1. Basic Atomic Theory Unit 2 (Chapter 3)
Electrons, protons, & neutrons, relative atomic mass (& isotopes), the Mole and molar mass

2. WKRP in Cincinnati Atom Lesson: https://youtu.be/WjcmU9Ijy1o
(2:13-6:02)

3. Relative electric charge
Particle
Symbol
Relative electric charge
Mass number
Relative mass (amu)
Actual mass (kg)
Electron
e-, 0-1e -1
9.109 x 10-31
Proton
p+, 11H +1 1
1.673 x 10-27
Neutron
n0, 10n
1.675 x 10-27
*1 amu (atomic mass unit) = x kg
Similar chart on page 76 in your textbook

4. Atomic number (Z) = the number of protons in the nucleus of an atom of that element.
Z is from the German Zahl or “number.”
Mass number (i.e. atomic mass) [A] = total number of protons and neutrons in the nucleus of an isotope.
A from German Atomgewichte or “atomic weight.”

5. What makes the atomic mass different?
Isotopes = atoms of the same element that have different masses (i.e. different numbers of neutrons). They may be naturally occurring or man-made in a laboratory. 
Nuclide is a general term for any isotope of an element

6. Ion: an atom or group of bonded atoms that has a positive or negative charge as a result of having either more or less electrons than the neutral atom from which it originated. Any process that results in the formation of an ion is called ionization.
Oxidation number/state/valence: a number assigned to an atom in a molecular compound or molecular ion that indicates the general distribution of electrons among the bonded atoms; i.e. the atom or ion’s charge.

7. Atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom. The atomic mass of any nuclide is determined by comparing it with the mass of the carbon-12 atom.
Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element.
This is what we see on the periodic table.

8. Law of conservation of mass: mass is neither created nor destroyed during ordinary chemical reactions or physical changes.
Law of definite proportions: a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.
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.

9. My.hrw.com video clips

10. Calculating average atomic mass
Multiplying the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results:
Example: Copper, Cu
( x amu) + ( x amu) = amu
Relative abundance is determined using a mass spectrometer

11. https://www. khanacademy
Figure 3: A diagram of a mass spectrometer. A sample is injected into the machine, vaporized by a heater, and then ionized by a stream of high-energy electrons. The resulting Ions are accelerated through parallel electric plates and then deflected in a magnetic field before they reach a detector. Imagefrom Openstax, CC BY 4.0.

12. Accelerator mass spectrometer at Lawrence Livermore National Laboratory

13. ueip.org/mass-spectrometer-isotope-analysis-u...

15. Figure 4: The simulated mass spectrum for a sample of elemental zirconium.Image from Openstax, CC BY 4.0.

16. Let’s Take a Trip Through Time!
Some slides taken from Ron Brandt’s PPT

17. Democritus 460 – 370 B.C.
There are various basic elements from which all matter is made
Everything is composed of small atoms moving in a void
Some atoms are round, pointy, oily, have hooks, etc. to account for their properties
Ideas rejected by leading philosophers because void = no existence

18. John Dalton 1766-1844 Introduced his ideas in 1808
Each element is composed of extremely small particles called atoms
All the atoms of a given element are identical, but they differ from those of any other element
Atoms are neither created nor destroyed in any chemical reaction
A given compound always has the same relative numbers and kinds of atoms
John Dalton

19. Elements can have atoms with different masses (isotopes)
Which parts of Dalton’s theory are no longer “true” today?
Elements can have atoms with different masses (isotopes)
Atoms are divisible into even smaller particles (like what?)

20. Discovered electron 1897 – Cathode Ray Experiment
Plum Pudding model 1904
Electrons in a soup of positive charges
Discovered isotopes 1913
J.J. Thomson

21. Discovery of the electron
Cathode ray tube
1. An object placed between the cathode and the opposite end of the tube cast a shadow on the glass.
2. A paddle wheel placed on rails between the electrodes rolled along the rails from the cathode toward the anode.
3. Cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current, which was known to have negative charge.
4. The rays were deflected away from a negatively charged object.

22. These observations led to the hypothesis that the particles that compose cathode rays are negatively charged. This was supported by experiments undertaken by J.J. Thompson in In one experiment he found that the ratio of the charge of cathode-ray particles to their mass was always the same, regardless of the cathode or the type of gas used.
He concluded that cathode rays are composed of identical negatively charged particles that were later named electrons.
Show videos in my.hrw.com for Thompson and Millikan

23. Mass of an electron: 9.109 x 10-31 kg
Millikan’s
Oil Drop
Experiment (1909)
He determined
the mass of an
electron.
Mass of an electron: x kg

24. Discovery of the atomic nucleus
Ernest Rutherford,
Hans Geiger and
Ernest Marsden (1911)
They bombarded a very thin
piece of gold foil with alpha particles
(about four times the mass of a hydrogen atom). They found that
the majority of the particles went straight through the foil and a
very few were deflected.
Ernest Rutherford ( )

25. Alpha Particle Experiment

26. Therefore, each atom in the gold foil had a very small, dense, positively charged nucleus surrounded by electrons.

27. James Chadwick 1891-1974 Worked with Rutherford
Interpreted work of the Curies
Discovered Neutron 1932
Nobel Prize in Physics 1935

28. Discovery of the neutron library.thinkquest.org/27954/neutron.html
In 1932, James Chadwick bombarded beryllium (Be) with alpha particles. He allowed the radiation emitted by beryllium to incident on a piece of paraffin wax.
It was found that protons were shot out from the paraffin wax. People began to look for what was in the "beryllium radiations".

29. Some people suggested that the radiations may be gamma radiation
Some people suggested that the radiations may be gamma radiation. However, Chadwick found that the radiation could not be gamma radiation since energy and momentum were not conserved in its production. He showed that all the observations could be explained if the radiation consisted of neutral particles of mass approximately equal to that of proton. This neutral particle was named neutron.
Equation of the nuclear reaction

30. This new tool in atomic disintegration did not have to overcome any electric barrier and was capable of penetrating and splitting the Nucleus of even the heaviest elements. Chadwick thus prepared the way towards the fission of uranium 235 and towards the creation of the atomic bomb. For this discovery he was awarded the Hughes Medal of the Royal Society in 1932, and subsequently the Nobel Prize for Physics in 1935.

31. Niels Bohr 1885-1962 Planetary Model 1913
Nucleus surrounded by orbiting electrons at different energy levels
Electrons have definite orbits
Utilized Planck’s Quantum Energy theory
Worked on the Manhattan Project (U.S. atomic bomb)
Niels Bohr
Actually discuss him more in the next unit

32. Discovery of the neutron led to the beginning of
current theories of nuclear structure. Immediately, the
neutron-proton model ( the Rutherford-Bohr model) of
the nucleus was adopted:
The nucleus is made up of protons and neutrons. These are bound together by a strong nuclear force.
Electrons and protons carry equal but opposite charges. In a neutral atom, the number of electrons is the same as the number of protons.
Electrons orbit the nucleus at certain fixed levels called shells.

33. Bohr’s Model
More in next unit

34. Nuclear forces Usually like charges repel, but protons and
neutrons can be very
close together. These
short-range proton
neutron, proton
proton, and neutron
neutron forces hold
the nuclear
particles together
And are called
nuclear forces.
Nuclear forces

36. Ernst Schrödinger 1887-1961 Werner Heisenberg 1901-1976
Quantum Mechanical Model 1926
Electrons are in probability zones called “orbitals”, not orbits and the location cannot be pinpointed
Electrons are particles and waves at the same time
Developed quantum numbers based on theories of Einstein and Planck
More in next unit

37. The Mole
a mole is the SI base unit used to measure the amount of a substance whose number of particles is the same as the number of atoms of carbon in exactly 12 g of carbon It is equal to x 1023.
Simply represents a number. Just as the term dozen refers to the
number twelve, the mole
represents the number
6.02 x 1023

38. The history of the MOLE
Amadeo Avogadro (1811) proposed that equal volumes of different gases at the same temperature contain equal numbers of molecules.
Show my.hrw.com videos

39. Stanislao Cannizzaro
About fifty years later, Stanislao Cannizzaro used Avogadro's hypothesis to develop a set of atomic weights for the known elements by comparing the masses of equal volumes of gas.

40. Johann Josef Loschmidt (Austrian high school teacher)
1865, calculated the size of a molecule of air, and thus developed an estimate for the number of molecules in a given volume of air.

41. 6.02 x 1023 Molar Mass – the mass of one mole of a pure substance
One mole of any substance will contain one mole of particles.
6.02 x 1023
Elements = atoms
Compounds = molecules 

42. Calculating Molar Mass
multiply atomic mass of each element by number of atoms of that element in the formula (shown by the subscript)
find the sum of all the atomic masses -- this is formula mass (unit is a.m.u.)
express formula mass in grams (unit is g/mol). This is the Molar Mass.