Unit 5

1. The Periodic Table & Periodic
Law

2. Objective:
▫ Trace the development of the periodic table
▫ Identify key features of the periodic table

3. • Sheldon & The Element Song

4. Development of the periodic table
• - John Newlands
• - Meyer and Meedeleev

5. John Newlands
• 1864 English chemist
• -Proposed a scheme for the elements
• -He noticed that when elements where arranged
based on atomic mass their properties repeated
every 8th element.

6. John Newlands
• -Newlands named the periodic relationship Law
of octaves. (after a musical, a name which was
frowned upon because it was unscientific)
• -Although not accepted Newland’s correct in the
aspect that elements do repeat each other.

7. Meyer & Mendeleev
• Meyer-Russian chemist-1869
• Dmitri Mendeleev- 1869
▫ Demonstrated a connection between atomic mass
and elemental properties.

8. Meyer
▫ Demonstrated a connection between atomic mass
and elemental properties
▫ Arranged the elements in order of increasing
atomic mass

9. Mendeleev
▫ Demonstrated a connection between atomic mass
and elemental properties
▫ Arranged the elements in order of increasing
atomic mass
▫ Predicted the existence of properties of
undiscovered elements.

10. Mendeleev
▫ Mendeleev is given more credit.
▫ History of the Periodic Table

11. Periodic Table Activity
▫ Recognizing patterns

12. What did Mendeleev miss?
▫ New elements were discovered, it became evident
that the periodic table was not in the correct
order.
▫ Henry Moseley (1887-1915), arrange elements by
atomic number.

13. Moseley
▫ Discovered that atoms contain a unique number of
protons called the atomic number
▫ Arranged element in order of increasing atomic
number, which resulted in a periodic pattern of
properties.

14. 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.

15. The Periodic Table & Periodic
Law
Elements & Modern periodic table

16. • Element Song

17. Modern Periodic Table
• Groups= columns “Vertical”
• Periods= rows “horizontal”
▫ Periodic properties
 Periods have similar characteristics.

18. 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?

19. Alkali Metals
• Li
• Na
• K
• Rb
• Cs
• Fr

20. Alkali Metals
▫ 1 valance electron very reactive
 Alkali metals with water

21. Alkali metals
▫ Very reactive, often exists as
compounds with other elements
▫ Two familiar alkali metals are..
Na- used in table salt
Li- Used in batteries

22. Alkaline Earth
▫ 2 valance
electron
▫ Reactive-forms
oxides
▫ Important in
living things
 Found on
planet in raw
forms

23. Halogens
▫ Not “super”
stable
▫ 7 valance
electrons
 Need 1 more

24. Halogens
Cl- Halogen
Used as a gas
in WWI

25. Noble Gases ▫ Most un-reactive
▫ Why?
 8 valance
electrons

26. Inert gases…
▫ Is a gas which does not undergo chemical
reactions under a set of given conditions
 Mig welder-
uses argon

27. CHNOPSS-nonmetals “important to life”
▫ Carbon- that's what we are “made of”
▫ Hydrogen-your body is mostly H2O
▫ Nitrogen – amino acids
▫ Oxygen-energy out of food
▫ Phosphorous-DNA
▫ Sulfur-Protein
▫ Selenium –micro amounts/ a deficiency
is thought to cause cancer

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

29. Transition Metals
▫ “weird” number of electrons
▫ Cu-copper
▫ Au-Gold
▫ Ag-Silver

30. Titanium
▫ Strong and light it is often used
to make frames for bicycles and
eyeglasses.

31. “Poor” metals
▫ Al- Aluminum
▫ Ga- Gallium
▫ In – Indium
▫ Sn-Tin
▫ Ti- thallium
▫ Pb- lead
▫ Bi- Bismuth

32. “inner” transition metals
▫ Actinide series
▫ Lanthanide series
 used extensively as phosphors, substances that emit
light when struck by electrons

33. Metals
▫ Malleable & ductile, meaning they can be
pounded into thin sheets & drawn into wire
 Metals met the nonmetals, forming a stair step on
the right hand side of the periodic table.

34. 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?

35. “The disappearing Spoon”
• Disappearing Spoon
Gallium Low melting point
Highly radioactive= good bye hand

36. Metalloids
▫ B-Boron
▫ Si-Silicon
▫ Ge-Germanium
▫ As-Arsenic
▫ Sb-Antimony
▫ Te-Tellurium
▫ Po-Polonium
▫ At-Astatine

37. 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?

38. Metalloids
▫ Semiconductors
▫ Properties of both
metals and
nonmetals.

39. Metalloids
▫ Two important
metalloids
 Ge-germanium
 Si-Silicon
Used in computers
chips and solar cells

40. 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.

41. 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.

42. The Periodic Table & Periodic
Law
Classification of Elements

43. Objective:
▫ Explain why elements in the same group have
similar properties
▫ Identify the four block of the periodic table based
in their electron configuration

44. 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.

45. Valance electrons

46. Electron configuration

47. s, p, d, and f blocks

48. The Periodic Table & Periodic
Law
Periodic Trends

49. Atomic Radius
Definition: Half of the distance
between nuclei in covalently
bonded diatomic molecule
Radius decreases across a period
 Increased effective nuclear charge
due to decreased shielding
Radius increases down a group
 Each row on the periodic table
adds a “shell” or energy level to the
atom

50. Table of
Atomic
Radii

51. Explanation Video-Atomic Radii

52. Period Trend:
Atomic Radius

53. Ionization Energy
Definition: the energy required to remove
an electron from an atom
 Tends to increase across a period
 As radius decreases across a
period, the electron you are removing
is closer to the nucleus and harder to
remove
 Tends to decrease down a group
 Outer electrons are farther from
the nucleus and easier to remove

54. Periodic Trend:
Ionization Energy

55. Electronegativity
Definition: A measure of the ability of an
atom in a chemical compound to attract
electrons
o Electronegativity tends to increase
across a period
o As radius decreases, electrons get
closer to the bonding atom’s nucleus
o Electronegativity tends to decrease down
a group or remain the same
o As radius increases, electrons are
farther from the bonding atom’s nucleus

56. Periodic Table of Electronegativities

57. Periodic Trend:
Electronegativity

58. Summary of
Periodic Trends

59. Ionic Radii
 Positively charged ions formed when
an atom of a metal loses one or
Cations more electrons
 Smaller than the corresponding
atom
 Negatively charged ions formed
when nonmetallic atoms gain one
Anions or more electrons
 Larger than the corresponding
atom

60. 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.

61. 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.

62. 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 He’s a “plussy” cat!
and negative.

63. 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.

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

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

66. Graphic courtesy Wikimedia Commons user Popnose