1. 2 Early Thoughts 3 Democritus (about 470-370 B.C.) thought that all forms of matter were made of...

1

2

Early ThoughtsEarly Thoughts

3

Democritus (about 470-370 B.C.) thought that all forms of matter were made of tiny particles called “atoms” from the Greek “atomos” indivisible.

• The earliest models of the atom were developed by the ancient Greek philosophers.

Leucippus of Miletus (490-??? B.C.). First to introduce the idea of the atom, an indivisible unit of matter. This idea was later extended by his student, Democritus.

4

According to Democritus atoms are:

• Unchangeable and indivisible.

• Identical except for their size and shape.

• Always in motion.

5

Democritus imagined that atoms of iron were shaped like coils, making iron rigid, strong, and malleable. Atoms of fire were sharp, lightweight, and yellow.

6

• Aristotle (384-322 B.C.) rejected the theory of Democritus and endorsed that of Empedocles that stated that matter was made of 4 elements: air, earth, fire , and water.

7

• Empedocles (492-432 B.C.) believed that these elements have always existed in fixed amounts, and that there two major forces which act upon these elements to both create and destroy: Love and Strife. According to legend, he died by falling into a volcano's crater after failing to become a god as he predicted.

8

– Aristotle’s influence dominated the thinking of scientists and philosophers until the beginning of the 17th century

9

10

Alchemical Symbols

antimony arsenic bismuth

copper gold iron

11

Alchemical Symbols

magnesium mercury phosphorus

platinum potassium silver

12

Alchemical Symbols

sulfur tin zinc

lead

13

Dalton’s Modelof the Atom

Dalton’s Modelof the Atom

14

2000 years after Aristotle, John Dalton, an English schoolmaster, proposed his model of the atom–which was based on experimentation.

15

16

2. Atoms of the same element are alike in mass and size.

3. Atoms of different elements have different masses and sizes.

4. Chemical compounds are formed by the union of two or more atoms of different elements.

Dalton’s Atomic Theory

Modern research has demonstrated that atoms are composed of subatomic particles.

Atoms under special circumstances can be decomposed.

1. Elements are composed of minute indivisible particles called atoms.

17

Dalton’s Atomic Theory

5. Atoms combine to form compounds in simple numerical ratios, such as one to one, two to two, two to three, and so on.

6. Atoms of two elements may combine in different ratios to form more than one compound.

18

Dalton’s atoms were individual particles.

Atoms of each element are alike in mass and size.

19

Dalton’s atoms were individual particles.

Atoms of different elements are not alike in mass and size.

20

Daltons atoms combine in specific ratios to form compounds.

H 2 =

O 1H 1

= O 1

21

Composition of Compounds

Composition of Compounds

22

The Law of Definite Composition

A compound always contains two or more elements combined in a definite proportion by mass.

23

• The percent by mass of hydrogen in water is 11.2%.

• The percent by mass of oxygen in water is 88.8%.

• Water always has these percentages. If the percentages were different the compound would not be water.

• Water always contains the same two elements: hydrogen and oxygen.

Composition of Water

24

• The percent by mass of hydrogen in hydrogen peroxide is 5.9%.

• The percent by mass of oxygen in hydrogen peroxide is 94.1%.

• Hydrogen peroxide always has these percentages. If the percentages were different the compound would not be hydrogen peroxide.

• Hydrogen peroxide always contains the same two elements: hydrogen and oxygen.

Composition of Hydrogen Peroxide

25

The Law of Multiple Proportions

Atoms of two or more elements may combine in different ratios to produce more than one compound.

26

Mass Hydrogen(g)

Mass Oxygen(g)

Water 1.0 8.0

Hydrogen Peroxide

1.0 16.0

mass of oxygen in hydrogen peroxide 16g 2= = ¹

mass of oxygen in water 8g 1

Combining Masses of Hydrogen and Oxygen

Hydrogen peroxide has twice as much oxygen (by mass) as does water.

27

• The formula for water is H2O.

• The formula for hydrogen peroxide is H2O2.

• Hydrogen peroxide has twice as many oxygens per hydrogen atom as does water.

Combining Ratios of Hydrogen and Oxygen

28

The Nature of Electric ChargeThe Nature of Electric Charge

29

Surrounding the atomic nucleus are electrons. The name electron comes from the Greek word for amber, a brownish-yellow fossil resin studied by the early Greeks. They found that when amber was rubbed by a piece of cloth, it attracted such things as bits of straw. This phenomenon, known as the amber effect, remained a mystery for almost 2000 years. In the late 1500s other materials that behaved like amber were called “electrics”. The concept of electric charge awaited experiments by Benjamin Franklin nearly two centuries later. Franklin experimented with electricity and postulated the existence of an electric fluid that could flow from place to place. An object with an excess of this fluid he called electrically positive, and one with a deficiency of the fluid he called electrically negative. The fluid was thought to attract ordinary matter but to repel itself. We still follow Franklin’s lead in how we define positive and negative electricity. Franklin’s 1752 experiment with the kite in the lightning storm showed that lightning is an electrical discharge between clouds and the ground. This discovery told him that electricity is not restricted to solid or liquid objects and that it can travel through a gas.

30

Ben Franklin (1746-52 ) flew kites to demonstrate that lightning is a form of static electricity (ESD). He would run a wire to the kite and produce sparks at the ground, or charge a Leyden jar. This led Franklin to invent the lightning rod. Franklin also made several electrostatic generators with rotating glass balls to experiment with. These experiments led him to formulate the single fluid (imponderable fluid) theory of electricity. Previous theories had held there were two electrical fluids and two magnetic fluids. Franklin theorized just one imponderable electrical fluid (a fluid under conservation) in the universe. The difference in electrical charges was explained by an excess ( + ) or defect ( - ) of the single electrical fluid. This is where the positive ( + ) and negative ( - ) symbols come from in electrical science.

31

Discovery of IonsDiscovery of Ions

32

• Michael Faraday discovered that certain substances, when dissolved in water, conducted an electric current.

• He found that atoms of some elements moved to the cathode (negative electrode) and some moved to the anode (positive electrode).

• He concluded they were electrically charged and called them ions (Greek wanderer).

33

Michael Faraday

34

• Svante Arrhenius reasoned that an ion is an atom (or a group of atoms) carrying a positive or negative electric charge.

• Arrhenius accounted for the electrical conduction of molten sodium chloride (NaCl) by proposing that melted NaCl dissociated into the charged ions Na+ and Cl-.

NaCl → Na+ + Cl-Δ

35

• In the melt the positive Na+ ions moved to the cathode (negative electrode). Thus positive ions are called cations.

• In the melt the negative Cl- ions moved to the anode (positive electrode). Thus negative ions are called anions.

NaCl → Na+ + Cl-

36

Svante Arrhenius

37

Subatomic Partsof the Atom

Subatomic Partsof the Atom

38

An atom is very SmallSmall

39

The diameter of an atom is 0.1 to 0.5 nm.

This is 1 to 5 ten billionths of a meter.

If the diameter of this dot is 1 mm then 10 million hydrogen atoms would form a line across the dot.

Even smaller particles than atoms exist. These are called subatomic particles.

40

Subatomic ParticlesSubatomic Particles

41

ElectronElectron

42

In 1875 Sir William Crookes invented the Crookes tube.

43

• Crookes tubes experiments led the way to an understanding of the subatomic structure of the atom.

44

• Crookes tube emissions are called cathode rays.

• Below are Crookes cathode-ray tubes. The cathode-rays (streams of electrons) can be clearly seen.

45

"Maltese Cross" Crookes Tube Demonstrates that radiant matter is

blocked by metal objects

46

Other Interesting Crookes Tubes May Be Found At the Sites Below:

• http://www.sparkmuseum.com/GLASS.HTM

• http://www.oneillselectronicmuseum.com/page9.html

47

In 1897 Sir Joseph Thomson demonstrated that cathode rays:

• travel in straight lines.

• are negative in charge.

• are deflected by electric and magnetic fields.

• produce sharp shadows

• are capable of moving a small paddle wheel.

48

Paddle Wheel

49

Thomson’s Apparatus

batteries

50

Thomson’s Lab

51

J.J. Thomson determined and is given credit for finding:

• The charge to mass (e/m) ratio of the cathode ray.

• The cathode ray was re-named the “electron”.

• Thomson “discovered” the electron.

http://www.aip.org/history/mod/fission/fission1/01.html

52

Can atoms be split apart? Does each atom have inner workings? Parts which can be separated? Parts which can perhaps be put to some use? These questions had already come to mind in 1898, when J. J. Thomson isolated the electron. That was the first solid proof that atoms are indeed built of much tinier pieces. Thomson speaks of the electron in this recorded passage...

Could anything at first sight seem more impractical than a body which is so small that its mass is an insignificant fraction of the mass of an atom of hydrogen, which itself is so small that a crowd of these atoms equal in number to the population of the whole world would be too small to have been detected by any means then known to science.

53

Robert Millikan

• Determined the charge of the electron.

• Experiment called the Oil Drop Experiment.

54

Oil Drop Apparatus

55

Apparatus Used by Millikan

56

Modern Apparatus

57

ProtonProton

58

• Eugen Goldstein, a German physicist, first observed protons in 1886:

• Thomson determined the protons’ characteristics.

• Thomson showed that atoms contained both positive and negative charges.

• This disproved the Dalton model of the atom which held that atoms were indivisible.

59

Thomson’s Plum-Pudding Model of the Atom

60

NeutronNeutron

61

• James Chadwick discovered the neutron in 1932.

• Its actual mass is slightly greater than the mass of a proton.

Particle symbol Actual charge (coulombs)

Relative charge

Actual mass (g)

Relative mass (amu)

Electron e– –1.602 10–19 –1 9.11 10–

28

0

Proton p +1.602 10–19 +1 1.67 10–

24

1

Neutron n

0 0 1.68 10–

24

1

Electric charge and mass of electrons, protons, & neutrons

63

IonsIons

64

• Positive ions were explained by assuming that a neutral atom loses electrons.

• Negative ions were explained by assuming that extra electrons can be added to atoms.

65

When one or more electrons are lost from an atom, a cation is formed.

66

When one or more electrons are added to a neutral atom, an anion is formed.

67

The Nuclear AtomThe Nuclear Atom

68

X-rays were discovered by Wilhelm Röentgen in 1895

69

• Röentgen observed that a vacuum discharge tube enclosed in a thin, black cardboard box had caused a nearby piece of paper coated with the salt barium platinocyanide to phosphorescence.

• From this and other experiments he concluded that certain rays, which he called X-rays, were emitted from the discharge tube, penetrated the box, and caused the salt to glow.

70

• Radioactivity was discovered by Henri Becquerel in 1896.

71

• Shortly after Röentgen’s discovery, Antoine Henri Becquerel attempted to show a relationship between X-rays and the phosphorescence of uranium salts.

• Becquerel wrapped a photographic plate in black paper, sprinkled a sample of a uranium salt on it, and exposed it to sunlight.

72

• When Becquerel attempted to repeat the experiment the sunlight was intermittent.

• He took the photographic plate wrapped in black paper with the uranium sample on it, and placed the whole setup in a drawer.

73

• Several days later he developed the film and was amazed to find an intense image of the uranium salt on the plate.

• He repeated the experiment in total darkness with the same result.

• This proved that the uranium salt emitted rays that affected the photographic plate, and that these rays were not a result of phosphorescence due to exposure to sunlight.

74

Radioactivity is the spontaneous emission of particles and/or rays from the nucleus of an atom.

• Two years later, in 1898, Marie Curie coined the name radioactivity.

75

Marie Curie, in a classic experiment, proved that alpha and beta particles are oppositely charged.

radiation passes between the poles of an electromagnet

a radioactive source was placed inside a lead block

Alpha rays are less strongly deflected to the negative pole.

Gamma rays are not deflected by the magnet.Beta rays are strongly deflected to the positive pole.

three types of radiation are detected by a photographic plate

76

The Rutherford Experiment

The Rutherford Experiment

77

Ernest Rutherford

78

• In 1899 Rutherford began to investigate the nature of the rays emitted by uranium.

• He found two particles in the rays. He called them alpha and beta particles.

79

• Rutherford in 1911 performed experiments that shot a stream of alpha particles at a gold foil.

• Most of the alpha particles passed through the foil with little or no deflection.

• He found that a few were deflected at large angles and some alpha particles even bounced back.

80Rutherford’s alpha particle scattering experiment.

81

• An electron with a mass of 1/1837 amu could not have deflected an alpha particle with a mass of 4 amu.

• Rutherford knew that like charges repel.

• Rutherford concluded that each gold atom contained a positively charged mass that occupied a tiny volume. He called this mass the nucleus.

82

• If a positive alpha particle approached close enough to the positive mass it was deflected.

• Most of the alpha particles passed through the gold foil. This led Rutherford to conclude that a gold atom was mostly empty space.

83

• Because alpha particles have relatively high masses, the extent of the reflections led Rutherford to conclude that the nucleus was very heavy and dense.

84Deflection and scattering of alpha particles by positive gold nuclei.

Deflection

Scattering

85

Ideas about the atom were refined by one of Thomson's students, Ernest Rutherford. He showed that the mass in an atom is not smeared out uniformly throughout the atom, but is concentrated in a tiny, inner kernel: the nucleus. Rutherford wanted to understand the nucleus, not for any practical purpose, but because he was attracted to the beauty of its simplicity. Fundamental things should be simple not complex. Here is how he explains himself in 1931...

The bother is that a nucleus, as you know, is a very small thing, and we know very little about it. Now, I had the opinion for a long time, that's a personal conviction, that if we knew more about the nucleus, we'd find it was a much simpler thing than we suppose, that these fundamental things I think have got to be fairly simple. But it's the non-fundamental things that are very complex usually. I am always a believer in simplicity being a simple person myself.

86

• The gamma ray, a third type of emission from radioactive material, was discovered by Paul Villard in 1900.

87

Alpha, Beta, and Gamma Radiation

NameNuclide Symbol

Particle Symbol

Mass (amu)

Charge

Alpha 4 +2

Beta –1

Gamma Ray 0 0

42 He

0-1e

00

1

1837

88

General Arrangement of

Subatomic Particles

General Arrangement of

Subatomic Particles

89

• Rutherford’s experiment showed that an atom had a dense, positively charged nucleus.

• Chadwick’s work in 1932 demonstrated the atom contains neutrons.

• Rutherford also noted that light, negatively charged electrons were present in an atom and offset the positive nuclear charge.

90

• Rutherford put forward a model of the atom in which a dense, positively charged nucleus is located at the atom’s center.

• The negative electrons surround the nucleus.

• The nucleus contains protons and neutrons

91

92

Atomic Numbers of the Elements

Atomic Numbers of the Elements

93

• The atomic number of an element is equal to the number of protons in the nucleus of that element.

• The atomic number of an atom determines which element the atom is.

94

Every atom with an atomic number of 1 is a hydrogen atom.

Every hydrogen atom contains 1 proton in its nucleus.

95

Every atom with an atomic number of 6 is a carbon atom.

Every carbon atom contains 6 protons in its nucleus.

96

1 proton in the nucleus

atomic number

Every atom with an atomic number of 1 is a hydrogen atom.

1H

97

6 protons in the nucleus 6C

atomic number

Every atom with an atomic number of 6 is a carbon atom.

98

92 protons in the nucleus 92U

atomic number

Every atom with an atomic number of 92 is a uranium atom.

99

Isotopes of the Elements

Isotopes of the Elements

100

• Atoms of the same element can have different masses.

• They always have the same number of protons, but they can have different numbers of neutrons in their nuclei.

• The difference in the number of neutrons accounts for the difference in mass.

• These are isotopes of the same element.

Greek roots isos (equal) and topos (place). Hence: "the same place," meaning that different isotopes of a single element occupy the same position on the periodic table.

102

Isotopes of the Same Element Have

Equal numbers of protons

Different numbers of neutrons

103

Isotopic Notation

Mass number is also the numberof nucleons in the nucleus.Nucleons = protons and/or neutrons

104

Relationship Between Mass Number and Atomic

Number

Relationship Between Mass Number and Atomic

Number

105

The mass number minus the atomic number equals the number of neutrons in the nucleus.

47109 Ag

mass numberatomic number

number of neutrons

- =

62109 - 47 =

mass number

atomic number

106

Isotopic Notation

88 protons

168 protons + 8 neutrons

O

107

Isotopic Notation

88 protons

178 protons + 9 neutrons

O

108

Isotopic Notation

88 protons

188 protons + 10 neutrons

O

109

Hydrogen has three isotopes

1 proton

0 neutrons

1 proton

1 neutron

1 proton

2 neutrons

110

Examples of Isotopes

Element Protons Electrons NeutronsSymbol

Hydrogen 1 1 0

Hydrogen 1 1 1

Hydrogen 1 1 2

Uranium 92 92 143

Uranium 92 92 146

Chlorine 17 17 18

Chlorine 17 17 20

11H21H31H

23592 U

23892 U

3517 Cl3717 Cl

111

Atomic WeightAtomic Weight

112

• The mass of a single atom is too small to measure on a balance.

• Using a mass spectrometer, the mass of the hydrogen atom was determined.

113

A Modern Mass Spectrometer

A mass spectrogram is recorded.

From the intensity and positions of the lines on the mass spectrogram, the different isotopes and their relative amounts can be determined.

Positive ions formed from sample. Electrical field

at slits accelerates positive ions.

Deflection of positive ions occurs at magnetic field.

114

A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given.

115

Using a mass spectrometer, the mass of one hydrogen atom was determined to be 1.673 x 10-24 g

116

smallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmallsmall

smallsmallsmallsmall

smallsmallsmallsmall

smallsmall

This number is very small.

117

1.673 x 101.673 x 10-24-24 g g

The mass of a hydrogen atom is very small.

Numbers of this size are too small for practical use. To overcome this problem a system of relative atomic weights using “atomic mass units” was devised to express the masses of elements using simple numbers.

118

126C

The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon.

119

126C

A mass of exactly 12 atomic mass units (amu) was assigned to

120

1 amu is defined as exactly equal to the mass of a carbon-12 atom 12

1

126C

1 amu = 1.6606 x 10-24 g

121

HAverage atomic weight 1.00797 amu.

122

KAverage atomic weight 39.098 amu.

123

UAverage atomic weight 248.029 amu.

124

Average RelativeAtomic Weight

Average RelativeAtomic Weight

125

• Most elements occur as mixtures of isotopes.

• Isotopes of the same element have different masses.

• The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12 (exactly 12.0000…amu)

126

To calculate the atomic mass multiply the atomic mass of each isotope by its percent abundance and add the results.

(62.9298 amu) 0.6909 = 43.48 amu

(64.9278 amu) 0.3091 = 20.07 amu63.55 amu

IsotopeIsotopic mass

(amu)Abundance

(%)

Average atomic mass

(amu)

62.9298 69.09

64.9278 30.91

6329 Cu6529 Cu

63.55

127

Isotope Practice(Fill-in the Blanks)

symbol atomic no

mass no # e # n # p

8 16 10

Pt 117

30P–3

53 74

48 36

34 45

40Ca+2

16 28O

8 8

78 195 78 7819578Pt

15 30 18 15 1530 315P

12753I 127 53 53

8436Kr 36 84 36

7934Se 34 79 34

40 220Ca 20 40 18 20 20

128