1. Periodic Table

2. I. History Stanislao Cannizzaro (1826-1910) (Italian Chemist)
1. reliable method to measure atomic
masses
B. Johann Dobereiner ( )
(German Scientist)
John Newlands ( English)
1. related atomic mass to properties
2. Newland’s Law of Octaves

3. John Newlands – Law of Octaves

4. John Newland’s Law of Octaves

5. Lothar Meyer (1835-1895 - German)
1. properties of elements show a
repetitive pattern when they are
arranged by atomic mass
D.Dimitri Mendeleev ( Russian)
(father of modern periodic table)
published system used today (1869)
2. elements arranged by increasing mass
3. left spaces for elements not yet
discovered - predicted properties

6. Dimitri Mendelev

7. Mendeleev’s Table His table re-organized

8. video

9. E. Henry Mosley (1887-1915) English
1.Arrange elements
by increasing atomic number – this led to the-periodic law
2. Periodic Law -
properties of elements
are periodic functions
of their atomic number

10. II. Arrangement of Elements
Periodic Table – arrangement of elements in order of increasing atomic number so that elements with similar properties are in the same column
1. period – horizontal row (7)
2. group(family)- vertical columns (1-18)
3. periodicity – reoccurrence of similar
properties of elements in groups

12. C. Special Groups on the Periodic Table
Group # and Name
1 - Alkali Metals
2 - Alkaline Earth
Metals
15 - Nitrogen
Family
- Oxygen
Halogens
Noble Gases

13. E. Metals – Metalloids - Nonmetals
1. Metals are on the left side – all are solids
except mercury (Hg)
a. elements near the left of a period are more metallic than those near the right
b. elements near the top of a group are more metallic than those near the bottom
2. Metalloids – group of elements between
metals and nonmetals(B,Si,Ge,As,Sb,Te)
3. Nonmetals are on the right side – all are
solids or gases except bromine(Br) liquid

14. Metals – Metalloids - Nonmetals

15. PROPERTY METAL NONMETAL
Luster high low
Deformability malleable brittle
and ductile
Conductivity good poor
Electron gain/lose lose gain
Ion formed cation (+) anion(-)
Ionization energy low high
Electronegativity low high

17. I. History of the Atomic Theory
Modern Atomic Theory
All matter is made up of small particles called atoms.
Atoms of the same element have the same chemical properties while atoms of different elements have different properties
(isotopes)
Not all atoms of an element have the same mass, but they all have a definite average mass which is characteristic.

18. I. History of the Atomic Theory
Atoms of different elements combine to form compounds and each element in the compound loses its characteristic properties.
Atoms cannot be subdivided by chemical or physical changes.
(nuclear reactions)

19. I. History of the Atomic Theory-time line
1803
1897
1909
1913
1935
Today
solid
particle
electron
proton
e- orbit nucleus
neutron
Quantum Atom theory
Dalton
Thomson
Rutherford
Bohr
Chadwick
Schrodinger and others

20. II. History of the Atomic Structure
J.J. Thomson (1887)
Cathode Ray Tube
Discovered matter contained negative charge
Electron
e -

21. Thomson’s Model Found the electron Said the atom was like plum pudding
A bunch of positive stuff (pudding), with the electrons suspended (plums)

22. II. History of the Atomic Structure
Robert Milikan (1909)
Oil Drop Experiment
Discovered mass and actual charge of electron
Mass is 1/1840 the mass of a hydrogen atom
e – has a mass of 9.11 x g

23. II. History of the Atomic Structure
So, at this point we know:
Dalton’s Atomic Theory
Electrons are negatively charged
The mass of an electron is very small

24. The main points of Dalton's atomic theory were:
Elements are made of extremely small particles called atoms.
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
Atoms cannot be subdivided, created, or destroyed.
Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
In chemical reactions, atoms are combined, separated, or rearranged.

25. II. History of the Atomic Structure
But:
Atoms are neutral, so there must be a positive charge.
Electron’s mass is so small…there must be more to an atom.

26. II. History of the Atomic Structure
Ernest Rutherford (1909)
Gold Foil Experiment
Discovered the proton p+
When alpha (+2) particles hit screen, the screen lights up
P.S.
#p+ = atomic number

27. Florescent
Screen
Lead block
Uranium
Gold Foil

28. He expected:
The alpha particles to pass straight through

29. He thought this would happen:

30. Because

31. Because, he thought the mass of the positive charge was evenly distributed in the atom

33. Here is what he observed:

34. So, he noticed: Most positive alpha particles pass right through
Only a few were deflected

35. He reasoned:
If a + particle hit a + point on the foil, it was repelled and deflected

36. He concluded: Atom is mostly empty space
Has a small, dense positive center

38. II. History of the Atomic Structure
At this point in 1909, we know:
p+ = 1.67 x g
e- = 9.11 x g
The charges are balance!
But,
How are the electrons arranged?
There is still mass that is unaccounted for

39. II. History of the Atomic Structure
Niels Bohr (1913)
Electrons orbit nucleus in predictable paths
Much more on him later

40. II. History of the Atomic Structure
Chadwick (1935)
Discovers neutron in nucleus
Neutron is neutral n0
Mass is 1.67 x g

41. II. History of the Atomic Structure
Charges balanced
Mass accounted for
But today, we subscribe to the Quantum Atom Theory to describe the atomic structure

42. II. History of the Atomic Structure
Quantum Atom Theory
The atom is mostly empty space
Two regions:
Nucleus- protons and neutrons
Electron cloud- region where you have a 90% chance of finding an electron

43. III. Subatomic Particles
Relative
mass
Actual
mass (g)
Name
Symbol
Charge
Electron e- -1
9.11 x 10-28
Proton p+ +1
1amu
1.67 x 10-24
Neutron n0
1amu
1.67 x 10-24

44. Positive-negative attraction electric force between protons in one atom and electrons in another atom hold atoms together  chemical bond

45. III. Subatomic Particles
Atomic number
The number of protons in the nucleus of an atom
Identifies the element
No two elements have the same atomic number
Mass number
The number of protons plus neutrons in the nucleus of an atom
Mass number is very close to the mass of an atom in amu (atomic mass units)
Two atoms with the same atomic number but different mass number are called isotopes
(mass number) – (atomic number) = #n 0 (number of neutrons)
Electrons and Ions
For neutral atoms, #e- = #p+
If there are more electrons, a negative ion forms
If there are less electrons, a positive ion forms
For now, we will work only with neutral atoms

46. III. Subatomic Particles
Electrons and Ions
For neutral atoms, #e- = #p+
If there are more electrons, a negative ion forms
If there are less electrons, a positive ion forms
For now, we will work only with neutral atoms

47. III. Subatomic Particles
You can never change the number of protons and have the same element
If you change the number of neutrons in an atom, you get
An isotope
If you change the number of electrons in an atom, you get
An ion

48. X III. Subatomic Particles Notation
Nuclear Notation- how we depict isotopes
contains the symbol of the element, the mass number, and the atomic number
Mass number
# of P +N X
Atomic number
# of P

49. Na III. Subatomic Particles 23 11 Notation Nuclear Notatioin
How many protons?
How many neutrons?
How many electrons? 23 Na 11

50. III. Subatomic Particles
Hyphenation Notation
Symbol or name of element – mass number
Fluorine-19
Protons? Neutrons? Electrons?
C-12

51. III. Subatomic Particles
Average Atomic Mass
Measured in grams (for a lot)
Measured in amu (for a few)

52. Atomic Mass
Mass of an atom
Too small to measure in grams
Use relative mass (amu)
Almost the same as mass number
Standard: 1 amu is defined as 1/12 the mass of one C-12 atom

53. III. Subatomic Particles
Average Atomic Mass
Weighted average mass of all known isotopes
Weighted means that the frequency of an isotope is considered
blackboard

54. 1. Horizontal rows are called periods
A. There are 7 periods

55. 2. Vertical columns are called groups.
a. Elements are placed in columns by similar properties.
i. Also called families

56. 3. The elements in the A groups are called the
representative elements
8A0 1A 2A 3A 4A 5A 6A 7A

57. 4. The group B are called the transition elements
a. These are called the inner transition elements and they belong here- lanthanide actinide series

58. 5. Group 1A are the alkali metals
6. Group 2A are the alkaline earth metals

59. 7. Group 7A is called the Halogens
8. Group 8A are the noble gases

60. The part of the atom another atom sees is the electron cloud.
More importantly the outside orbitals.
The orbitals fill up in a regular pattern.
The outside orbital electron configuration repeats.
The properties of atoms repeat.

61. Classification of the elements
A. Periodic Table- arrangements of elements so that elements with similar properties are in the same column
1. Valence electrons- electrons that exist on the highest principal energy level of an atom
a. these are the e- that can be lost or gained in formation of a compound

62. 2. energy level and period
a. the energy level of valence electrons is indicated by which period it is found in.
Ie: period 4 elements have valence electrons on the 4th energy level
Lets work through a few.

63. 3. Valence electrons and group number
a. Representative element’s group number and number of valence electrons it contains also are related.
Ie: Na has one e- on valence shell
Lets work on some

64. D. Periodic Table Showing s,p,d,f Blocks