1. History of the Atom; Modern Atomic Theory, Subatomic Particles
Atomic Theory CP
History of the Atom; Modern Atomic Theory, Subatomic Particles

2. Reasons to look at history
Shows that present theory is the work of many scientists over many years
To become aware of the logic used in obtaining evidence to support the theory
To gain a better understanding of the properties of atoms

3. Greek Philosophers Democritus (460 – 371 B.C.)
Stated that the world was made of tiny particles that are mostly empty space and that they are the building blocks of all matter
Tiny particles could not be divided
Termed them atomos, or atoms
Means “uncut” or “indivisible”
Aristotle (384 – 322 B.C.)
Proposed that matter was continuous (i.e. came from only four elements – earth, air, fire, and water)

4. Dalton’s Atomic Theory
John Dalton (1766 – 1844)
Work marks the beginning of the development of the modern atomic theory
Revived Democritus’ ideas and backed them up with scientific experimentation and data!
Dalton’s Atomic Theory:
Matter is composed of extremely small particles called atoms
Atoms are indivisible and indestructible
Atoms of a given element are identical in size, mass, and chemical properties
Atoms of a specific element are different from those of another element
Different atoms combine in simple whole-number ratios to form compounds
In a chemical reaction, atoms are separated, combined, and rearranged

5. Dalton’s Atomic Theory
Dalton’s theory successfully explained:
The law of conservation of mass (proposed by Lavoisier)
States that mass is conserved in any process, such as a chemical reaction.
Conservation is the result of the separation, combination, or rearrangement of atoms
The law of multiple proportions (proposed by Proust)
Certain elements can combine to form two or more different chemical compounds
Experimental evidence led to general acceptance of his atomic theory!
Not all of the theory was accurate.
It has had to be revised with new information

6. Modern Atomic Theory
All matter is made up of a very small particles called atoms
Atoms of the same element have the same chemical properties, atoms of different elements have different properties
Any natural sample of an element has an average mass characteristic of that element
Compounds are formed when atoms of two or more elements unite
Atoms are not subdivided in physical or chemical reactions

7. Protons, Neutrons, and Electrons
Defining the atom

8. The Atom
Atom:
Smallest particle of an element that retains the properties of the element
Size:
Consider this:
In 2006 the population of Earth was about 6.5 x 109 people
A solid copper penny contains 2.9 x 1022 atoms
This is 5 trillion times larger than the worlds population!!!!!

9. Electrons J.J. Thomson 1879 Cathode Ray Tube
Provided the first evidence that atoms are made of even smaller particles
Cathode-Ray Tube:
Vacuum tube, all gases removed
Metal electrodes at each end
Electricity runs from the negative cathode to the positive anode
Using a magnet, the negative charged cathode ray bent towards a positively charged plate

10. Electrons
Observations of the cathode-ray behavior led to the hypothesis that cathode rays consist of identical negatively charged particles
Thomson proposed the “Plum Pudding” model of the atom

11. Thomson’s Model Found the electron
1 unit of negative charge
Mass 1/2000 of hydrogen atom
Later refined by Millikan to 1/1840
Concluded that there must be a positive charge since atom was neutral
Atom was like plum pudding
A bunch of positive stuff, with electrons able to be removed

12. Nucleus Ernest Rutherford – 1908 Discovered the nucleus
The core of the atom and contains protons and neutrons
Rutherford’s nuclear model proposed that an atom is mostly empty space
This is based on studies of the bombardment of thin metal foils with fast-moving positively charged particles

13. Rutherford’s Experiment

14. Rutherford’s Experiment
His experiment demonstrated that:
An atom is mostly empty space
There is a small dense, positive piece at the center
Alpha particles are deflected if they got close enough to the positive center

15. Protons and Neutrons Protons: Neutrons:
Rutherford also concluded that the nucleus also contained positively charged particles call protons.
Subatomic particle with a positive charge equal in magnitude to the negative charge of an electron
The number of protons determines the identity of an atom!!
Neutrons:
Discovered by James Chadwick (Rutherford’s coworker)
Showed the nucleus also contained a neutral subatomic particle, the neutron
Electrically neutral subatomic particle
The number of neutrons can vary within an element  Isotopes!

16. Rutherford Activity

17. Structure of the Atom There are two regions The nucleus:
Protons and Neutrons
Positive charge
Almost all of the mass
Electron Cloud:
Most of the volume of an atom
Region were electrons can be found

18. Comparing Subatomic Particles
Properties of Subatomic Particles
Particle
Symbol
Relative Charge
Relative Mass
(Mass of Proton = 1)
Actual Mass (g)
Electron e- 1-
1/1840
9.11 x 10-28
Proton p+ 1+ 1
1.67 x 10-24
Neutron n0

19. Atomic Number
Atoms of each element contain a unique positive charge in their nuclei
Therefore, the number of protons in an atom identifies that particular atom!
The number of protons in an atom is referred to as the atomic number
Because all atoms are neutral:
The number of protons = the number of electrons
If you know the atomic number of an element you now know both the number of protons and the number of electrons!
Atomic Number

20. Mass Number
The total number of protons and neutrons in the nucleus of an atom
Mass # = p + n
Usually found at the bottom of the atomic symbol
How to find the number of neutrons:
Mass number – atomic number = # neutrons
Example:
Nitrogen:
Mass number: 14
Atomic number: 7
Number of Neutrons: 7

21. Fill in the Gaps
Atomic #
Mass # P N E 3 7 39 19 6 10

22. Mass Number Continued
Although a given type of atom will usually contain a certain number of neutrons in the nucleus, a small percentage will not
Ex: most hydrogen atoms contain no neutrons
A small percentage contain one neutron and a smaller percentage two neutrons
What do we call atoms with a different number of neutrons?

23. Isotopes In nature, most elements are found as a mixture of isotopes.
Atoms that have the same number of protons but different numbers of neutrons
This means that the number of neutrons in a atom can change!
Protons MUST stay the same
Different isotopes of the same element have different mass numbers (i.e. different neutrons)
A given element has at least two isotopes, and they all have different mass numbers
Example:
Gadolinium (Gd) has 6 stable isotopes

24. Isotopes
Isotopes are chemically alike because they have identical numbers of protons and electrons
It is useful to compare the relative masses of atoms to a standard reference isotope
Carbon-12 is the standard reference isotope
Carbon-12 has a mass of exactly 12 atomic mass units (amu)

25. Atomic Mass
A weighted average mass of the atoms in a naturally occurring sample of the element
Mass of an atom in atomic mass units (amu)
Equal to 1/12th of the mass of carbon
Weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature
This is the number which appears on the periodic table

26. Atomic Mass Number Example:
99% of all carbon atoms are the isotope containing 6 neutrons, the remaining 1% is the heavier isotope containing 7 neutrons, which raises the aver mass of carbon from to

27. Calculations Carbon has two major, stable isotopes.
Carbon-12 (98.93% abundance)
Carbon-13 (1.07% abundance)
Carbon-14 is in trace amounts and is radioactive
Calculate the average atomic mass of carbon
Carbon 12: ( x 12) =
Carbon 13: ( x 13) =
AAM:

28. Calculations
Rubidium has two common isotopes, 85Rb and 87Rb. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium?
Average atomic mass = (0.722 x 85) + (0.278 x 87)
Average atomic mass = (61.37) + (24.186)
Average atomic mass = amu