Key Concept SummaryLaw of mass conservation

Law of definiteproportion

Law of multipleproportion

AtomAtomicstructure Element

Nucleus Electrons

Protons Neutrons

Ions

Cations Anions

Isotopes

Elementsymbol,

Atomicmass

Atomicnumber, Z

Massnumber, A

explained by

havewith samenumber ofprotons are

with different numbers ofneutrons are

Atomic number

Periodic table

Symbols

containing

comprising

represented by

average mass is

having

lose/ gainelectrons

are

represented by

arranged in the

divided into

Periods Groups Classes

which arewhich are

Transition Metal

Inner transitionmetalsMain Groups

Metals,Semimetals,or Nonmetals

Chapter 2 Atoms and the Atomic Theory 

All the matters can be broken down into elements. Is matter

continuously divisible into ever smaller and smaller pieces, or is there an

ultimate limit? What is an element made of?

Although most of thinkers including Plato and Aristotle, believed that

matter is continuous, the Greek philosopher Democritus disagreed.

Democritus proposed the elements are composed of tiny particles that we

now call atoms, from the Greek word atomos, meaning indivisible.

I. Early Chemical Discoveries and Atomic Theory

Three important fundamental laws in chemistry

1) Law of conservation of mass

2) Law of constant composition

3) Law of multiple proportions

Law of conservation of mass

1774 Antoine Lavoisier –showed heating the red power HgO causes it to decompose into the silvery liquid mercury and the colorless gas oxygen. 2HgO 2Hg +O2 then show that oxygen is the key substance involved in

combustion.

Furthermore, he demonstrated by careful measurements that when combustion is carried out in a closed container, the mass of the combustion products is exactly equal to the mass of the starting reactants. (tin + air+ sealed glassed vessel) (tin oxide + remaining air + glass

vessel)

Law of conservation of mass~ The total mass of substances present after a chemical reaction is the same as the total mass of substances before the reaction. Matter is neither created nor destroyed in a chemical reaction.

E.g. A 0.382g sample of magnesium reacts with 2.652g of nitrogen gas. The sole product is magnesium nitride. After reaction, the mass of unreacted nitrogen is 2.505g. What mass of magnesium nitride is produced?

mass before reaction = 0.382g magnesium + 2.652g nitrogen gas = 3.034g

mass after reaction = ?g magnesium nitride + 2.505 nitrogen gas 3.034g =?g magnesium nitride + 2.505 nitrogen gas ?g magnesium nitride = 3.034g – 2.505g = 0.529g

quiz 1.doc

Law of constant composition

1799 Joseph Proust – reported one hundred pounds of copper, dissolved in sulfuric or nitric acids and precipitated by the carbonates of soda or potash, invariably gives 180 pounds of green carbonate.

Law of constant composition~ All samples of a compound have the same composition- the same proportion by mass of the constituent elements. This means that the relative amount of each element in a particular compound is always the same, regardless of the source of the compound or how the compound is prepared.

E. g. Water is made up of two elements H and O. The two sample of water below have the same proportions of the two elements, expressed as percentages by mass. Every sample of water contains 1 part hydrogen and 8 parts oxygen by mass. _________________________________________

sample A Composition sample B10.000g 27.000g1.119g H %H = 11.19 3.031g H8.881g O %O = 88.81 23.979g O

Dolton’s Atomic TheoryHow can the Law of conservation of mass and Law of constant composition be explain? Why do element behave as they do?To answer these questions:

1803-1808 John Dalton : proposed a new theory of matter. 1 Each chemical element is composed of minute, indestructible particles called atoms. Atoms can be neither created nor destroyed during a chemical change. (If atoms of an element are indestructible, then the same remains unchanged. This explains the law of conservation of mass).

2 All atoms of an element are alike in mass and other properties, but the atoms of one element are different from those of all other elements.

3 In each of their compounds, different elements combine in a simple numerical ratio: e.g. one atom of A to one of B (AB) or one atom of A to two of B (AB2). If all atoms of an element are alike in mass (assumption 2)

and if atoms unite in fixed numerical ratio (assumption 3), the percent composition of a compound must have a unique value, regardless of the origin of the sample analyzed. This explains the law of constant composition.

Law of multiple proportions

Dalton’s theory leads to a prediction- the law of multiple proportions.~ If two elements form more than a single compound, the masses of one element combined with a fixed mass of the second are in the ratio of small whole numbers. Same elements to combine in different ratios to give different substances.

E.g. Oxygen and carbon can combine either in a 1: 1.333 mass ratio to make a substance or in a 1: 2.667 mass ratio to make a substance.

first 1 g carbon per 1.333 g oxygen C:O mass ratio = 1: 1.333second 1 g carbon per 2.667 g oxygen C:O mass ratio = 1: 2.667

comparison C:O mass ratio in first sample = (1 g C)/(1.333g O) = 2

of C:O ratios C:O mass ratio in second sample (1 g C)/(2.667g O)

Compare two substances clearly the second substance contains exactly twice as

much oxygen as the first for a given number of carbon. If the first oxide has the

molecular formula CO then the second oxide will be CO2.

There are two compound both contain nitrogen and hydrogen. Compound A contains 1.50g of N and 0.216g H. Compound B contains 2.00g of N and 0.144g H. If the formula of compound B is N2H2, what is the formula

of compound A?

N:H ratio in A = 1.50: 0.216g = 1.00: 0.144N:H ratio in B = 2:00: 0.144 = 1.00: 0.0720H in A is (0.144/0.0720 = 2 ) twice as much in B

IF B is N2H2 then A is N2H4

Atomic Mass Unit

Dalton’s theory enables us to set up a scale of relative atomic masses. He cannot measure the exact mass of atoms but relative mass.E.g. Consider calcium sulfide, which consists of 55.6% calcium by mass and 44.4% sulfur by mass. Suppose there is one calcium atom for each sulfur atom in calcium sulfide. Because we know that the mass of a calcium atom relative to that of a sulfur atom must be the same as the mass % in calcium, we know that the ratio of the mass of a calcium atom to that of a sulfur atom is

mass of Ca atom = 55.6 = 1.25mass of a S atom 44.4 or mass of a Ca atom = 1.25 x mass of a sulfur atom

By continuing in this manner with other compounds, it is possible to build up a table of relative atomic masses. We define a quantity called atomic mass ratio (usually referred to as atomic mass), which is the ratio of the mass of a given atom to the mass of some particular reference atom. The unit of atomic mass is assigned atomic mass unit (amu).

The Structure of AtomsWhat is an atom made of ?

Discovery of subatomic particle

i) The Discovery of Electrons

*1897 J.J. Thomson- cathode ray experiment

Thomson’s experiment involved the use of cathode-ray tube. When a

sufficiently high voltage is applied across the electrode, an electric current

flows through the tube from negatively charged electrode ( the cathode) to the

positively charged electrode (the anode).

Later experiments had shown that cathode-ray can be deflected by electric or

magnetic field. Because the beam is produced at a negative electrode and is

deflected toward a positive plate, Thomson proposed that cathode rays are

negatively charged fundamental particles found in all atoms which, we now

called electrons.

Furthermore, because electrons are emitted from electrodes made of many

different metals, all these substances must contain electrons. By careful

measuring the amount of deflection caused by electric and magnetic fields of

known strength, He established the ratio of mass to electric charge for cathode ray

that is, m/e = -5.6857x10-9 g/coulomb.

Millikan’s Oil-Drop Experiment: Mass of Electron

1909 Robert Millikan determined the electronic charge through a series of oil-drop experiments. The currently accepted value of the charge of the e is –1.6022x10-19C. Substituting into Thomson’s mass to charge ratio then gives the mass of electron as 1/1836(= 9.1094x10-28g).

X-Ray and Radioactivity

Ernest Rutherford identified two type of radiation from radioactive materials, alpha () and beta (). -particles (He2+)carry two fundamental units of positive charge and have essentially the same mass as He atoms. -particles are negatively charged particles produced by changes occurring within the nuclei of radioactive atoms and have the same properties as electrons.

A third form of radiation, that is not affected by an electric field was discovered in 1900 by Paul Villard. This radiation, called -ray, is not made up of particles; it is electromagnetic radiation of extremely high penetrating power. Properties of the three radioactive emissions discovered Original name Modern name Mass (amu) Charge-ray -particle 4.00 +2-ray -particle (electron) 5.49x10-4 -1-ray -ray 0 0_______

Rutherford’s Scattering Experiment

1909 Ernest Rutherford~ use particle to study the inner structure of atoms. When he directed a beam of -particles at a thin gold foil, he found that

The majority of -particles penetrated the foil undeflected. Some particles experienced slightly deflections. A few (about one in every 20,000) suffered rather serious deflections as they penetrated the foil. A similar number did not pass through the foil at all, but bounced back in the direction from which they had come.

The Nuclear Atom: Protons and Neutrons

1911 Rutherford explained his results by proposing a model of the atom known as the nuclear atom and having these features. 1 Most of the mass and all of the positive charge of an atom are centered in a very small region called the nucleus. The atom is mostly empty space.

2 The magnitude of the positive charge is different for different atoms and is approximately one-half the atomic weight of the element.

3 There are as many electrons outside the nucleus as there are units of positive charge on the nucleus. The atom as a whole is electrically neutral.  Rutherford’s nuclear atom suggested the existence of positively charged fundamental particles of matter in the nuclei of atoms- called protons. He predicted the existence in the nucleus of electrically neutral particles.

* 1932 Jame Chadwick~ verified that there is another type of particles in atom called neutron.

The Structure of AtomsTherefore•Modern picture of an atom, then, consist of three types of particles-electrons, protons and neutron.

Electric Charge Mass Particle SI (C ) Atomic SI (g) amu LocatedElectron -1.602x10-19 -1 9.109x10-28 5.49x10-4 outside nucleusProton +1.602x10-19 +1 1.673x10-24 1.0073 in nucleusNeutron 0 0 1.675x10-24 1.0087 in nucleus

Conclusion:

• Modern physics has revealed successively deeper layers of structure in ordinary matter. Matter is composed, on a tiny scale, of particles called atoms. Atoms are in turn made up of minuscule nuclei surrounded by a cloud of particles called electrons. Nuclei are composed of particles called protons and neutrons, which are themselves made up of even smaller particles called quarks. Quarks are believed to be fundamental, meaning that they cannot be broken up into smaller particles.

Chemical Elements

i) Atomic numberWhat is that makes one atom different from another?Elements differ from one another according to the number of protons in their atoms, a value called the element’s atomic number. All atoms of particular element have the same atomic number, Z, on the other hand all atoms with thesame number of protons are the same element. • atomic number (Z) = Number of proton in atom’s nucleus• mass number (A) = # of protons (Z) + # of neutrons (N)

Isotopes

Contrary to what Dalton thought, we know that atoms of an element do not necessarily all have the same mass. In 1912, J. J. Thomson measured the mass-to charge ratios of positive ions formed in neon gas. He found that about 91% of the atoms had one mass and that the remaining atoms were about 10% heavier. All neon atoms have 10 protons in their nuclei, and most have 10 neutrons as well. A few neon atoms, however have 11 neutrons and some have 12.

The naturally occurring percentages of isotopes of a particular element are referred to as the percent natural abundance of that element. Atoms that have the same atomic # (Z) but different mass numbers (A) are called isotopes.Most elements occur naturally as a mixture of different isotopes.

E.g. Hydrogen has three isotopeshttp://www.prenhall.com/petrucci

An isotope is specified by its atomic # and its mass #. The notation used to designate isotopes is the chemical symbol of the element written with its atomic # as a left subscript and its mass # as a left superscript.

EAZ

Symbol of element

number p + number n

number p

C12

6

Mass #Symbol of element

Atomic #

E.g.

Ions

~ is an electrically charged particle obtained from an atom or chemically bonded group of atoms lose or gain electrons. The charge on an ion is equal to the # of protons minus the # of electrons. An atom that gains extra electrons becomes a negatively charged ion, called an anion. An atom that loses electrons becomes positively charged ion, called a cation.

EAZ

number p + number n

number p

+ ? number p - number e

E.g. Determine numbers of electrons in Mg2+ cation and the S2- anion?

Mg2+ 12-number e = +2 number e =10

S2- 16-number e = -2 number e =18

Isotopic Masseswe arbitrarily choose one atom and assign it a certain mass. By international agreement, this standard is an atom of the isotope carbon-12, which is assigned a mass of exactly 12 atomic mass unit. Next, we determine the mass of other atoms relative to carbon-12 by mass spectrometer. All atomic masses are given relative to the mass of carbon-12 isotope.

Mass of one 12C atom = 12 amu (exactly)

1 amu = Mass of one 12C atom = 1.660539 x 10-24 g12

E.g From the mass spectral data, the ratio of the mass of 16O to 12C is found to be 1.33291. What is the mass of an 16O atom?

16O/12C = 1.33291mass of 16O = 1.33291 x 12amu = 15.9949amu

E.g. If the standard is H =1.00amu then what would be the atomic mass of carbon be?

Atomic Mass

v) Atomic Massis the average of the isotopic masses, weighted according to the naturally occurring abundances of the isotopes of the element.

atomic mass natural mass of natural mass of of an = abundance x isotope (1) + abundance x isotope (2) + .....element of isotope(1) of isotope(2)

E.g. Chlorine consists of the following isotopes: Isotope isotope mass(amu) fractional abundanceChlorine-35 34.96885 0.75771Chlorine-37 36.96590 0.24229

What is the atomic mass (weight) of chlorine?

• The average atomic mass of an imaginary element, Z, is 12.50 amu. Z has only two stable isotopes: 10Z (10.20 amu) and 13Z (13.00 amu). What is the natural percent abundance of 13Z?

Introduction to Periodic Table

With discovery of many elements1869 Mendeleev and Meyer ~ independently proposed periodic table organized the elements In modern periodic table, The periodic table of the elements is organized into 18 groups and 7 periods. Elements are represented by one or two-letter symbols and are arranged according to atomic number.

* a horizontal row of elements- a period

* a vertical row of elements- a group or family

 

 

Periodic Table of Elements

Group 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

Period 1A 8A

1 1 H

1.008

2A 3A 4A 5A 6A 7A

2 He

4.003

2 3 Li

6.941

4 Be

9.012

5 B

10.81

6 C

12.01

7 N

14.01

8 O

16.00

9 F

19.00

10 Ne

20.18

3 11 Na

22.99

12 Mg

24.30

3B 4B 5B 6B 7B --- --8B -- --- 1B 2B

13 Al

26.98

14 Si

28.09

15 P

30.97

16 S

32.07

17 Cl

35.45

18 Ar

39.95

4 19 K

39.10

20 Ca

40.08

21 Sc

44.96

22 Ti

47.88

23 V

50.94

24 Cr

52.00

25 Mn

54.94

26 Fe

55.85

27 Co

58.93

28 Ni

58.69

29 Cu

63.55

30 Zn

65.39

31 Ga

69.72

32 Ge

72.61

33 As

74.92

34 Se

78.96

35 Br

79.90

36 Kr

83.80

5 37 Rb

85.47

38 Sr

87.62

39 Y

88.91

40 Zr

91.22

41 Nb

92.91

42 Mo

95.94

43 Tc (98)

44 Ru

101.1

45 Rh

102.9

46 Pd

106.4

47 Ag

107.9

48 Cd

112.4

49 I n

114.8

50 Sn

118.7

51 Sb

121.8

52 Te

127.6

53 I

126.9

54 Xe

131.3

6 55 Cs

132.9

56 Ba

137.3 *

71 Lu

138.91

72 Hf

178.5

73 Ta

180.9

74 W

183.8

75 Re

186.2

76 Os

190.2

77 I r

192.2

78 Pt

195.1

79 Au

197.0

80 Hg

200.6

81 Tl

204.4

82 Pb

207.2

83 Bi

209.0

84 Po

(209)

85 At

(210)

86 Rn

(222)

7 87 Fr

(223)

88 Ra

226.0 **

103 Lr

227.0

104 Rf

(216)

105 Db

(262)

106 Sg

(266)

107 Bh

(264)

108 Hs

(269)

109 Mt

(268)

110 Ds

111 Rg

112 Uub

113 Uut

114 Uuq

115 Uup

116 Uuh

117 Uus

118 Uuo

* Lanthanoids * 57 La

58 Ce

59 Pr

60 Nd

61 Pm

62 Sm

63 Eu

64 Gd

65 Tb

66 Dy

67 Ho

68 Er

69 Tm

70 Yb

* * Actinoids ** 89 Ac

90 Th

91 Pa

92 U

93 Np

94 Pu

95 Am

96 Cm

97 Bk

98 Cf

99 Es

100 Fm

101 Md

102 No

It is customary also to divide the elements into two broad categories known as Metals: Except mercury (liquid), metals are solid s at room temperature. They are generally malleable, ductile, good conductors of heat and electricity, and have a lustrous or shiny appearance. Nonmetals: generally have opposite properties of metals; e.g. poor conductors of heat and electricity.Metalloid (semimetal): is an element having both metallic and nonmetallic properties.Or into three groupsMain group elements are those in groups 1, 2 and 13-18. when form ions, group 1, 2 lose the same # e as their group #; group 13 lose group #-10; group 14-18 gain 18-group #.Transition elements: from group 3 to 12, and because all of them are metals, they are also called the transition metals. The # of electrons lost in TM is not related to their group #.

Inner transition metals which include Lanthanides and Actinindes.

Mole : S.I. Unit for amount of substance

We buy a quantity of groceries in several ways. by number such as eggs, apples and oranges by mass such as rice and peanuts

The quantity of an element or a compound, which like grocery items, can be measured by number or by mass. Molecules and atoms are extremely small objects - both in size and mass. Consequently, working with them in the laboratory requires a large collection of them.

How large does this collection need to be?

Chemists have adopted the mole concept as a convenient way to

deal with the enormous numbers of atoms, molecules or ions in

the sample they work with.

Quick-Think• 1) Complete a sentence starter:

– One mole of substance contains _______ number of elementary entities such as people, atoms, molecules, etc.

• 2) Select the best response: For one gram of each element, which one has least number of atoms– (a) Li– (b) Al– (c) Ti– (d) W

• 3) Correct the error: Avogadro’s number is larger than one mole

The Concept of the Mole and the Avogadro Constantdefinition: A mole is an amount of substance that contains the same number of elementary entities (atoms, molecules or formula units) as the number of atoms in exactly 12g of carbon-12. (the quantity of a substance whose mass in gram is numerically equal to the formula mass of the substance). The value of Avogadro’s number (the # of elementary entities) is based on a definition and a measurement. A mole of carbon-12 is defined to be 12 gram. The mass of one carbon-12 atom is measured using mass spectrometer and found to be 1.9927x10-23g. The ratio of this two mass is Avogadro’s number, NA.

NA =12g/mol

1.99927x10 23g= 6.02214 x1023 mol 1

Therefore, one mole of C = 6.02214 x 10 23 atoms = 12.011g one mole of O = 6.02214 x 10 23 atoms= 15.9994gThe mass of one mole of atoms, called molar mass (g/mole).

E.g. How many atoms of gold are present in 5.07x10-3mol Au?

E.g. How many lead-206 atoms are present in 1.71g of lead? The nature abundance of lead-206 is 24.1%.

First convert the mass to mol