awarded
the Noble prize in 1951 and
died in 1999.

10. 6-2: Reading the Periodic Table
A. Information in each square:
B. Organizing the Squares
1. Vertical Columns – groups & families
2. Horizontal Rows - periods
13
Al
Aluminum
26.9815
Atomic # (protons)
Symbol
Name
Atomic Mass
Weighted average of all an element’s isotopes
Mass # - protons & neutrons
particular isotope of
aluminum like
Al-26 or Al-27

11. Parts of the Periodic Table

12. E. Metal, Nonmetals and Metalloids (Semimetals):
1. Metals
2. Nonmetals
3. Metalloids - Properties of both metals & nonmetals
- good conductors of heat & electricity
- most solids at room temperature (except
Hg)- high luster (shiny)
- ductile (can be drawn into thin wire)
- malleable (bends without breaking)
- high melting points
- high densities
- react with acids
- brittle (shatters when struck)
- low luster (dull)
- neither ductile nor malleable
- nonreactive with acids
- nonconductors
(Semimetals)

14. C. Electron Configuration & Families
1. Valence electrons –
outermostoutermost electrons
responsible for
bonding.
2. Elements in the
same groupgroup have
the same number
of valence
electrons.
Carbon has 4 valence electrons

15. Atomic Families:
Alkali Metals*
Alkaline Earth Metals*
Nobel Gases*
Halogens*
Oxygen Family
Carbon Family
Nitrogen Family
Boron Family
Transition Metals*
*Valance electrons?
*Physical Properties?
*Chemical Properties?
http://www.learner.org/interactives/periodic/groups.html

17. 6-3: Trends in the Periodic Table
A. Atomic Radius
1. The distance from
the center of thethe center of the
nucleus to thenucleus to the
outermost electronoutermost electron.
2. Atoms get largerlarger
going down a group
and smallersmaller going
across a period.
Ex) Na is larger than Mg
Na is smaller than K

18. Atomic Radii of the
Representative Elements

19. Atomic Radii vs Atomic Number

20. Ionic Size
1. When atoms gaingain
electrons, they
become (-) and get
larger.

21. Positive Ion Size
1. When atoms
loselose electrons,
they become
(+) and get
smallersmaller.
2. Ions get largerlarger
as you go
down a group.

22. Relative Sizes of Positive &
Negative Ions
The sodium ion lost an electron, and
therefore the positive protons in the
nucleus exert a stronger pull on the
remaining negative electrons,
shrinking the orbitals. Thus positive
ions are smaller than their atoms.
The chloride ion gained an electron, and
therefore the fewer positive protons in
the nucleus exert a weaker pull on the
extra negative electrons, increasing the
size of the orbitals. Thus negative ions
are larger than their atoms.

23. Electron Attraction in a Bond &
Ion Size

24. C. Ionization Energy:
1. The energy needed to removeremove one electron
from an atom.
2. Elements that do not want to lose their
electrons have highhigh ionization energies.
3. Elements that easily lose electrons have lowlow
ionization energies.
4. I.E. decreasesdecreases down a group (opposite of
atomic radius)
5. I.E. increasesincreases across a period. (opposite of
atomic radius)
Ex) Na IE smaller than Mg
Na IE larger than K

26. Ionization Energy of the
1st
20 Elements

27. Ionization Energy vs. Atomic Number

28. Sublevels by the Periodic Table

29. Which element would have the lowest ionization energy?Which element would have the lowest ionization energy?
Which element would have the highest ionization energy?Which element would have the highest ionization energy?
Will the Lithium ion loseWill the Lithium ion lose
any more electrons?any more electrons?

30. D. Successive Ionization Energies:
1. Energy required to removeremove electrons beyond the 1st
electron.
2. Ionization energies will increaseincrease for every electron removed.
3. Na [Ne]3s1
Na• 1st = ____ kJ 2nd = ____ kJ
4. Mg [Ne]3s2
Mg: 1st = ____ kJ 2nd = ____ kJ 3rd = ____kJ
5. Al [Ne]3s2
3p1
Al: 1st = ____kJ 2nd = ____kJ 3rd = ____kJ 4th = ___kJ
738
4560496
1450 7730
577 1814 2742 11,600

31. E. Electron Affinity:
1. Energy change that occurs when an atom gainsgains
an electron.
2. A highly negativenegative electron affinity attracts
electrons. (nonmetals)
3. A positivepositive electron affinity does not attract
electrons. (metals)

32. Electron Affinity
decrease
s

33. F. Electronegativity:
1. Reflects an atoms ability to attractattract electrons
in a chemical bond.
2. E.N. decreasesdecreases going down a group
3. E.N. increasesincreases going across a period.
4. Examples: NaCl and H2

36. G. Octet Rule:
1. An atom will tend to loselose, gaingain or shareshare
electrons in order to acquire a full set of
valence electrons.
2. “Octet” = 8 = s2
p6
configuration
H. Oxidation Number:
The charge on an ion when it gains or
loses electrons to acquire a stable octet.

37. Which of the following would have the largest?
• Atomic Radii?
• Ionization Energy?
• Electron Affinity?
Element D
Element C
Element A

38. Which of the following would have the smallest?
• Atomic Radii?
• Ionization Energy?
• Electron Affinity?
Element B
Element D
Element D

39. Electron Shielding