1. Chemistry Chapter 5

2. Democritus – 4th century B. C
Democritus – 4th century B.C. Greek philosopher who first proposed that matter is composed of tiny, indivisible particles called atoms.
The first atomic model was not proposed until 2000 years later!

3. John Dalton, an English schoolteacher, proposed the following atomic theory, in 1808.

4. Dalton’s Atomic Theory:
1. All elements are composed of tiny, indivisible particles called atoms.
2. Atoms of the same element are identical and differ from those of other elements.

5. 3. Atoms cannot be subdivided, created, nor destroyed.
4. Different elements combine in simple whole-number ratios to form chemical compounds.

6. 5. Chemical reactions occur when atoms are separated, joined, or rearranged.

7. Law of conservation of mass – mass is neither created nor destroyed
Law of conservation of mass – mass is neither created nor destroyed. (See #3 and 5 above)

8. Law of definite proportions – different samples of any pure compound contain the same elements in the same proportions by mass.
For instance, both 23g of water and 72,000g of water will be 11.2% hydrogen and 88.9% oxygen by mass. (See #4 above)

9. Law of multiple proportions – if two or more different compounds are composed of the same two elements, A and B, then the ratio of the masses of element B that combine with a certain mass of element A is always a ratio of small whole numbers. (See #4 above)
Examples: NO, NO2, N2O4, N2O5

10. The Modern Atomic Theory
Not all aspects of Dalton’s theory have been proven correct:
1) atoms are divisible and
2) a given element can have atoms with different masses.

11. The important aspects from Dalton’s theory are that:
1) all matter is composed of atoms and
2) atoms of any one element differ in properties from atoms of another element.

12. An atom is defined as the smallest particle of an element that retains the properties of that element.
We now know that atoms are not indivisible. They can be broken down into even smaller, more fundamental particles.

13. 3 Subatomic Particles: electron, proton, and neutron.

14. Discovery of the Electron
Electrons are negatively charged subatomic particles.

15. J. J. Thomson, an English physicist discovered them in 1897 by doing experiments involving passing an electric current through a tube containing gases at low pressure.

17. Thomson found that a beam, called a cathode ray tube, traveled from the cathode (-) to the anode (+), and that this beam was repelled by negative electrical charge.
Since he knew opposite charges attract and like charges repel, he proposed that this cathode ray was a stream of tiny negatively charged particles moving at high speed. He named them electrons.

18. Further studies by Robert Millikan, an American scientist, led to the determination of the properties of the electron in 1909 with his oil-drop experiment .
Oil-drop experiment

19. Oil-Drop Experiment:
His experiment involved measuring the charge on tiny oil drops. The charge of each drop was a multiple of 1.60x10-19 C. This was the charge of an individual electron. He used this charge, along with Thomson’s charge/mass ratio, to determine the mass of an electron.
1.6x10-19C x g = 9.11x10-28g
1.76x108C

20. An electron has a charge of -1 and a mass of 1/1840, or 9
An electron has a charge of -1 and a mass of 1/1840, or 9.109x10-31 kg, the mass of a hydrogen atom.

21. Discovery of the Atomic Nucleus
Protons: Scientists knew that atoms are neutral (have no charge). So, if there are electrons with negative charge, there must also be particles with a positive charge. This led to the discovery of the proton.

22. Protons have a +1 charge and are 1840 times more massive than the electron.

23. Neutrons: Sir James Chadwick discovered them in 1932.
Neutrons have no charge, but have essentially the same mass as a proton.

24. Once subatomic particles were discovered, Dalton’s model of the atom had to be modified.

25. The second model of the atom (Thomson’s “Plum Pudding Model”) proposed that electrons were evenly distributed throughout an atom filled uniformly with positively charged material.

26. In 1911, Ernest Rutherford, a New Zealand native, tested this model with his gold foil experiment.

27. Gold Foil Experiment:
He shot a beam of massive alpha particles (He+2) at a very thin sheet of gold foil.
He expected the alpha particles to pass easily through the foil, with little deflection.
He was shocked to see that even though most passed through without deflection, a small fraction was deflected at large angles and some even bounced straight backwards!
To explain this, he modified Thomson’s atomic model.

28. Figure 02-10
Title:
Rutherford's experiment on the scattering of α particles.
Caption:
The red lines represent the paths of the α particles. When the incoming beam strikes the gold foil, most particles pass straight through the foil, but some are scattered.
Notes:
Keywords:

29. Gold-Foil Experiment Figure 02-11 Title:
Rutherford's model explaining the scattering of α particles.
Caption:
The gold foil is several thousand atoms thick. Because most of the volume of each atom is empty space, most α particles pass through the foil without deflection. When an α particle passes very close to a gold nucleus, however, it is repelled, causing its path to be altered.
Notes:
Keywords:
Gold-Foil Experiment

30. Rutherford’s model of the atom stated that the atom is mostly empty space with all the positive charge and almost all of the mass concentrated in a small region, which he called the nucleus.

32. The tiny nucleus is composed of protons and neutrons.
This is an image of silicon atoms arranged on a face of a crystal. It is impossible to "see" atoms this way using ordinary light. The image was made by a Scanning Tunneling Microscope, a device that "feels" the cloud of electrons that form the outer surface of atoms.

33. How small is the nucleus?
If the atom were the size of a football stadium, the nucleus would be about the size of a marble …10,000 times smaller!

34. Electron e- -1 9.109x10-31 Proton p+ +1 1 1.673x10-27 Neutron no
Properties of Subatomic Particles (Table 3-1, pg. 74)
Particle
Symbols
Charge
Mass Number
Actual Mass (kg)
Electron e- -1
9.109x10-31
Proton p+ +1 1
1.673x10-27
Neutron no
1.675x10-27

36. Going Further

37. MORE SUBATOMIC PARTICLES???!!!
Leptons (elementary particles)
Electron
Mu-meson (muon) – more massive than electrons
Tau-meson (tau) – more massive than electrons
3 types of neutrinos – almost massless

38. Hadrons (made of quarks)
Hadrons (made of quarks)
Mesons of many types
Composed of a quark and an antiquark
Baryons
Composed of 3 quarks of different colors
protons (2 up quarks + 1 down quark)
neutrons (1 up quark and 2 down quarks)

39. Every particle has an antiparticle. The
antiparticle of the electron is the positron (e+)
Gluons – hold quarks together.

40.  6 “flavors” of quarks
1. up 3. top (truth) 5. strange 2. down 4. bottom (beauty) 6. charm

41. “Colors” of quarks (charge)
+2/3 up, top, charm
-1/3 down, bottom, strange

42. Going Further

43. Atomic number – the number of protons in the nucleus of an atom.
The atomic number identifies an element.

44. 1) What is the atomic number of aluminum?13
2) Which element has 26 protons?
Fe - iron

45. In a neutral atom, the number of protons (p+) equals the number of electrons (e-).
3) How many p+ does a neon atom have? 10
4)How many e- does a neon atom have?
5)An element with atomic number = 6 has how many e-? 6

46. Mass Number – The total number of protons and neutrons in the nucleus of an atom.
mass number – atomic number = number of neutrons

47. 6) Oxygen-18 has a mass number of 18
6) Oxygen-18 has a mass number of 18. How many protons, neutrons and electrons does it have?
Atomic # = 8
Therefore, p+ = 8
e- = 8
n0 = =10

48. Isotopes – Atoms that have the same number of protons but different numbers of neutrons.

49. When writing nuclear chemical symbols for isotopes, the mass number is written as a superscript and the atomic number is written as a subscript before the element symbol. Ex. 146C. This can also be written in hyphen notation as carbon-14.

51. Ex. Hydrogen has three isotopes:
Name:
Hydrogen – 1
Hydrogen – 2 (protium)
Hydrogen – 3 (tritium)
Percent:
99.985%
0.015%
negligible
Symbol: 1H 2H 3H
Atomic #: 1
Mass #: 2 3
p+:
e-:
no:

52. Atomic Mass is given in units called atomic mass units.
Atomic Mass Unit(amu) – One amu is defined as 1/12 the mass of a carbon-12 atom.

53. An atomic mass unit can also be called a Dalton.
A carbon-12 atom has 6 p+ and 6 n0, and its mass is set at 12 amu’s.

54. But, the atomic masses on the periodic table are not whole numbers (mainly) because most elements occur in nature as a mixture of isotopes. The masses are a weighted average reflecting the relative abundance of each isotope.

55. Average atomic mass – weighted average mass of the atoms in a naturally occurring sample of the element.

56. Average atomic mass = (Percent abundance 1)(Isotopic mass 1) + (Percent abundance 2)(Isotopic mass 2)+ etc.
*when doing weighted averages, change percents to decimal form: ex instead of 55%.*

57. The average atomic mass of an element depends on both the mass and the relative abundance of each of the element’s isotopes.

58. 7) Calculate the atomic mass of bromine
7) Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative abundance of amu (50.69%) and (49.31%).
(78.92 amu) X (0.5069) = amu
(80.92 amu) X (0.4931) = amu
Atomic mass Br = amu

59. 8) Naturally occurring iron consists of four isotopes with the abundances indicated here. From the masses and relative abundances of these isotopes, calculate the atomic weight of naturally occurring iron.
Isotope Isotopic Mass (amu) % Abundance
Iron
Iron
Iron
Iron
( amu) x (0.0582) = amu
( amu) x (0.9166) = amu
( amu) x (0.0219) = amu
( amu) x (0.0033) = amu
Atomic mass of Fe amu

60. 9) The atomic weight of gallium is 69. 72 amu
9) The atomic weight of gallium is amu. The masses of the naturally occurring isotopes are amu and for 6931Ga and amu for 7131Ga. Calculate the percent abundance of each isotope.
Let x = fraction of 6931Ga. Then (1- x) = fraction of7131Ga.
x( amu) + (1 – x)( amu) = amu
x – x = 69.72
x = -1.20
x = 0.600
Multiply by 100 to get percentage
x = = fraction of 6931Ga; therefore, 60.0% 6931Ga
(1-x) = = fraction of 7131Ga; therefore, 40.0%7131Ga

61. 10) The atomic weight of rubidium is 85. 4678 amu
10) The atomic weight of rubidium is amu. The two naturally occurring isotopes of rubidium have the following masses: rubidium-85, amu; rubidium-87, amu. Calculate the percent of rubidium-85 in naturally occurring rubidium.
Let x = fraction of rubidium-85. Then (1- x) = fraction of rubidium-87.
x( amu) + (1 – x)( amu) = amu
x – x =
x =
x = 0.722
72.2% rubidium-85

63. The Periodic Table:
A Russian chemist, Dmitri Mendeleev ( ), constructed the first periodic table by listing the elements in order of increasing atomic mass and arranging them according to similarities in their properties. He was able to predict the physical and chemical properties of missing elements.

75. Nuclear Chemistry

76. Types of Radiation

77. 1. alpha radiation -helium nuclei emitted from a radioactive source -2 protons and 2 neutrons - +2 charge - symbol is  -don't travel far and aren't very penetrating -stopped by a sheet of paper -can't penetrate skin cells but dangerous if ingested -very common with very heavy radioactive nuclides

79. 2. Beta radiation - fast moving electron formed by the decomposition of a neutron of an atom. The neutron breaks into a proton and an electron. The proton stays in the nucleus and the electron is ejected. (net effect: neutron changes into proton) -much smaller than alpha particles -symbol is  -charge is -1 -much more penetrating than alpha particles -stopped by aluminum foil or thin pieces of wood.

80. -Nuclides that have too high of a neutron/proton ratio tend to undergo beta particle emission.

82. 3. Gamma radiation - electromagnetic radiation (high energy) emitted from a nucleus as it changes from an excited state to a ground energy state -often emitted along with alpha or beta radiation -symbol is  -have no mass and no charge -high energy photon

83. -the emission of gamma rays is one way that a nucleus with excess energy (in an excited nuclear state) can relax to its ground state -extremely penetrating, very dangerous -stopped somewhat by several feet of concrete or several inches of lead

84. Fission- splitting a heavy nucleus into two nuclei with smaller mass numbers. - used for nuclear energy Ex production of neutrons causes a chain reaction (which must be controlled) -1 kg of uranium-235 is equivalent to 20,000 tons of dynamite

87. Fission in a nuclear reactor is carefully controlled
*Fission in a nuclear reactor is carefully controlled. Much of the energy is heat. This energy is used to produce steam and subsequently, electricity. *A coolant (usually water) is needed. *The water (or carbon) also acts as a moderator. It slows the neutrons down so that they can be captured by the U-235 fuel. *Control rods made of cadmium are present to absorb excess neutrons to slow down the reaction. They can be raised or lowered into the reactor core.

90. Fusion- combining two light nuclei to form a heavier, more stable nucleus -stars produce their energy this way -currently, high temperatures are necessary in order to initiate fusion -possible future energy source

92. Going Further

93. Methods of detection:

94. Geiger counter- uses a gas-filled metal tube to detect radiation
Geiger counter- uses a gas-filled metal tube to detect radiation -primarily detects beta

95. Scintillation counter- -uses a specially coated screen todetect radiation -can detect all types

96. Film badges- -detect beta and gamma

97. Radioisotopes in Research and Medicine
Tracer: Iodine-131 is used to check for thyroid problems
Tracer: Radioactive barium is used to check for digestive system problems
Radiation sources: Cobalt-60 and Cesium-137 (among many others) are used as radiation sources for cancer treatment