1. The Structure of the Atom
4/21/2017
Chapter 4
The Structure of the Atom

2. Early Theories of Matter
The concept of the atom developed over a few thousand years.
Consider that there were no controlled experiments and few tools for scientific exploration.

3. The power of the mind and intellectual thought were the primary ways to truth about the universe.
Curiosity drove discovery 
Philosophers, scholarly thinkers, speculated about the nature of matter.
Ideas came from their own life experiences.

4. Early Greek philosophers thought that matter was made of Earth, Wind, Water, and Fire.

5. No methods to test the validity of these ideas.
Commonly accepted that matter could be divided endlessly into smaller and smaller pieces…..
The Continuous Theory of Matter
No methods to test the validity of these ideas.

6. Democritus a Greek Philosopher
the first to propose the idea that matter was not infinitely divisible. discontinuous theory
He proposed the idea that matter was made up of tiny individual particles called atomos. 

7. The English word for atom comes from this Greek word atomos.
 He believed that atoms could not be created, destroyed, or further divided (discontinuous theory).

8. His belief in atoms was amazingly accurate.
However, his ideas were met with much criticism from other philosophers.
The harshest criticism was from the influential Greek philosopher, Aristotle.

9. Aristotle
Aristotle rejected the theory of atoms because it did not agree with own ideas on nature.
He did not believe that atoms could move through empty space.
He did not believe that the “nothingness” of empty space could exist.

10. Democritus Rejected
Because Democritus had no way to answer the challenges to his ideas, his theory was eventually rejected.
It is important to realize that these ideas were just that—ideas and not science.
Incredibly, the ideas of Aristotle were so great and the science so primitive that his denial of the existence of the atom went unchallenged for two thousand years

11. John Dalton In the 19th century John Dalton an English School teacher
revived and revised the work of Democritus.
This time it was based on scientific research conducted by Dalton.

12. Dalton’s Atomic Theory
1808
marks the beginning of modern atomic theory. 

13. Dalton was wrong about:
the atom being divisible (atoms are divisible into subatomic particles)
all atoms of an element having identical properties. (atoms of an element can have slightly different masses.)

14. Dalton’s Model of the Atom

15. Defining the Atom
An atom is the smallest part of an element that retains the characteristics of the element.
How small is a typical atom?
In the year 2000, the world population was about 6 billion people.
In comparison, a copper penny contain about 5 billion times as many atoms of copper.
     World population
  Atoms in a Penny
      Diameter of a copper atom 1.28 x m

16. Next Set of Questions What is an atom like? How are atoms shaped?
Is the atom composed of other particles?

17. The Discovery of the Electron
British scientist named J. J.Thompson
Used a glass tube, called a cathode-ray tube, connected to two electrodes.

18. Cathode Ray Tube

19. One electrode was attached to one end of the tube.
It was negative in charge and was thus called the cathode.
At the other end was attached a positive electrode, the anode.

21. Thomson’s Experiment
Thomson placed a paddle wheel inside the tube, between the cathode and the anode.
He passed an electric current through a variety of gases in the tube. The gases within the tube were changed with each experiment.

22. Thomson made the following observations
The current passing through the tube creates different colors in different gases.
The anode (positive end) glows where as the cathode (negative end) does not.
 An object placed within the cathode-ray casts a shadow.
The paddle-wheel spins in the direction of the anode.

23. Conclusions:
The different colors that were created by using different gases showed that atoms of different elements possessed different energies.  
The cast shadow was thought to be due to the beam of light created by the cathode-ray. However, the experiment made with the spinning paddle-wheel showed that the cathode-ray was composed of particles.
The fact that the anode glowed and the cathode did not shows that the cathode-ray travels from a negative potential to a positive potential. This was one of the most important findings of J. J. Thomson's experiment and was repeated using another apparatus.

24. To test the polarity of the cathode ray, Thompson replaced the paddle-wheel with another pair of electrodes.
When the current was passed through this new apparatus, Thompson found that the cathode-ray was "bent" towards the positive electrode and repelled by the negative electrode.  
Thompson attributed the "bending" of the cathode-ray to charge-charge repulsion of the second cathode and the negatively charged particles in the cathode-ray

26. Electrons He named these negatively charged particles electrons.
Since Thompson could get these results regardless of what element he used within the tube, he concluded that all atoms have electrons.

27. Thomson’s Model of the Atom
The Plum Pudding Model

28. Oil Drop Experiment
In 1909, American scientist, Robert Millikan, determined the charge of an electron in his famous Oil Drop Experiment.

29. Discovery of the Nucleus
1911 Ernst Rutherford
Rutherford’s Gold Foil Experiment
determined that nearly all of the mass of the atom consisted of a positively charged nucleus
    

30. Rutherford’s Gold Foil Experiment
He did this by measuring the deflection of helium nuclei by gold atoms.

31. Predicted Results

32.   While most atoms were not deflected much at all, a few were deflected by 180 degrees.

33. How was this possible?
Rutherford calculated that the only way this was possible was if the gold atoms consisted of a cloud of electrons surrounding the very dense, positively charged nucleus, for only then could the gold atom transfer enough momentum to the ions to turn them around.

34. Rutherford’s Conclusions
Atoms are mostly empty space and the nucleus is incredibly dense.
If the nucleus were a marble, the atom would be the size of a football field

35. Rutherford ‘s Nuclear Model
By 1919 Rutherford refined his concept of the nucleus and concluded that it contained protons.
Protons have a charge of +1.
This model could not account for the mass of the atom.
20 more years before this mystery was solved.

36. Discovery of the Neutron
1932
James Chadwick, an English Physicist
He showed that the nucleus also contained another particle, the neutron
The neutron is neutral, meaning it has no charge.
The neutron has a mass equal to the proton.

37. Composition of the Nucleus
Protons
Mass of the proton = X 10-24grams
Positive Charge
Neutrons
Mass = X 10-24grams
No charge 

38. Properties of Subatomic Particles
Symbol
Electrical charge
Relative mass
Actual mass (g)
Electron e 1-
1/1840
9.1110-28
Proton p 1+ 1
1.67310-24
Neutron n0
1.67510-24

39. Nuclear Model of the Atom
An atom is an electrically neutral particle composed of protons, neutrons, and electrons. 
Atoms are spherical in shape with a tiny, dense nucleus of positive charge surrounded by one or more negatively charged electrons 
Nucleus contains 99.7 % of the mass of an atom.

40. Subatomic Particles
Most of the atom consists of fast-moving electrons traveling through the space around the nucleus.
The number of protons is equal to the number of electrons since atoms are neutral in charge.
The number of Protons in the nucleus determines the identity of an atom. (Atomic Number) 
An Element consists of atoms with the same number of protons. 

41. How Atoms Differ
Henry Moseley discovered that atoms of each element contain a unique positive charge in their nucleus.
Therefore, the proton number of an atom identifies the atom.
The number of protons in an atom is known as atomic number.
Example: All Hydrogen atoms have one proton

42. Atomic Number = Protons
Atomic Number(Z): the number of protons in the nucleus of each atom of the element
Atomic Number = Protons
Elements are arranged in the periodic table from left to right in order of increasing atomic number.
Atomic number identifies the element.

43. # of protons = # of electrons
  Atoms are neutral, therefore:
# of protons = # of electrons 

44. Isotopes and Mass Number
Naturally occurring elements are a mixture of atoms that have different numbers of neutrons.
Atoms with the same number of protons, but different number of neutrons are called Isotopes.

45. Hydrogen Has Three Isotopes:
Protium 1 proton 1 electron 0 neutrons
Deuterium 1 proton electron 1 neutron
Tritium 1 proton 1 electron 2 neutrons

46. Isotopes and Mass Number
Isotopes: atoms of the same element that have different masses.
Different number of neutrons so mass is different.
Atoms of different isotopes have different masses so identity is given by name and mass.
Mass Number: the total number of protons and neutrons in the nucleus of an isotope

47. Isotope (Nuclear) Symbols:
Consists of three parts
the symbol of the element
the atomic number of the element
the mass number of the specific isotope.

48. Nuclear Symbol

49. Isotope Names Element Name-Mass Number Examples: Helium-4 Potassium-39
Hydrogen-3

50. Neutrons = Mass Number – Atomic Number
# of Protons: _________
# of Electrons: _______
# of Neutrons: ________
Name of Isotope: ____________

51. Example Atomic Number: _________ Mass Number: __________
# of Protons: _________
# of Electrons: _______
# of Neutrons: ________
Name of Isotope: ____________

52. Example Atomic Number: _________ Mass Number: __________
Atomic Number: _________
Mass Number: __________
# of Protons: _________
# of Electrons: _______
# of Neutrons: ________
Isotope Name: _________

53. Draw the isotope symbol for Calcium with 21 neutrons.

54. Comparing Potassium Isotopes
protons = ________
electrons = ______
neutrons = _______
Potassium-40
Potassium-39
Protons = __________
Electrons = _________
Neutrons = __________ 

55. Relative Atomic Masses
Atomic masses measured in grams are very small. 
Example: An atom of Oxygen-16 has a mass of x 19-23grams.
It is most convenient to measure mass in relative atomic masses.

56. To set up a relative scale of atomic masses:
One atom is chosen and assigned a relative mass value,
All the other masses are expressed in relation to this defined standard.
Carbon-12 is the chosen standard.
A single atom of C-12 is assigned a mass of exactly 12 atomic mass units(u).

57. An AMU One amu is exactly 1/12th of the mass of a carbon-12 atom
x 10-24grams.
The atomic mass of a carbon-12 atom is exactly 12 u. 
Atomic Mass: The mass of an atom expressed in atomic mass units.

58. Average Atomic Mass Average Atomic Mass Average Atomic Mass depends on
given on the periodic table
weighted average for the naturally occurring mixtures of isotopes in each element
Average Atomic Mass depends on
mass and relative abundance of the isotopes.

59. Calculating average atomic mass:
Step 1:
Isotope 1: Atomic mass x relative abundance = mass contribution 1
Isotope 2: Atomic mass x relative abundance = mass contribution 2
Step 2: Add the mass contributions results.

60. Example Problem: Cu-63 0.6917 x 62.939 598u = 43.535u
Naturally occurring copper consists of 69.17% Cu-63, mass of u and % Cu-65, mass of u
Cu x u = u
Cu x u =
63.552u

61. Example Problem 2
Gallium occurs in nature as a mixture of two isotopes. They are Ga-69 with a % abundance and a mass of amu and Ga-71 with a % abundance and an atomic mass of Calculate the atomic mass of gallium.

62. Example 3
The element chlorine occurs in nature as a mixture of two isotopes. Chlorine-35 has an atomic mass of amu and makes up 75.77% of chlorine atoms. Chlorine-37 atoms make up the remaining 24.23% of all chlorine. Use the average atomic mass of chlorine from the periodic table to calculate the atomic mass of Cl-37 atoms.

63. Example 4
The atomic mass of bromine given in the periodic table is amu, which is very close to 80 amu. Use a reference book to find the percent of Br-80 in naturally occurring bromine. Explain the value of the atomic mass of bromine from the data you find.