1. SCH 3UY - Atomic Theory Review
The Greeks:
Democritus determined in 300 BC that matter cut into smaller and smaller pieces would eventually reach what they called the atom – literally meaning indivisible (atomos).
This idea was reintroduced over two thousand years later by an English chemist/schoolteacher named John Dalton in 1805.

2. John Dalton proposed:
All matter is composed of tiny indivisible particles called atoms.
Atoms of one element cannot be converted into atoms of another element.
Atoms of a given element are identical to each other.
Compounds result from the combination of elements in specific whole number ratios.

3. 1) Law of Mass Conservation
Total mass of substances does not change during a chemical reaction.
2) Law of Constant Composition
A particular compound is composed of the same elements in similar proportions or ratios.
3) Law of Multiple Proportions
If element A and B react to form 2 compounds, the masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers.

4. Cathode Ray Tube

5. Cathode rays are emitted from the cathode when electricity is passed through an evacuated tube.
They are emitted in a straight line, perpendicular to the cathode surface.
They cause glass and other materials to fluoresce.
They are deflected by a magnet in the direction expected for negatively charged particles. (attracted to positively charged magnets, repelled by negatively charged magnets.)
Their properties do not depend on the composition of the cathode. For example, the cathode rays from an aluminum cathode are the same as those from a silver cathode.

6. Observations were made by J. J
Observations were made by J. J. Thomson that required a revision of Dalton’s atomic theory.
Thomson found that cathode rays emitted from a charged plate were composed of negatively charged particles that had a mass to charge ratio (m/z) much less than any known atom.
It was proposed that these “rays” were composed of subatomic particles.
We now call these particles electrons and often give them the symbol e- or e.
The observation of subatomic particles was in violation of Dalton’s ideas. A new atomic theory was required.

7. Millikan determined that the mass of the electron was less than 1/1000th of the lightest known atom.
Millikan’s oil drop
Experiment.

8. George Stoney: names the cathode-ray particle the electron.
Robert Millikan: determines a value for the electron’s charge:
e = –1.602 × 10–19 C
Concluded that there exists a “light” carrier of negative charge that is indivisible and a fraction of an atom.
Charged droplet can move either up or down, depending on the charge on the plates.
Radiation ionizes a droplet of oil.
Magnitude of charge on the plates lets us calculate the charge on the droplet.

9. Thomson determined the mass-to-charge ratio; Millikan found the charge; we can now find the mass of an electron:
me = × 10–31 kg/electron
This is almost 2000 times less than the mass of a hydrogen atom (1.79 × 10–27 kg)
Some investigators thought that cathode rays (electrons) were negatively charged ions.
But the mass of an electron is shown to be much smaller than even a hydrogen atom, so an electron cannot be an ion.
Since electrons are the same regardless of the cathode material, these tiny particles must be a negative part of all matter.

10. Thomson’s “Plum Pudding” Model of the Atom.
In this model the light, negatively charged electrons are embedded in a “sphere of positive charge” to produce a neutral atom.

11. It was known that some elements emitted -particles these are positively charged particles.
They have twice the charge of an electron and four times the mass of a hydrogen atom.
Alpha particles consist of 2 positively charged
protons and 2 neutrons.

12. Ernest Rutherford was interested in studying how -particles interacted with matter.
Ernest Rutherford had members of his research group expose a sheet of gold foil to a beam of -particles.
If Plum-Pudding Model was correct the α‑particles should go through the foil unimpeded.
Predicted Behaviour

13. Actual Observations A very few “bounced back” to the source!
A few particles were deflected slightly by the foil.
Most of the alpha particles passed through the foil.
Alpha particles were “shot” into thin metal foil.

14. The only way Rutherford could satisfactorily explain the results of this experiment was as follows:
Most -particles pass through foil.
Atoms must be mostly empty space.
Some α-particles are deflected or bounce back.
The atom must contain a positively charged center containing nearly all the mass of the atom.
He called this the nucleus.

15. Rutherford proposed that the positive charge of the nucleus was a result of positively charged particles called protons.
He proposed that the small negatively charged electrons moved around the nucleus at a great distance from it like planets around a star.
However, Rutherford’s model could not account for all the mass of the nucleus.

16. Rutherford was able to predict the size of the atom and its nucleus:
The nucleus is like a marble in the middle of an Olympic stadium.
What takes up the rest of the space?
Has to have little mass
Has to be negatively charged

17. Modern atomic theory considers the atom to be comprised of a dense nucleus containing nearly all the mass of the atom. It is made of neutrons and protons surrounded at great distance by the small negatively charged electrons.

18. Atoms have no overall charge and so contain the same number of protons (+) as electrons (-).
The number of protons in the nucleus of an atom is called the atomic number. It is given the symbol Z.
All atoms of a given type or element have the same atomic
number (Z).
The periodic table arranges atoms in order of increasing atomic number.

19. Neutrons have no electrical charge
Neutrons have no electrical charge. For a given atom type the number of neutrons may vary.
The sum of the number of neutrons and protons in an atom is called the mass number. It is given the symbol A.
Atoms with the same number of protons but different numbers of neutrons in their nucleus are called isotopes.
number of neutrons (N)
= (mass number)-(atomic number) = A - Z

20. James Chadwick (Rutherford’s student) resolved the differences in nuclear masses in isotopes with the discovery of the neutron, a dense, uncharged particle that also resides in the nucleus.
The planetary or Rutherford model of the atom was now complete.

21. We can distinguish isotopes with the following notation:
Where:
E is the atom’s symbol
A is the mass number
Z is the atomic number
e.g. 612C, 613C, 614C, 11H, 12H and 13H
As all atoms with a given symbol have the same atomic number sometimes we skip writing Z.
e.g. 12C (carbon-12), 13C (carbon-13), 14C (carbon-14)

23. If we assign a mass of 12 atomic mass units (u or sometimes amu) to a carbon-12 atom then we can compute the relative atomic weight for any other atom.
One atomic mass unit is understood to be 1/12 of the mass of one atom of carbon -12.
all elements are compared by weight to this standard.
We can define the isotopic mass the mass of a specific atom in amu.
The periodic table gives atomic masses which are the weighted average overall all isotopes of an element in a natural sample.

24. The Russian chemist Dmitri Mendeleev organized the elements such that elements with similar chemical and physical properties were placed next to one another.
While no subatomic particles were known at the time, it is now recognized he was placing elements in order of atomic number.
Mendeleev noticed that after a series of elements the properties would repeat.

25. Mendeleev predicted that there should be an element with
properties similar to aluminum which he called eka-aluminum.
17 years after the development of Mendeleev’s table scientists discovered this missing element and named it gallium.
The experimental properties of gallium are very close to those Mendeleev predicted:

26. Mendeleev’s original periodic table. The elements labelled with
Mendeleev’s original periodic table. The elements labelled with ? Were predicted
but had not yet been isolated.

27. eka-aluminium
germanium

29. Homework
Nelson, Section 3.1 Questions #1-8, p. 166