1. Do Atoms exist?

2. Law of Conservation of Mass
The total mass of substances present at the end of a chemical process is the same as the total mass of substances present before the process took place.
What are the implications of this Law?

3. Combust 11.5 g ethanol uses 24.0 g of oxygen = 35.5 g reactants
Collect 22.0 g CO2 and 13.5 g H2O = 35.5 g of products
3.6

4. Law of Constant Composition Joseph Proust (1754–1826)
The elemental composition as determined by mass measurements of a pure substance does not vary.
Usually reported as a mass percent of each element.

5. Percent composition of an element in a compound =
mass of element
molar mass of compound
x 100%
When we analyze pure ethanol, the percent by mass is the following always:
%C = 52.14%
%H = 13.13%
%O = 34.73%
52.14% % % = 100.0%
3.5

6. Multiple Proportions
Some atoms (mostly nonmetals) can combine in different mass ratios with each other. In the event this happens, the mass ratios are whole numbers of each other.
Link to Movie

7. Atomic Theory of Matter
In the early 19th century, John Dalton considered what these known Laws could mean.
His thoughts are now known as Dalton’s atomic theory.
Figure 2.1 John Dalton ( )

8. Dalton’s Postulates
Each element is composed of extremely small particles called atoms.

9. Dalton’s Postulates
All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

10. Dalton’s Postulates
Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

11. Dalton’s Postulates
Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.
So what happened next?

12. The Electron
Figure 2.4
Thompson measured the charge/mass ratio of the electron to be 1.76  108 coulombs/g.

13. Millikan Oil Drop Experiment
Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other.
Figure 2.5

14. Millikan Oil Drop Experiment
Robert Millikan (University of Chicago) determined the charge on the electron in 1909.
Figure 2.5
Link to movie

15. Radioactivity:
The spontaneous emission of particles and/or high frequency waves by an atom.
First observed by Henri Becquerel.
Also studied by Marie and Pierre Curie.

16. Radioactivity
Three types of radiation were discovered by Ernest Rutherford:
 particles (mass about 4 times that of hydrogen)
 particles (mass +- 1/2000 that of hydrogen)
 rays (a high frequency wave)
Link to video
Figure 2.8

17. The Atom, circa 1900: “Plum pudding” model, put forward by Thompson.
Positive sphere of matter with negative electrons imbedded in it.
Figure 2.9

18. Discovery of the Nucleus
Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.
Figure 2.10
Link to Video

19. The Nuclear Atom
Since some particles were deflected at large angles, Thompson’s model could not be correct.
Figure 2.11

20. The Nuclear Atom
Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom.
Most of the volume of the atom is empty space.
Figure 2.12

21. Other Subatomic Particles
Protons were discovered by Rutherford in 1919.
Neutrons were discovered by James Chadwick in 1932.

22. Subatomic Particles
Protons and electrons are the only particles that have a charge.
Protons and neutrons are very close in mass.
The mass of an electron is 1/1837 of a proton or neutron.
Table 2.1

23. X H H (D) H (T) U Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei
Mass Number X A Z
Element Symbol
Atomic Number H 1
H (D) 2
H (T) 3 U
235 92
238
2.3

24. Isotopes of Hydrogen
Link to activity

25. Atomic mass is the mass of an atom in atomic mass units (amu)
Micro World
atoms & molecules
Macro World
grams
Atomic mass is the mass of an atom in atomic mass units (amu)
By definition:
1 atom 12C “weighs” 12 amu
On this scale
1H = amu
16O = amu
3.1

26. Atomic Mass
Atomic and molecular masses can be measured with great precision with a mass spectrometer.
1 amu = 1/12 carbon-12
Link to Movie
Figure 2.13

27. Heavy
Light
3.4

28. Average Mass
Because in the real world we use large amounts of atoms (the elements), we use average masses in calculations.
Average mass is calculated from the isotopes of an element weighted by their relative abundances.

29. Average atomic mass of lithium:
Natural lithium is:
7.42% 6Li (6.015 amu)
92.58% 7Li (7.016 amu)
Average atomic mass of lithium:
7.42 x x 7.016
100
= amu
3.1

30. Average atomic mass (6.941)

31. Do You Understand Isotopes?
How many protons, neutrons, and electrons are in C 14 6 ?
6 protons, 8 (14 - 6) neutrons, 6 electrons
How many protons, neutrons, and electrons are in C 11 6 ?
6 protons, 5 (11 - 6) neutrons, 6 electrons
2.3