1. Atoms and their structure
Unit 2
Atoms and their structure

2. How we started to think about atoms
Original idea came from Ancient Greece (400 B.C.)
Democritus and Leucippus were Greek philosophers

3. How we started to think about atoms
Democritus looked at the beach
Made of sand
If you cut sand particles, you get smaller sand particles

4. How we started to think about atoms
There must be a smallest possible piece
Called those pieces “Atomos” – not able to be cut

5. 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

6. Who Was Right?
None of the philosophers experimented to determine who was right
Greeks settled disagreements by argument
Aristotle was a better debater - He won
His ideas carried through middle ages
Later on, alchemists tried to change lead to gold (they did not understand atoms)

7. Who’s Next? England in the late 1700’s - John Dalton
Teacher who summarized results of his experiments and those of others
Elements are substances that can’t be broken down
In Dalton’s Atomic Theory, he combined the idea of elements with that of atoms

8. 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.

9. 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
A limited amount of other gases are put in and an electric current is applied to the tube

10. Thomson’s Cathode Ray Tube
Voltage source - +
Metal Disks

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

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

13. Thomson & his atomic model
Discovered the electron
He did not know where positive charges were
Said the atom was like plum pudding
A bunch of positive stuff, with the electrons able to be removed

14. Rutherford’s Experiment
Ernest Rutherford - English physicist, 1910
Believed the plum pudding model of the atom was correct
Wanted to see how big atoms are
Used radioactivity
Alpha particles - positively charged pieces given off by uranium
Shot them at gold foil which can be made a few atoms thick

15. Rutherford’s experiment
When the alpha particles hit a fluorescent screen, it glows
Here’s what it looked like
Fluorescent
Screen
Lead block
Uranium
Gold Foil

16. Fluorescent
Screen
Lead block
Uranium
Gold Foil

17. He expected that…
The alpha particles would pass through without changing direction very much
Because…
The positive charges were spread out evenly - alone they would not be enough to stop the alpha particles

18. What he expected

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

20. What he got

21. How he explained it Atom is mostly empty space
There is a small dense, positive piece at the center
Alpha particles are deflected by it, if they get close enough +

22. +

23. Modern View The atom is mostly empty space Two regions
Nucleus - protons and neutrons
Electron cloud - region where you might find an electron

24. Density and the Atom
Since most of the particles went straight through the gold foil, the atom was mostly empty space
Because the alpha particles turned so much, the positive particles must have been heavy
Small volume and big mass = big density.
This small dense positive area is the nucleus

25. Subatomic particles
Electron – located outside of nucleus, has negative charge
Proton - positively charged particles inside of nucleus that are many times heavier than the electron
Neutron - no charge but about the same mass as a proton, located inside of nucleus

26. Subatomic particles Relative mass (amu) Particle Name Actual mass (g)
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

27. 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

28. Size of an atom Atoms are small Measured in picometers, 10-12 meters
Hydrogen atom, 32 pm radius
Nucleus is very 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 m
Density near 1014 g/cm3

29. 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
Includes all the particles with mass
NOT found on the periodic table

30. What about when Electrons ≠ Protons
Electrons may be gained or lost
Gaining electrons gives a negatively charged ion called an anion
Losing electrons gives a positively charged ion called a cation

31. What about when Electrons ≠ Protons
Practice with ions…
Magnesium makes ions with a 2+ charge. Are electrons lost or gained? How many electrons are moved?
Fluorine makes ions with a 1- charge. Are electrons lost or gained? How many electrons are moved?
An ion has 13 p+ and 10 e-. Give the symbol and charge for the ion.
An ion has 34 p+ and 36 e-. Give the symbol and charge for the ion.

32. Isotopes Dalton was wrong
Atoms of the same element can have different numbers of neutrons
different mass numbers (will have the same atomic number
called isotopes

33. X Symbols for Isotopes Mass number Atomic number Nuclear Notation
Contains the symbol of the element, the mass number and the atomic number X
Mass
number
Atomic
number

34. Symbols for Isotopes carbon- 12 carbon -14 uranium-235 Hyphen Notation
Contains the symbol (or name) of the element and the mass number.
carbon- 12
carbon -14
uranium-235

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

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

37. Symbols for Isotopes
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
Symbol – Nuclear & Hyphen notation
Name

38. Symbols for Isotopes
if an element has 91 protons and 140 neutrons what is the
Atomic number
Mass number
number of electrons
Symbol – Nuclear & Hyphen notation
Name

39. Symbols for Isotopes
if an element has 78 electrons and 117 neutrons what is the
Atomic number
Mass number
number of electrons
Symbol – Nuclear & Hyphen notation
Name

40. 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

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

42. 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 of an element has its own atomic mass
we get the average atomic mass of an element using weighted averages (need mass & percent abundance)

43. 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 = (4 x 50) + (1 x 60) = 260 g
Average mass = (4 x 50) + (1 x 60) = 260 g

44. Calculating averages (% as decimal x mass) + (% as decimal x mass) + …
If 80% of the rocks were 50 grams and 20% of the rocks were 60 grams what is the weighted average mass of the rocks?
Weighted Average =
(% as decimal x mass) +
(% as decimal x mass) + …
(0.8 x 50 g) + (0.2 x 60 g) = 52 g

45. Homework – due 10/18 Pg 76 (#3) Pg 87 (#2-3)
Qs with * need to have answers in the form of complete sentences. Please show your math work as well.