Download Chapter 2 Atoms and the Atomic Theory

Key Concept Summary
Law of mass
conservation
Law of definite
proportion
Law of multiple
proportion
explained by
Atomic
structure
have
with same
number of
protons are
Atom
comprising
with different
numbers of
neutrons are
Element
represented by
lose/ gain
electrons
Nucleus
containing
Atomic
number
Electrons
Symbols
Isotopes
represented
average mass is
by
Ions
arranged in the
are
Protons
Neutrons
Cations
Anions
Periods
having
divided into
Groups
Atomic
mass
Element
symbol,
Periodic
table
Classes
Atomic
number, Z
Mass
number, A
which are
which are
Main Groups
Transition
Metal
Inner transition
metals
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 B
10.000g
27.000g
1.119g H
%H = 11.19
3.031g H
8.881g O
%O = 88.81
23.979g O
Dolton’s Atomic Theory
How 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
second 1 g carbon per 2.667 g oxygen
C:O mass ratio = 1: 1.333
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: 0.144
N:H ratio in B = 2:00: 0.144 = 1.00: 0.0720
H 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 = 1.25
mass 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 Atoms
What 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)
-ray
-particle
4.00
-ray
-particle (electron) 5.49x10-4
-ray
-ray
0
Charge
+2
-1
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 Atoms
Therefore
•Modern picture of an atom, then, consist of three types of particles-electrons,
protons and neutron.
Electric Charge
Particle SI (C )
Atomic
Electron -1.602x10-19
-1
Proton +1.602x10-19 +1
Neutron
0
0
Mass
SI (g)
9.109x10-28
1.673x10-24
1.675x10-24
amu
5.49x10-4
1.0073
1.0087
Located
outside nucleus
in nucleus
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 number
What 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 the
same 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 isotopes
http://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.
number p + number n
number p
A
E
Z
Symbol of element
E.g.
Mass #
Atomic #
12
C
6
Symbol of element
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.
number p + number n
number p
A +?
E
Z
number p - number e
E.g. Determine numbers of electrons in Mg2+ cation and the S2- anion?
Mg2+ 12-number e = +2
S216-number e = -2
number e =10
number e =18
Isotopic Masses
we 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 carbon12 isotope.
Mass of one 12C atom = 12 amu (exactly)
1 amu = Mass of one 12C atom = 1.660539 x 10-24 g
12
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 = 1.33291
mass 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 Mass
is the average of the isotopic masses, weighted according to the
naturally occurring abundances of the isotopes of the element.
atomic mass
natural
of an
= abundance
element
of isotope(1)
x
mass of
isotope (1)
natural
+ abundance
of isotope(2)
mass of
x isotope (2)
E.g. Chlorine consists of the following isotopes:
Isotope
isotope mass(amu) fractional abundance
Chlorine-35 34.96885
0.75771
Chlorine-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 elements
1869 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
Period
1A
1
2
1
H
He
Li
Be
6.941
9.012
Mg
22.99
24.30
7
8
9
10
11
12
13
14
15
3A
4A
5A
5
6
7
16
17
18
6A
7A
8
9
4.003
10
F
Ne
10.81 12.01 14.01 16.00 19.00 20.18
B
13
3B
20
21
4B
5B
6B
16
S
Cl
Ar
26.98
28.09
30.97
32.07
35.45
39.95
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
36
Kr
78.96 79.90 83.80
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
52
53
I
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
132.9
137.3
87
88
Fr
Ra
(223)
226.0
*
**
51
35
Br
39
56
50
34
Se
38
Ba
49
33
18
Sr
55
48
32
17
Rb
Cs
47
31
P
Cr
46
30
15
V
45
29
14
23
44
28
Si
22
43
27
2B
Al
Ti
42
26
1B
O
Sc
41
25
--- --8B -- ---
N
Ca
40
24
7B
C
K
37
6
7
12
Na
19
5
6
4
11
4
5
2A
3
3
4
8A
1.008
2
3
54
71
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Lu
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
105
106
107
108
109
110
111
112
113
114
115
116
117
118
138.91 178.5
103
104
Rf
Db
Sg
Bh
Hs
Mt
227.0
(216)
(262)
(266)
(264)
(269)
(268)
Lr
57
*Lanthanoids
*
La
**Actinoids
**
Ac
89
58
Ce
90
Th
59
60
Pr
Nd
91
92
Pa
U
61
Pm
93
Np
62
Sm
94
Pu
63
Eu
95
Am
Ds
64
Gd
96
Cm
Rg
65
Tb
97
Bk
Uub
66
Dy
98
Cf
Uut
67
Ho
99
Es
Uuq
68
Er
100
Fm
Uup
69
Tm
101
Md
Uuh
70
Yb
102
No
Uus
Uuo
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 groups
Main 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 Constant
definition: 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= 15.9994g
The 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