GC-S007-Atom

Henry R. Kang (1/2010)

General Chemistry

Lecture 7

Atom


Contents

• Atomic Theory

• Law of Mass Conservation

• Structure of Atom

• Atomic Number

• Mass Number

• Isotopes

• Periodic Table


Atomic Theoryof

Matter


Brief History of Atomic Theory

• Greek philosopher Democritus (460-370 B.C.) expressed the belief that all matter consists of tiny indivisible particles, which he named “atomos” (meaning indivisible). However, this view of matter was not a mainstream philosophy.

• Plato (427?-347 B.C.) and Aristotle (384-322 B.C.) believed that there can be no ultimately indivisible particles. Therefore, the “atomic” view of matter faded for several

millenniums.

• The modern era of atomic theory started from the work of John Dalton (1766-1844).


Dalton’s Atomic Theory (1808)• All matter is composed of indivisible atoms.

No longer true because atoms can be split into subatomic particles.

• Elements are composed of extremely small particles called atoms. All atoms of a given element are identical, having the same size, mass and

chemical properties. No longer true because of isotopes.

The atoms of one element are different from the atoms of other elements.

• Compounds are composed of atoms of more than one element. The ratios of the numbers of atoms from all elements in a compound are

integers or simple fractions (Law of Definite Proportions).

• Chemical reactions only involve the separation, combination, or rearrangement of atoms. Atoms are not created or destroyed in chemical reactions (Law of mass

conservation).


Atomic Symbols and Models

• Atom is represented by one- or two-letter taken from its name. The first letter is capitalized from the name of the element.

H (hydrogen), C (carbon), O (oxygen), N (nitrogen), etc.

Sometimes, the first letter is followed by an additional letter from the name to distinguish elements having the same first capital letter. Cl (chlorine), Ca (calcium), and Cu (copper from cuprum) Ne (neon), Na (sodium), and Ni (nickel)

• Models Dalton used spheres of different sizes to represent atoms and

combinations of theses spheres to represent compounds. Dalton’s approach has been evolved to three-dimensional models.


3Li

lithium

37

19K

potassium

87

55

20Ca

calcium

10Ne

neon

3231 3635Br

bromine

3433As

arsenic

15P

phosphorus

16S

sulfur

17Cl

chlorine

18

42

14Si

silicon

38

89

41

57

4039

88

56Ba

barium

49 50Sntin

51 52 53I

iodine

54

105

73

104

72 74W

tungsten

106

43 44 45 46 47Ag

silver

48

109

76 79Augold

108107

83Bi

bismuth

84 85 8675 80Hg

mercury

78 8177 82Pblead

11A

22A

55B

44B

66B

33B

88B

77B

98B

108B

111B

122B

133A

144A

155A

166A

177A

188A

24Cr

chromium

232221

12Mg

magnesium

4Be

beryllium

1H

hydrogen

9F

fluorine

2He

helium

30Znzinc

29Cu

copper

28Ni

nickel

27Co

cobalt

26Feiron

25Mn

manganese

5B

boron

13Al

aluminum

6C

carbon

7N

nitrogen

8O

oxygen

110 111 114112 118116

11Na

sodium

Symbols & Names of Common Elements


Law ofConservation of Mass


Law of Conservation of Mass

• The total mass remains the same during a chemical change (or chemical reaction). Proposed by A. Lavoisier (1743-1794)

• Mass of reactants = mass of products Reactants are the original matters

before the chemical change. Products are the matters formed after

the chemical change.


Law of Conservation of Mass: Illustration

• Atoms are not created or destroyed in chemical reactions (Law of conservation of mass)

8 X2Y16 X 8 Y+

Atoms of element X

Atoms of element Y

Compound of elements X and Y


Law of Conservation of Mass: Example• A sample of 1.28 grams magnesium is burned in air

and 2.12 grams of a white ash-like residue (magnesium oxide) is produced at the completion of the reaction. What is the mass of oxygen that reacts?

Magnesium + Oxygen Magnesium oxide

2 Mg + O2 2 MgO

• Answer: (mass of magnesium) + (mass of oxygen)

= (mass of magnesium oxide)

1.28 g + (mass of oxygen) = 2.12 g

(mass of oxygen) = 2.12 g – 1.28 g = 0.84 g


Law of Definite Proportions

• Law of definite proportions was proposed by Joseph Proust (1754-1826) in 1799. The law can be expressed in two ways:

1. Different samples of the same compound always contain elements in the same proportion by mass, regardless where they come from. Example: samples of carbon dioxide gas obtained from different sources

contain exactly the same ratio by mass of carbon to oxygen. Mass ratio = 12/32 = 3/8 = 0.375

2. The relative number of atoms of each element in a given compound is always the same. Example: samples of carbon dioxide gas obtained from different sources

contain exactly the same ratio of carbon atom to oxygen atom. C/O atom ratio = 1/2 = 0.5


Laws of Multiple Proportions• If two elements can combine to form more than one

compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.

Nitrogen monoxide

Nitrogen dioxide

Oxygen in NO and NO2 has a ratio of 1/2.

Oxygen in NO, NO2 and N2O5 has a ratio of 2/4/5.

OC

OC

11

21

= =

= =


Structureof

the Atom


Structure of Atom• Atom is the basic unit of an element that can

enter into chemical combination. However, atom is not indivisible; it can be split into

subatomic particles.

• Atom possesses internal structure that consists an inner core, nucleus, and surrounding electrons. Nucleus

Proton (p)

Neutron (n)

Electron (e–)


Subatomic Particles: Leptons & Hadrons

• There are two broad categories of the subatomic particles

• Leptons (Greek for “light” or “small”) They can be viewed as a point particle with very little size or no size at all They have no internal structure They are not affected by the strong force interaction Example:

electron, positron, neutrino, quarks

• Hadrons (Greek for “heavy” or “strong”) They have definite sizes They have internal structure They are subject to the strong force interaction Examples:

Proton, neutron, etc. Proton consists of one d quark (-1/3 e) and two u quarks (+2/3 e) Neutron consists of one u quark and two d quarks

d(-1/3)

n (0)

p (1)

Proton

Neutron

d(-1/3)

d(-1/3)

u (+2/3)

u (+2/3)

u (+2/3)


Discovery of Electron

• J.J. Thomson used cathode-ray tube to demonstrate the existence of the charged particles and measured charge/mass ratio of e– (1906 Nobel Prize in Physics)

• The cathode ray is attracted by the plate bearing positive charges and repelled by the plate bearing negative charges. Thomson concluded that cathode rays are streams of

negatively charged particles.

• Using electromagnetic theory, Thomson determined the charge (Ce) to mass (me) ratio of an electron.

Ce/me = -1.76×108 coulomb/gram


Millikan’s Experiment

• Millikan (1868-1953) used the oil drop experiment for measuring charge of e– (1923 Nobel Prize in Physics)

• He determined the value of the electronic charge by monitoring the motions of charged oil drops under an electric field. The charge on each electron is exactly the same.

Ce = -1.6022 ×10-19 coulomb

• Knowing the charge and charge/mass ratio (Thomson’s result), he calculated the mass of the electron Ce/me = -1.76×108 coulomb/gram (Thomson’s result)

me = -1.6022 ×10–19 coulomb / (-1.76×108 coulomb/gram)

= 9.10×10–28 gram


Radioactivity & Fundamental Particles

• Henri Becquerel (1852-1908) and Marie Curie (1845-1923, 1903 Nobel Prize in Physics with husband, Pierre, and Becquerel; and 1911 Nobel Prize in Chemistry) coined the name radioactivity to describe the emission of particles and radiation from some radioactive elements.

• Three types of rays are produced by the decay of radioactive elements such as uranium. Alpha (α) ray is positively charged and is identified as the

helium nuclei. Beta (β) ray is negatively charged and is identified as the

electron. gamma (γ) ray has no charge and is identified as the high

energy photons.


Thomson’s Atomic Model• Thomson proposed (in 1904) that an atom is a uniform,

positive sphere in which electrons are embedded like raisins in a cake; Therefore, this model is sometimes referred to as the “raisin pudding” model.

Positive charge spread over the entire sphere –

–

–

–

––

–

–


Geiger-Marsden-Rutherford’s Experimental Design & Observations

• Marsden under the supervision of Geiger bombarded a thin gold foil with α particles (velocity is about 1.4×107 m/s, 5% speed of light).

• Majority of α particles penetrated the foil un-deflected or with a slight deflection.

• Occasionally, a few (about 1 in 8000) α particle was scattered (or deflected) at a large angle.

• In rare instances, α particle was actually bounced backward.

• Thomson’s model contradicted to this phenomenon.


Rutherford’s Atomic Model• Based on the results from Ernest Marsden and

Hans Geiger, Rutherford (1871-1937) proposed (in 1911) that the majority of the mass and positive charges of the atom was located in a small, dense region, the nucleus, with negatively charged electrons occupying a much larger volume outside of the nucleus. The positive charge of atoms is concentrated

in the nucleus. Proton (p) has positive charge as opposite to

the negative charge of electron. Mass of proton is about 1840 times of the

electron (1.67×10-24 g). Nuclei have diameters of about 10-15 m,

whereas atomic diameters are about 10-10 m. If the nucleus is represented by a golf ball,

then the atom would be about 3 miles in diameter.

A region of mostly empty space where electrons reside

Dense, positively charged nucleus at the center

–

–

––

––

––


Comparison of Atom & Nuclear Radii

• For hydrogen atom Atomic radius = 3.1×10-11 m (31 pm)

Radius of hydrogen nucleus (proton) = 8.768×10-16 m

The ratio of atomic radius to nuclear radius is about 35,000.

• For a larger atom: Atomic radius ~ 100 pm = 1×10-10 m

Nuclear radius ~ 5×10-3 pm = 5×10-15 m

The ratio of radii is about 20,000.

• Imagine “If an atom has the size of the Houston Astrodome, then the nucleus is a marble in the center.”


Chadwick’s Experiment (1932)• After the discovery of electron and proton, a problem arose that

the mass ratio of hydrogen to helium did not add up: H atoms - 1 proton He atom - 2 protons The mass ratio of (He mass) to (H mass) should be 2. But, the measured mass ratio was 4.

• Chadwick (1891-1974) bombarded a thin sheet of beryllium with α particles, a very high-energy radiation was emitted by the metal. α + 9Be 1n + 12C + energy

• Later, it was shown that the rays were a third type of subatomic particle named “neutron” by Chadwick.

• Neutron (n) is neutral (charge = 0) with a mass slightly higher than proton. n mass ~ p mass = 1.67×10-24 g


Mass & Charge of Subatomic Particles

• (mass p) / (mass n) = 1.67262×10-24 / 1.67493×10-24

= 0.998621

• (mass p) / (mass e-) = 1.67262×10-24 / 9.10939×10-28 = 1836

• (mass n) / (mass e-) = 1.67493×10-24 / 9.10939×10-28 = 1839 They differ by about 3 electron-masses

• (mass p) (mass n) = 1840 × mass e-

Particle Mass(gram)

Charge(Coulomb)

Chargeunit

Electron 9.10939×10-28 -1.6022×10-19 -1

Proton 1.67262×10-24 +1.6022×10-19 +1

Neutron 1.67493×10-24 0 0


Nuclear Structureand

Isotopes


Nuclear Structure

• Nucleus consists of protons and neutrons. Except the hydrogen nucleus

Hydrogen nucleus has only one proton and no neutron.

• For any neutral atom, the number of proton equals the number of electrons. #proton = #electron

• The number of protons is called “atomic number”.


Atomic Number, Mass Number & Isotopes• All atoms can be identified by two numbers

The number of protons and the number of neutrons

• Atomic Number (Z) The number of protons in the nucleus: Z = #protons

In a neutral atom, the number of protons is equal to the number of electrons.

• Mass Number (A) The total number of neutrons and protons in the nucleus A = #protons + #neutrons

• Nuclide A nuclide is an atom characterized by atomic number and mass number,

represented by a symbol, AZX. Example: 32

16S

• Isotopes Atoms have the same atomic number but different mass numbers (or

different number of neutrons).


Isotopes: Definition & Examples

• Isotopes are the same element (same number of protons) with different numbers of neutrons in their nuclei. Atoms have the same atomic number but different mass

numbers.

1 proton

H11 H (D)2

1 H (T)31

1 proton1 neutron

1 proton2 neutron

Hydrogen Deuterium Tritium

U23592 U238

92

XAZ

Uranium-235 Uranium-238

Element SymbolMass Number

Atomic Number

11H

21H 3

1H


Isotopes - Computation

• How many protons, neutrons, and electrons are in 14

6C?6 protons, 8 (14 - 6) neutrons, and

6 electrons• How many protons, neutrons, and

electrons are in 116C?

6 protons, 5 (11 - 6) neutrons, and6 electrons


Examples of Nuclide Symbol

• Give the number of protons, neutrons, and electrons in the following atoms: Number of protons = Number of electrons = Atomic number Number of neutrons = mass number – atomic number 17

8O Number of protons = Number of electrons = 8

Number of neutrons = 17 – 8 = 9 199

80Hg Number of protons = Number of electrons = 80


80Hg Number of protons = Number of electrons = 80


29Cu Number of protons = Number of electrons = 29

Number of neutrons = 63 – 29 = 34


Periodic Tableof

Elements


Periodic Table• Many elements show strong similarities to one another.

They process periodic regularities in physical and chemical properties.

• Periodic table is a chart in which elements, having similar chemical and physical properties, are group together. Horizontal rows are called period. Vertical columns are called group or family.

• Three categories Metal

Elements are good conductor of heat and electricity Nonmetal (17 elements)

Usually, poor conductor of heat and electricity. Metalloid (8 elements)

Intermediate elements between metals and nonmetals From left to right across any period, the properties of the elements change

gradually from metallic to nonmetallic.


Modern Periodic Table1

1A188A

1H

1.008

22A

133A

144A

155A

166A

177A

2He

4.003

3Li

6.941

4Be

9.012

5B

10.81

6C

12/01

7N

14.01

8O

16.00

9F

19.00

10Ne

20.18

11Na

22.99

12Mg

24.31

33B

44B

55B

66B

77B

88B

98B

108B

111B

1212B

13Al

26.98

14Si

28.09

15P

30.97

16S

32.07

17Cl

35.45

18Ar

39.95

19K

39.10

20Ca

40.08

21Sc

44.96

22Ti

47.88

23V

50.94

24Cr

52.00

25Mn

54.94

26Fe

55.85

27Co

58.93

28Ni

58.69

29Cu

63.55

30Zn

65.39

31Ga

69.72

32Ge

72.59

33As

74.92

34Se

78.96

35Br

79.90

36Kr

83.80

37Rb

85.47

38Sr

87.62

39Y

88.91

40Zr

91.22

41Nb

92.91

42Mo

95.94

43Tc

(98)

44Ru

101.1

45Rh

102.9

46Pd

106.4

47Ag

107.9

48Cd

112.4

49In

114.8

50Sn

118.7

51Sb

121.8

52Te

127.6

53I

126.9

54Xe

131.3

55Cs

132.9

56Ba

137.3

57La

138.9

72Hf

178.5

73Ta

180.9

74W

183.9

75Re

186.2

76Os

190.2

77Ir

192.2

78Pt

195.1

79Au

197.0

80Hg

200.5

81Tl

204.4

82Pb

207.2

83Bi

208.9

84Po

(209)

85At

(210)

86Rn

(222)

87Fr

(223)

88Ra

(226)

89Ac

(227)

104Rf

(257)

105Db

(260)

106Sg

(263)

107Bh

(262)

108Hs

(265)

109Mt

(266)

110Ds

(271)

111Uuu(272)

112Uub(277)

114Uuq(296)

116Uuh(298)

118Uuo(?)

58Ce

140.1

59Pr

140.9

60Nd

144.2

61Pm

(147)

62Sm

(150.4)

63Eu

152.0

64Gd

157.3

65Tb

158.9

66Dy

162.5

67Ho

164.9

68Er

167.3

69Tm

168.9

70Yb

173.0

71Lu

175.0

90Th

232.0

91Pa

(231)

92U

(238)

93Np

(237)

94Pu

(242)

95Am

(243)

96Cm

(247)

97Bk

(247)

98Cf

(249)

99Es

(254)

100Fm

(253)

101Md

(256)

102No

(254)

103Lr

(257)

Lanthanides

Actinides

Metals

Nonmetals

Metalloids

Halogen

Nob

le Gas

Alk

ali Metal

Alk

aline E

arth

Metal

Period

Grou

p

Note that the table is organized in order of the atomic number.


Distribution of Elements on Earth & Body• Natural abundance of elements in earth’s crust

Oxygen: 45.5% Silicon: 27.2% Aluminum: 8.3% Iron: 6.2% Calcium: 4.7% Magnesium: 2.8% All others: 5.3%

• Natural abundance of elements in human body Oxygen: 65% Carbon: 18% Hydrogen: 10% Nitrogen: 3% Calcium: 1.6% Phosphorus: 1.2% All others: 1.2%

GC-S007-Atom

Documents

indivisible atoms

atomic theory of matter

b boron

rearrangement of atoms

numbers of atoms

atomic symbols

daltons atomic theory

atomic view of matter