1. Unit 4
THE ATOM

2. EVOLUTION OF The Atomic Theory

3. Main idea
The development of the atomic theory is a perfect explain of how the scientific method works.
This example illustrates the continuous changes many scientific theories undergo over time.
This example uses scientific models.
This example uses numerous tests and experiments.

4. Democritus – 460 B.C. Greek Philosopher
Only talk and thoughts, no evidence or experiments.
At some point you can’t divide matter anymore.
Named these smallest pieces an atom, meaning “indivisible.”
Atomic model: smallest pieces

5. Fast forward to 1800’s
Greek society collapses and harder times force science to takes a backseat to basic survival.
Atomic theory lays dormant for about 2000 years.
The industrial revolution (and technology) during the 1800’s allows science to prosper once again.
Getting quantitative….
Law of conservation of mass
Law of definite proportions
Law of multiple proportions

6. Law of conservation of mass - 1782
Mass is neither destroyed nor created during ordinary chemical and physical reactions.
mass of reactants = mass of products
Carbon + Oxygen  Carbon Dioxide
C O2  CO2
12g g g

7. Law of DEFINITE PROPORTIONS- 1799
A given compound always contains the same fixed ratio of elements.
The recipes for chemical compounds never change.
Water is always
11.20% Hydrogen
88.80% Oxygen

8. Dalton’s atomic theory - 1803
English School Teacher
Developed 1st atomic theory
Based on experiments
Atomic Model: hard sphere

9. Dalton’s atomic theory - 1808
All matter is composed of extremely small particles called atoms.
Atoms of a given element are identical in size, mass and other properties.
Atoms cannot be subdivided, created or destroyed.
Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
In chemical reactions, atoms can be combined, separated and rearranged.

10. Conclusion
All matter is composed of extremely small particles called atoms.
Atoms of a given element are identical in size, mass and other properties.
Atoms cannot be subdivided, created or destroyed.
Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
In chemical reactions, atoms can be combined, separated and rearranged.

11. Discovery of the electrons – 1897
English Physicist
Thomson
Preformed experiments with cathode-ray tubes.

12. (a stream of negative particles)
Cathode-Ray Tube
Voltage Source
Cathode Ray
(a stream of negative particles)
CATHODE
ANODE

13. An object placed between the cathode ray and the opposite end
of the tube cast a shadow on
the glass.
This showed movement and
direction.

14. A paddle wheel placed on rails between
the electrodes rolled along the rails from
the cathode towards the anode.
This showed cathode ray had a mass or force.

15. Cathode rays could be manipulated by magnets
Cathode rays could be manipulated by magnets. The ray was attracted to the positive side of a magnet and repelled by the negative side of the magnet.

16. Conclusions
1. atoms have negatively charged particle called electron 2. atoms are divisible 3. hypothesized that atoms must also contain a positively charge particle that balance the atoms. Atomic Model: Plum Pudding

17. Discovery of the nucleus
Ernest Rutherford (1911)
Gold Foil Experiments
Atomic Model:
Nuclear Model

18. Gold foil Experiment

19. Most of the alpha particles
pass straight through.
Some hit something that is dense and positive.

20. Conclusion
Rutherford concluded that the rebounded alpha particles must have experienced some powerful force within the atom. And he figured that the source of this force must occupy a very small amount of space because only a few of the total number of alpha particles had been affected by it.
Atoms have a very densely packed bundle of positive matter called the nucleus. Later, it will be found to be the proton.

21. Planetary model Niels Bohr (1912) Spectroscopy (Light) Experiments
Atomic Model: Planetary Model

22. Planetary model
Electrons move around the nucleus in fixed energy levels.
These energy levels are like rungs on a ladder where energy is needed to move up “the ladder”
Think of the electrons as a moving fan or bee!

23. Conclusion
Suggested that the electrons orbiting the nucleus of atoms can only have certain discrete energies.
Electrons lose/gain energy in order to move from one energy level to another.
The emission of light (photon) occurs when an electron moves from a higher to a lower energy orbit (emission spectra).

24. According to Bohr, electrons exist as sub-atomic particles found in neat orbits around the nucleus.
Many scientists soon found that this idea of the electron’s location was an over simplification.
Electrons can have properties of both particles and waves which make them much more complex to understand.

25. Structure of the atom

26. The Atom
The smallest unit of an element that maintains the properties of that element.
So what does all of this mean for the shape, structure and contents of the atom…

27. Properties OF subatomic particles
Electron
Neutron
Proton
Charge -1
Negative
Neutral +1
Positive
Mass
1/1836 amu
(Small)
1 amu
(Large)
Location
Outside Nucleus
(Electron Cloud)
Nucleus
AMU = Atomic Mass Unit

28. The Atom
Electrons are located outside the nucleus in the electron cloud.
Most of the volume of an atom is occupied by the electron cloud.
The protons and neutrons are located inside in the center of the atom in the nucleus.
The majority of the mass of an atom is found in the nucleus (neutrons & protons).

29. The Atom Typically, particles with the same charge repel each other.
Nuclear forces - short range forces that help hold the atom together.
Proton – proton
Neutron – neutron
Neutron – proton

30. Atomic number & atomic mass

31. Atomic Number Atomic Number = Number Of Proton
Different For Each Element

32. Atomic NUMBER Number of protons = Number of electrons
Result is a neutral atom

33. Isotopes
Atoms of a given element that contain different numbers of neutrons.
Remember: Neutrons have a sizable mass.
Therefore: Atoms of the same element that have different masses.
Same number of protons (+) and electrons (-).

35. Isotope Notation Hydrogen – 3 Hydrogen – 2 Hydrogen – 1
Number after element’s name represents isotopes atomic mass

37. Average mass of all the isotopes of an element.
Atomic mass
Average mass of all the isotopes of an element.

38. Atomic Mass
Atomic Mass = number of protons + number of neutrons number of neutrons = atomic mass - atomic number

39. The Mole SI unit for amount of substance.
the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon - 12
Counting unit like a dozen or π.

40. Avogadro’s Number 6.0221367 X 1023 rounded to 6.02 X 1023
The number of particles in exactly one mole of a pure substance.

41. Molar mass (Formula Weight)
The mass of one mole of a pure substance (g/mol)
Equal to the atomic mass of an element or the elements in a compound.
Examples:
Hydrogen – g/mol or 1g/mol
Oxygen – g/mol or 16g/mol

42. Atom/Mass/Mole Conversions
Sample Problem
Finding mass from moles
Finding moles from mass
Finding atoms from moles
Finding moles from atoms