periodic trends chemistry chapter 3

1. How are
most
library
books
classified?
Why is such a
classification
system
useful?

2. Chapter 6 and 7
The Periodic Table
and Periodic Law
Harry Potter Sings the Element
Song
Jim Lehrer The Real Periodic
Table Song

3. 1789 - Lavoisier
Acid making
phosphorous
sulfur
carbon
Gas like
light
heat
oxygen
Azote (nitrogen)
hydrogen

4. 1789 - Lavoisier
Metallic
Copper, nickel, iron
Cobalt, mercury, tin
Gold, lead, silver, zinc
Earthy
Lime (calcium hydroxide)
Magnesia (magnesium oxide)
Manganese, tungsten
(platina) platinum
Barytes (barium sulphate)
Argilla (Aluminum oxide)
Silex (silicon dioxide)

5. I. Development of the Modern
Periodic Table 6.1 pg. 151-158
1. Compiled a list of 23 elements known
at the time.
2. 1800’s: changes in the world
3. Electricity used to break compounds,
spectrometer, industrial revolution
4. Tripled Lavoisier’s number of elements
5. 1860 chemists agreed on atomic masses
A. Lavoisier – 1790’s

6. B. John Newland - 1864
1. Proposed an organization scheme
2. Arranged by increasing atomic mass
3. The elements’ properties repeated
every eighth element
4. The law of octaves.
Musical
Analogy

7. B. John Newland - 1864

8. C. Meyer, Mendeleev, & Moseley
1. Meyer & Mendeleev each
demonstrated connections between
atomic mass and elemental properties.
2. Mendeleev published first and showed
the connections’ usefulness
3. Mendeleev predicted the existence
and properties of undiscovered
elements.
4. He left blanks for undiscovered
elements
Can you find the errors in his system?

9. C. Meyer, Mendeleev, & Moseley

10. BBC History of the
Periodic Table

11. "...if all the elements be arranged in order of their atomic
weights a periodic repetition of properties is obtained." -
Mendeleev
Mendeleev’s First Periodic Table

12. 5. Henry Moseley (1913):
British chemist - discoveries resulted in a
more accurate positioning of elements
by determination of atomic numbers.
(Tragically for the development of
science, Moseley was killed in action at
Gallipoli in 1915).

13. 5. Moseley (continued)
5. When atoms were arranged according to
increasing atomic number, the few problems
with Mendeleev's periodic table had
disappeared. Because of Moseley's work, the
modern periodic table is based on the atomic
numbers of the elements
6. Periodic Law: There is a periodic repetition of
the chemical and physical properties of the
elements when they are arranged by increasing
atomic number.

14. 5. Moseley

15. II. The Modern Periodic Table

16. A. Columns on the periodic table are
called groups or families.
Atomic number increases as you
move down on the periodic table.
Each group is numbered one though
eight, followed by the letter A or B.
II. The Modern Periodic Table
What group is Chlorine in?
VIIA

17. B. The rows on the periodic table are
called periods.
Each row on periodic table (except the
first) begins with a metal and ends with
a noble gas.
Beginning with hydrogen in period 1
there are a total of seven periods. In
between, the properties of the
elements change in an orderly
progression from left to right.
The pattern in properties repeats after
group VIIIA (18).
II. The Modern Periodic Table

18. C. Why does the first period on the
periodic table only have two elements?
Only two electrons can occupy the first
energy level in an atom.
The third electron in lithium must be at
a higher energy level.
Lithium is the first element in Group IA
and in Period 2.
II. The Modern Periodic Table

19. D. The groups designated with an A (IA
through VIIIA) are often referred to as
the main group or representative
elements because they possess a wide
range of chemical and physical
properties.
The groups designated with a B (IB
through VIIIB) are referred to as the
transition elements.
II. The Modern Periodic Table

20. Vanadium a transition element.
II. The Modern Periodic Table

21. III. Classification of the Elements
A.Valence Electrons
1. electrons in the highest principal
energy level
2. What would group IA look like?
3. Atoms in the same group have
similar chemical properties because
they have the same number of
valence electrons.
B. The s-, p-, d- and f- block elements
1s1, 2s1….

22. III. Classification of the Elements

23. A. Metals are elements that have luster,
conduct heat and electricity, and
usually bend without breaking, and are
solid at room temperature. With the
exception of tin, lead, and bismuth,
metals have one, two, or three valence
electrons.
IV. Classification of the Elements

24. The Metals

25. Metals are to the left of the heavy
stair-step line that zigzags down
from Boron (B, in column IIIA) to
astatine (At) at the bottom of group
VIIA (hydrogen is the exception).
Alkali metals – group IA, the most
reactive of all metals. They react with
water to form alkaline solutions.
IV. Classification of the Elements

26. Alkali metals

27. Alkaline earth metals group IIA.
These elements form compounds
with oxygen, called oxides.
All transition elements are metals.
IV. Classification of the Elements

28. All transition elements are metals.
IV. Classification of the Elements

29. Inner transition metals are know as
lanthanide and actinide series and
are located along the bottom of the
periodic table. Because of their
natural abundance on Earth is less
than 0.01 percent, the lanthanides
are sometimes called the rare earth
elements. All of the lanthanides
have similar properties.
IV. Classification of the Elements

30. All of the actinides are radioactive,
and none beyond uranium (92) occur
in nature.
IV. Classification of the Elements

31. Dr. Glenn Seaborg on Transuranium Elements
Meet Dr. Glenn Seaborg
0:59-4:43

32. B. Nonmetals are elements that are
generally gases or brittle, dull-looking
solids. They are poor conductors of heat
and electricity.
The only liquid nonmetal at room
temperature is bromine (Br).
The highly reactive group VIIA is the
halogens, “salt formers.” The extremely
un-reactive group is the noble gases.
The nonmetals oxygen and nitrogen make
up 99 percent of Earth’s atmosphere.
III. Classification of the Elements

33. Carbon, another nonmetal, is found in
more compounds than all the other
elements combined.
Their melting points tend to be lower
than those of metals.
III. Classification of the Elements

35. C. Metalloids or semimetals.
Metalloids are elements with
physical and chemical properties
of both metals and nonmetals.
Silicon and germanium are two of
the most important metalloids as
they are used extensively in
computer chips and solar cells.
III. Classification of the Elements

36. Some metalloids such as silicon,
germanium (Ge), and arsenic
(As) are semiconductors.
III. Classification of the Elements

37. A semiconductor is an element that
does not conduct electricity as well
as a metal, but does conduct slightly
better than a nonmetal.
The ability of a semiconductor to
conduct an electrical current can be
increased by adding a small amount
of certain other elements.
III. Classification of the Elements

38. Silicon’s semi conducting properties
made the computer revolution
possible.
III. Classification of the Elements
Your television,
computer, handheld
electronic games, and
calculator are electrical
devices that depend on
silicon semiconductors.

39. All have miniature electrical circuits
that use silicon’s properties as a
semiconductor.
You learned that metals generally
are good conductors of electricity,
nonmetals are poor conductors, and
semiconductors fall in between the
two extremes.
III. Classification of the Elements

40. The ability of a semiconductor to
conduct an electrical current can be
increased by adding a small amount
of certain other elements.
Silicon’s semi conducting properties
made the computer revolution
possible.
Your television, computer, handheld
electronic games, and calculator are
electrical devices that depend on
silicon semiconductors.
III. Classification of the Elements

41. III. semiconductors
Awesome New Technology –
Cool Gadgets

42. What clues does the arrangement of the football players on the
field give about the functions of their positions?

43. What characteristics does a football player have based on the
position played?

44. V. Periodic Trends 6.3 pgs. 163-169
A. Atomic Size or Radius
1. How closely an atom lies to a
neighboring atom.
2. Metals – half the distance between two
adjacent atoms in a crystal
3. Nonmetal – determined from a diatomic
molecule of an element

45. Time
out!
What’s a
diatomic
element?

46. Time out! What’s a diatomic
element?
Mr. Brinklehoff
BrINClHOF
gases
N2, H2, F2…
an element that when not
chemically bonded with any other
elements, will form a molecule
having two atoms of the element.

47. A. Atomic Radii - trend (cont.)
Size increases going across – right to left.
(Across same energy level, but add protons
and the ‘pull’ of nucleus is greater and pulls
electrons closer.)
Size increases going down a group. Electrons
are at progressively higher levels, and shielding
effect is increasing.

48. A. Atomic Radii (continued)
Shielding Effect – lots of inner electrons
shield or protect the outer electrons from
the ‘pull’ of the positive nucleus.
What happens to shielding effect as
you go down the periodic table?
increases
What happens to shielding effect as
you go across the periodic table?
Stays the same
Quick time video

49. B. Ionic Radius
Positive ions are always smaller than
neutral atoms
+1 +2 +3
-1 -2
-3
Negative ions are always larger than
neutral atoms because their nuclear
attraction is less

50. Going down a group – both anions and
cations get bigger
Going across: + decrease, - increase
Smaller Larger

51. Atomic and ionic radii trends

52. C. Ionization Energy
Energy required to remove an
electron from a gaseous atom.
Think of ionization energy as an
indication of how strongly an atom’s
nucleus holds onto its valence
electrons.
Bigger your ionization energy the
harder it is to rip an electron away.
How easy is it to become an ion?

53. Ionization energy

54. I.E. increases going up (harder to pull off outer
electrons when atoms are smaller)
I.E. increases going across (harder to pull off
electrons because of nuclear attraction is
hanging onto them.
Metals have lower I.E., nonmetals have high I.E.
High High

55. Successive ionization energies
Quick time video
Video

56. D. Electronegativity
indicates the relative ability of an
atom to attract electrons in a
chemical bond “hogs electrons”
The bigger the electronegativity
the bigger the ‘hog’

57. Polar Covalent Bonds
Though atoms often form
compounds by sharing
electrons, the electrons
are not always shared
equally.
• Fluorine pulls harder on the electrons it
shares with hydrogen than hydrogen does.
• Therefore, the fluorine end of the molecule
has more electron density than the hydrogen
end.

58. Electronegativity
Electronegativity is the
ability of atoms in a
molecule to attract
electrons to themselves.
On the periodic chart,
electronegativity increases
as you go…
…from left to right across a
row.
…from the bottom to the
top of a column.

59. D. Electronegativity (continued)
moving across, electronegativity increases
moving up electronegativity increases
high highest
Video

61. Following Slides not covered in chapter 6
Glencoe

62. E. Electron Affinity
“Electron Grabbiness” or how easy it is
for an atom to gain electrons (make a
negative ion = anion)
(-) negative = energy released
(+) positive = energy absorbed
Video

63. Electron Affinity

64. E. Electron Affinity (continued)
Generally decreases going down a group
because atomic size increases
Increases going across because atoms are
smaller and nucleus charge increases
Video

65. F. Melting Points
metallic side – decreases as you go down
non-metals- increase going down
melting pt of gases are very low (-100 C)

66. G. Metallic Character
increase to left and down least
metallic
most
metallic

67. Metallic Character

68. H. Chemical Activity
metals – increase going down, most active? Cs
non metals – decreases going down, most active?
F