1. How do we know that atoms exist?
Atomic Structure
How do we know that atoms exist?

2. Democritus Greek philosopher ~ 400 B.C.
First person to believe that all matter is composed of tiny indivisible particles
Atomos – not to be cut

3. Aristotle More famous Greek philosopher
Did not believe Democritus was correct
All matter is continuous
All matter is made up of 4 elements
Earth
Air
Water
Fire

4. Who Was Right? Neither view was supported by experimentation.
Greeks settled disagreements by debate.
Aristotle was a better debater - He won.
His ideas carried through middle ages.

5. Foundations of Atomic Theory
Late 1700’s – most chemists accepted the modern definition of an element as a substance that cannot be broken down further by ordinary chemical means.
Elements combine to form compounds whose properties are different from the elements that form them.

6. Foundations of Atomic Theory
Great controversy over whether elements always combine in the same ratio when forming a particular compound.
Significant improvements to the available technologies

7. Foundations of Atomic Theory
Led to a more quantitative study of elements, compounds, and chemical reactions
Led to the discovery of several basic laws

8. Law of Conservation of Mass
Mass is neither created nor destroyed during ordinary chemical reactions or physical changes.
Antoine Lavoisier – French chemist

9. Law of Definite Proportions
Proust’s Law
Regardless of where or how a pure chemical compound is prepared, it is composed of a fixed proportion of. elements.
It is a ratio by mass.
All salt crystals, NaCl, regardless of sample size contains exactly 39.34% sodium and 60.66% chlorine.

10. Law of Multiple Proportions
If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.
Sometimes called Dalton’s Law

11. Law of Multiple Proportions
1.00g of carbon combines with 2.66g of oxygen in carbon dioxide and 1.33g of oxygen in carbon monoxide.
2.66/1.33 = 2 to1
The ratio of the ratios is a whole number.

12. John Dalton English school teacher
Proposed an explanation for the law of conservation of mass, law of definite proportions and law of multiple proportions.
Elements are composed of atoms and that only whole numbers of atoms can combine to form compounds.

13. Dalton’s Atomic Theory
1) All matter is composed of extremely small particles called atoms.
2) Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.

14. Dalton’s Atomic Theory
3) Atoms cannot be subdivided, created, or destroyed.
4) Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
5) In chemical reactions, atoms are combined, separated, or rearranged.

15. Where Dalton Was Wrong
Atoms are divisible into even smaller particles!
We now know that atoms are composed of subatomic particles.
A given element can have atoms with different masses!
Isotopes – atoms with the same number of protons but different numbers of neutrons.

16. The Structure of the Atom
Late 1800’s scientific advances allowed for a deeper exploration into the nature of matter.
Atoms are composed of several basic types of smaller particles.
The number and arrangement of these particles within an atom determines that atom’s chemical properties.

17. J. J. Thomson English physicist. 1897
Investigated relationship between electricity and matter.
Made a piece of equipment called a cathode ray tube.
It is a vacuum tube - all the air has been pumped out and a limited amount of other gases are put in

18. Thomson’s Experiment
Voltage source - +
Metal Disks

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

20. Thomson’s Experiment
Voltage source + -
By adding a magnetic field

21. Thomson’s Experiment Voltage source + - By adding a magnetic field
By adding a magnetic field he found that the moving pieces were negative.

22. Thomson’s Experiment Used many different metals and gases
Beam was always the same
By the amount it bent he could find the ratio of charge to mass
He found the ratio was the same with every material
He concluded that the rays were composed of identical negatively charged particles.

23. Thomsom’s Plum Pudding Model
Said the atom was like plum pudding.
For us, like seeds in a watermelon
A bunch of positive stuff, with the electrons able to be removed.

24. Robert Millkan
Robert Millikan measured the charge of an electron.
- Scientists used this information and the charge-to-mass ratio to determine the mass of an electron.

25. Millikan’s Experiment
Atomizer - +
Metal Plates
Oil
Microscope

26. Millikan’s Experiment
Atomizer
Oil droplets - +
Oil
Microscope

27. Millikan’s Experiment
X-rays
X-rays give some drops a charge by knocking off electrons

28. Millikan’s Experiment+ -

29. Millikan’s Experiment- - + +
They put an electric charge on the plates

30. Millikan’s Experiment- - + +
Some drops would hover

31. Millikan’s Experiment- - - - - - - +
Some drops would hover + + + + + + +

32. Millikan’s Experiment- - + +
From the mass of the drop and the charge on
the plates, he calculated the charge on an electron

33. Inferences About Atomic Structure
Atoms are electrically neutral, therefore, they must contain a positive charge to balance the negative electrons.
Electrons have so much less mass than atoms, atoms must contain other particles that account for most of the atoms mass.

34. Lord Ernest Rutherford
English Physicist 1911
Worked with
Hans Geiger and Ernest Marsden
Student of J.J. Thomson
Believed the plum pudding model of the atom was correct
Set out to prove it!

35. Rutherford’s Experiment
Alpha Particle Scattering Experiment
Bombarded a thin piece of gold foil with fast-moving alpha particles, positively charged particle 4 times the mass of a hydrogen atom.
Expected the particles to go through the gold atoms with only slight deflection.

36. Rutherford’s Experiment
Most did what was expected but 1 in 8000 was deflected back toward the source.

37. Flourescent
Screen
Lead block
Uranium
Gold Foil

38. What he expected

39. Because, he thought the mass was evenly distributed in the atom

41. What he got

42. Rutherford’s Conclusion
When the alpha particles hit a florescent screen, it glows.
The deflected particles experienced a powerful force within the atom.
He figured this force occupied a very small space because most of the particles went through.

43. Rutherford’s Conclusion
Atom is mostly empty space.
Small densely packed bundle of matter with a positive electric charge.
Alpha particles are deflected by it if they get close enough. +

44. +

45. What We Know So Far
Except for the simplest hydrogen atom, all nuclei are composed of two kinds of particles
positive Protons
neutral Neutrons
Atoms are electrically neutral, they contain the same number of protons and electrons.

46. What We Know So Far
The nuclei of atoms of different elements differ in their number of protons and therefore in the amount of positive charge.
The number of protons determines the atoms identity.

47. What We Don’t Know Yet
Where are the electrons?

48. Forces In The Nucleus Like charges repel each other
Except in the nucleus
When like charges are extremely close together there is a strong attraction between them.
More than 100 protons can exist together to help form a nucleus.
These short range forces hold the nuclear particles together, they are called nuclear forces.

49. Subatomic Particles
Proton - positively charged pieces 1836 times heavier than the electron.
Neutron - no charge but the slightly larger than the mass of a proton.
Electron – negatively charged pieces ~1/1837 times the mass of the hydrogen atom.
Where are the pieces?

50. Subatomic particles Relative mass Actual mass (g) Name Symbol Charge
Electron e- -1
1/1837
9.109 x 10-28
Proton p+ +1 1
1.673 x 10-24
Neutron n0 1
1.675 x 10-24

51. The Structure of the Atom
An atom is the smallest particle of an element that retains the chemical properties of that element.
All atoms are composed of two regions:
Nucleus: very small dense region located in the center of the atom.
Made up of at least one positive particle, proton, and usually one or more neutral particles called neutrons.

52. The Structure of the Atom
Surrounding the nucleus is a large region occupied by negatively charged particles called electrons.
This region is called the electron cloud.
Protons, Neutrons, and Electrons are referred to as subatomic particles.

53. Structure of the Atom There are two regions. The nucleus.
With protons and neutrons.
Positive charge.
Almost all the mass.
Electron cloud- most of the volume of an atom.
The region where the electron can be found.

54. Size of an atom Atoms are small. Measured in picometers, 10-12 meters.
Hydrogen atom, 32 pm radius.
Nucleus tiny compared to atom.
IF the atom was the size of a stadium, the nucleus would be the size of a marble.
Radius of the nucleus is near 10-15m.
Density near 1014 g/cm3.

55. Counting the Pieces Atomic Number = number of protons
# of protons determines kind of atom.
the same as the number of electrons in the neutral atom.
Mass Number = the number of protons + neutrons.
All the things with mass.
NOT on the periodic table

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

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

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

59. Naming Isotopes Put the mass number after the name of the element.
carbon- 12
carbon -14
uranium-235

60. Na Symbols 24 11 Find the number of protons number of neutrons
number of electrons
Atomic number
Mass Number
Name 24 Na 11

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

62. 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
Name

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

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

65. 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.

66. Measuring Atomic Mass Unit is the Atomic Mass Unit (amu)
One twelfth the mass of a carbon-12 atom.
6 p+ and 6 n0
Each isotope has its own atomic mass
we get the average using percent abundance.

67. Calculating averages
You have five rocks, four with a mass of 50 g, and one with a mass of 60 g. What is the average mass of the rocks?
Total mass = x x 60 = 260 g
Average mass = 4 x x 60 = 260 g
Average mass = 4 x x 60 = 260 g

68. Calculating averages Average mass = 4 x 50 + 1 x 60 = 260 g 5 5 5
80% of the rocks were 50 grams
20% of the rocks were 60 grams
Average = % as decimal x mass % as decimal x mass % as decimal x mass +

69. 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.

70. 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

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