1. Chapter 5 Atomic Structure and the Periodic Table
Charles Page High School
Dr. Stephen L. Cotton

2. Section 5.1 Atoms
OBJECTIVES:
Summarize Dalton’s atomic theory.

3. Section 5.1 Atoms
OBJECTIVES:
Describe the size of an atom.

4. History of the atom Not the history of atom, but the idea of the atom.
Original idea Ancient Greece (400 B.C.)
Democritus and Leucippus- Greek philosophers.

5. History of Atom Smallest possible piece? Looked at beach Made of sand
Atomos - not to be cut
Looked at beach
Made of sand
Cut sand - smaller sand

6. Another Greek Aristotle - Famous philosopher
All substances are made of 4 elements
Fire - Hot
Air - light
Earth - cool, heavy
Water - wet
Blend these in different proportions to get all substances

7. Who Was Right? Greek society was slave based.
Beneath famous to work with hands.
Did not experiment.
Greeks settled disagreements by argument.
Aristotle was more famous.
He won.
His ideas carried through middle ages.
Alchemists change lead to gold.

8. Who’s Next? Late 1700’s - John Dalton- England.
Teacher- summarized results of his experiments and those of others.
Dalton’s Atomic Theory
Combined ideas of elements with that of atoms.

9. Dalton’s Atomic Theory
All matter is made of tiny indivisible particles called atoms.
Atoms of the same element are identical, those of different atoms are different.
Atoms of different elements combine in whole number ratios to form compounds.
Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed.

10. Just How Small Is an Atom?
Think of cutting a piece of lead into smaller and smaller pieces
How far can it be cut?
An atom is the smallest particle of an element that retains the properties of that element
Atoms-very small: Fig. 5.2, p. 108
still observable with proper instruments: Fig. 5.3, page 108

11. Section 5.2 Structure of the Nuclear Atom
OBJECTIVES:
Distinguish among protons, electrons, and neutrons in terms of relative mass and charge.

12. Section 5.2 Structure of the Nuclear Atom
OBJECTIVES:
Describe the structure of an atom, including the location of the protons, electrons, and neutrons with respect to the nucleus.

13. Parts of Atoms J. J. Thomson - English physicist. 1897
Made a piece of equipment called a cathode ray tube.
It is a vacuum tube - all the air has been pumped out.

14. Thomson’s Experiment
Voltage source - +
Vacuum tube
Metal Disks

15. Thomson’s Experiment
Voltage source - +

16. Thomson’s Experiment
Voltage source - +

17. Thomson’s Experiment
Voltage source - +

18. - + Thomson’s Experiment
Voltage source - +
Passing an electric current makes a beam appear to move from the negative to the positive end

19. - + Thomson’s Experiment
Voltage source - +
Passing an electric current makes a beam appear to move from the negative to the positive end

20. - + Thomson’s Experiment
Voltage source - +
Passing an electric current makes a beam appear to move from the negative to the positive end

21. - + Thomson’s Experiment
Voltage source - +
Passing an electric current makes a beam appear to move from the negative to the positive end

22. Thomson’s Experiment
Voltage source
By adding an electric field

23. Thomson’s Experiment
Voltage source + -
By adding an electric field

24. Thomson’s Experiment
Voltage source + -
By adding an electric field

25. Thomson’s Experiment
Voltage source + -
By adding an electric field

26. Thomson’s Experiment
Voltage source + -
By adding an electric field

27. Thomson’s Experiment
Voltage source + -
By adding an electric field

28. Thomson’s Experiment
Voltage source + -
By adding an electric field he found that the moving pieces were negative

29. Other particles
Proton - positively charged pieces 1840 times heavier than the electron – by E. Goldstein
Neutron - no charge but the same mass as a proton – by J. Chadwick
Where are the pieces?

30. Rutherford’s experiment
Ernest Rutherford -English physicist. (1910)
Believed in the plum pudding model of the atom (discussed in Chapter 13).
Wanted to see how big they are.
Used radioactivity.
Alpha particles - positively charged pieces- helium atoms minus electrons
Shot them at gold foil which can be made a few atoms thick.

31. Rutherford’s experiment
When an alpha particle hits a fluorescent screen, it glows.
Here’s what it looked like (page 111)

32. Fluorescent
Screen
Lead block
Uranium
Gold Foil

33. He Expected
The alpha particles to pass through without changing direction very much.
Because…?
…the positive charges were thought to be spread out evenly. Alone they were not enough to stop the alpha particles.

34. What he expected

35. Because

36. He thought the mass was evenly distributed in the atom

37. Since he thought the mass was evenly distributed in the atom

38. What he got

39. How he explained it Atom is mostly empty.
Small dense, positive piece at center.
Alpha particles are deflected by it if they get close enough. +

40. +

41. Density and the Atom
Since most of the particles went through, it was mostly empty space.
Because the pieces turned so much, the positive pieces were heavy.
Small volume, big mass, big density.
This small dense positive area is the nucleus.

42. Subatomic particles – p.111
Relative
mass
Actual
mass (g)
Name
Symbol
Charge
Electron e- -1
1/1840
9.11 x 10-28
Proton p+ +1 1
1.67 x 10-24
Neutron n0 1
1.67 x 10-24

43. Section 5.3 Distinguishing Between Atoms
OBJECTIVES:
Explain how the atomic number identifies an element.

44. Section 5.3 Distinguishing Between Atoms
OBJECTIVES:
Use the atomic number and mass number of an element to find the numbers of protons, electrons, and neutrons.

45. Section 5.3 Distinguishing Between Atoms
OBJECTIVES:
Explain how isotopes differ, and why the atomic masses of elements are not whole numbers.

46. Section 5.3 Distinguishing Between Atoms
OBJECTIVES:
Calculate the average atomic mass of an element from isotope data.

47. Counting the Pieces Atomic Number = number of protons in the nucleus
# of protons determines kind of atom (since all protons are alike!)
the same as the number of electrons in the neutral atom.
Mass Number = the number of protons + neutrons.
These account for most of mass

48. Symbols
Contain the symbol of the element, the mass number and the atomic number.

49. Symbols
Contain the symbol of the element, the mass number and the atomic number.
Mass
number X
Atomic
number

50. F Symbols 19 9 Find the number of protons number of neutrons
number of electrons
Atomic number
Mass Number 19 F 9

51. Br Symbols 80 35 Find the number of protons number of neutrons
number of electrons
Atomic number
Mass Number 80 Br 35

52. Symbols
if an element has an atomic number of 34 and a mass number of 78 what is the
number of protons
number of neutrons
number of electrons
Complete symbol

53. Symbols if an element has 91 protons and 140 neutrons what is the
Atomic number
Mass number
number of electrons
Complete symbol

54. Symbols if an element has 78 electrons and 117 neutrons what is the
Atomic number
Mass number
number of protons
Complete symbol

55. Isotopes Dalton was wrong.
Atoms of the same element can have different numbers of neutrons.
different mass numbers.
called isotopes.

56. Naming Isotopes
We can also put the mass number after the name of the element.
carbon- 12
carbon -14
uranium-235

57. Atomic Mass How heavy is an atom of oxygen?
There are different kinds of oxygen atoms.
More concerned with average atomic mass.
Based on abundance of each element in nature.
Don’t use grams because the numbers would be too small.

58. Measuring Atomic Mass Unit is the Atomic Mass Unit (amu)
One twelfth the mass of a carbon-12 atom.
Each isotope has its own atomic mass, thus we determine the average from percent abundance.

59. Calculating averages
Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results.
Sample 5-5, p.120

60. Atomic Mass
Calculate the atomic mass of copper if copper has two isotopes. 69.1% has a mass of amu and the rest has a mass of amu.

61. Atomic Mass
Magnesium has three isotopes % magnesium 24 with a mass of amu, 10.00% magnesium 25 with a mass of amu, and the rest magnesium 25 with a mass of amu. What is the atomic mass of magnesium?
If not told otherwise, the mass of the isotope is the mass number in amu

62. Atomic Mass Is not a whole number because it is an average.
are the decimal numbers on the periodic table.

63. Section 5.4 The Periodic Table: Organizing the Elements
OBJECTIVES:
Describe the origin of the periodic table.

64. Section 5.4 The Periodic Table: Organizing the Elements
OBJECTIVES:
Identify the position of groups, periods, and the transition metals in the periodic table.

65. Development of the Periodic Table
mid-1800s, about 70 elements
Dmitri Mendeleev – Russian chemist
Arranged elements in order of increasing atomic mass
Thus, the first “Periodic Table”

66. Mendeleev Left blanks for undiscovered elements
When discovered, good prediction
Problems?
Co and Ni; Ar and K; Te and I

67. New way Henry Moseley – British physicist
Arranged elements according to increasing atomic number
The arrangement today
P.124 – long form
Symbol, atomic number & mass

68. Periodic table Horizontal rows = periods There are 7 periods
Periodic law:
Vertical column = group (or family)
Similar physical & chemical prop.
Identified by number & letter

69. Areas of the periodic table
Group A elements = representative elements
Wide range of phys & chem prop.
Metals: electrical conductors, have luster, ductile, malleable

70. Metals Group IA – alkali metals Group 2A – alkaline earth metals
Transition metals and Inner transition metals – Group B
All metals are solids at room temperature, except _____.

71. Nonmetals
Nonmetals: generally nonlustrous, poor conductors of electricity
Some gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br)
Group 7A – halogens
Group 0 – noble gases

72. Division between metal & nonmetal
Heavy, stair-step line
Metalloids border the line
Properties intermediate between metals and nonmetals
Learn the general behavior and trends of the elements, instead of memorizing each element property