The document summarizes the development and key features of the periodic table. It discusses how early chemists like Newlands, Meyer, and Mendeleev arranged elements in order of atomic mass and properties, recognizing repeating patterns. Mendeleev predicted new elements and Moseley later arranged elements by atomic number. The periodic table is organized into blocks by electron configuration and periods and groups that demonstrate trends in properties. Metals are on the left, nonmetals on the right, and metalloids border the stair-step line between them.

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