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Dalton’s Atomic Theory (1808)
1. All matter is made of tiny indivisible particles
called atoms.
2. Atoms of the same element are identical.
The atoms of any one element are different
from those of any other element.
3. Atoms of different elements can combine
with one another in simple whole number
ratios to form compounds.
4. Chemical reactions occur when atoms are
separated, joined, or rearranged;however,
atoms of one element are not changed into
V.Montgomery
R.Smith
atoms of another
by a& chemical
reaction.
1
Summary of the Atom
• atoms are the smallest particles that can be
uniquely associated with an element
• each element has unique atoms
• atoms are composed of e-, p and n
• atoms are electrically neutral (# of e- = # of p)
• for a single element, isotopes differ only in
number of n (neutrons)
• atoms have characteristic masses (atomic
weights)
• atoms combine with one another in definite,
whole number proportions to make compounds
~ 10-10 m
electron
Mass 9 x 10-31 kg
nucleus
Mass > 10-26 kg
~ 1 – 7 x 10-15 m (1 – 7 fermi)
The Spacious Atom
atoms are dominantly empty space:
electron orbits
If an oxygen atom
had a total radius of
100 km, the nucleus
would be a ~1 m
diameter sphere in
the middle.
Electrons in Orbit
In a simplistic model, electrons float around the nucleus in
energy levels called shells.
electron orbits
As the number of
electrons increases,
they start to fill shells
farther out from the
nucleus.
In most cases,
electrons are lost or
gained only from the
outermost shell.
Atom
Nucleus
The Nuclear Model
of the atom
• Mass of an electron is approximately
1/1840th of a proton or neutron.
• Mass of a neutron is very close to the mass
of a proton.
• 1 atomic mass unit (amu or u)=
1.66054x10-24g
• 1 g= 6.022 x1023amu
• 1amu is defined as 1/12th the mass of
an atom of carbon-12 .
Subatomic Particles
Particle
Symbol
Charge
Relative
Mass
Electron
e-
1-
0
Proton
p+
1+
1
Neutron
n
0
1
Atomic Number(nuclear charge)
Atomic Number
Symbol
11
Na
All atoms of an element have
the same number of protons
11 protons
Sodium
11
Na
Number of Electrons
 If an atom is neutral ;
 The net charge is zero
 Number of protons = Number of electrons
 Atomic number = Number of electrons in a
neutral atom
Ions
• Have a net electrical charge since the total
number of electrons isn’t equal to the
number of protons.
• Can be anions (e->p+ ;formed as a result of
gaining electrons; negatively charged),
cations (p+>e-;formed as a result of losing
electrons; positively charged).
• e- + q(charge of the ion) = proton number
• In chemical rxns, atom never gain or
lose protons. It’s the interaction of
electrons.
Mass Number(nucleon
number)
Counts the number
of
protons and neutrons
in an atom
A= p+ + n0
Atomic Symbols
 Show the mass number and atomic number
 Give the symbol of the element
(A)mass number
23Na
11
(Z)atomic number
sodium-23
Notation for Atoms
12C
only one isotope of carbon
13C
only one isotope of carbon
C
all isotopes of
carbon
Basic Definitions
• “atomic number” = number of protons in the
nucleus;
• “mass number” = sum of protons + neutrons
in the nucleus
• “isotopic mass” = mass of a single isotope
More Atomic Symbols
16
O
31
P
65
8
15
30
15 p+
16 n
15 e-
30 p+
35 n
30 e-
8 p+
8n
8 e-
Zn
Isotopes
 Atoms with the same number of protons,
but different numbers of neutrons.
 Atoms of the same element (same atomic
number) with different mass numbers
Isotopes of chlorine
35Cl
37Cl
17
17
chlorine - 35
chlorine - 37
• Since both isotopes have the same
number of protons and electrons, they
have identical chemical properties.
• Since physical properties depend on
the mass of particles as well, isotopes
will often have slightly different
physical properties such as density,
mass, rate of diffusion etc.
Natural abundances of
isotopes
Natural chlorine contains :
75 %
35Cl
17
and 25 %
37Cl
.
17
• These percentages are known as the
natural abundances of the isotopes and
determined by “mass spectrometry.”
Relative atomic mass
• is weighted mean.
(75x 35)  (25x 37)
 35.5
100
Relative atomic mass has no unit since the
atomic mass unit, amu is cancelled out in
calculation w/ respect to C-12.
Exercise 1
• The molar mass of iridium is 192.2 g / mol.
What are the naturally occurring
percentages of the two isotopes of Ir-191
and Ir-193?
solution
191( x )  193(100  x )
 192.2
100
x  40
Iridium is a mixture of 40% 191Ir and
60 % 193Ir.
Isotopes of Hydrogen element
3
1
H
2
1
H
1
1
H
has the biggest abundance in nature.
Radioactive isotopes
• Are produced by exposing the natural
element to a flux(flow) of neutrons in a
nuclear reactor. The nucleus of an atom
captures an additional neutron and form
radioisotope.
Usages of radioisotopes
1. The rate of radioactive decay is used to
date objects (C-14).
 Naturally occurring C has a fixed
proportion of C-14 due to exchange w/ C in
the atmosphere. When the plant is dead,
the exchange stops & the proportion of C14 starts to decrease in the plant due to
radioactive decay. This decayed amount is
used to date the plant. After about 5,700
yrs, the proportion of C-14 falls to about
half its initial value.
Usages of radioisotopes
2. as tracers.
Radioactive isotope reacts chemically &
biologically. For example the activity of
thyroid gland can be measured w/
monitoring the increase in radioactivity of
the gland after taking a drink including
iodine radioisotopes (I-125 and I-131) since
the thyroid gland absorbs the radioactive
iodine when it works.
Usages of radioisotopes
3. Source of gamma rays and therefore,
source intense radioactivity.
Cobalt-60 is an example of such a
radioactivity source. It’s used in radiation
treatment for cancer and industrially as well
Mass spectrometry
• A mass spectrometer is an instrument
which separates particles according to
their masses , records the relative
proportions of these, and determine
natural abundances of the isotopes of an
element. Therefore, it also allows us to
calculate the atomic mass of an element.
• The most accurate way for determining
atomic and molecular weights is provided
by mass spectrometer.
Mass spectrometry
• mass spectrometer, invented by the English
physicist Francis William Aston (1877-1945)
when he was working in Cambridge with J. J.
Thomson. It was in his use of this instrument
that the existence of isotopes of elements was
discovered.
• Aston eventually discovered many of the
naturally occurring isotopes of non-radioactive
elements.
• He was awarded the Nobel Prize for Chemistry
in 1922.
Mass spectrometry
accelerating
C
A
B
F
D
E
A: a gaseous sample is very slowly
introduced to the mass spectrometer.
B: atoms/molecules are bombarded by a
stream of high energy electrons to produce
positive ions, mostly w/ a 1+ charge. These
electrons collide w/ electrons in the particle
knocking them out and leaving a positive
ion.
C: positively charged ions are accelerated
high enough to make the particles pass
through the slits and magnetic field by high
electrical voltage on the negatively charged
grid. With the slits, the ions were made a
beam of ions.
D: Fast moving ions enter a magnetic field
produced by an electromagnet. Ions are
deflected by a magnetic field into a curved
path. The deflection of the ions depends on
“charge to mass ratio(q/m). ”The more
massive the ion, the less the deflection. The
ions w/ equal mass and charge will deflect
the same.
E: By changing the strength of the magnetic
field or the accelerating voltage on the
negatively charged grid, ions of varying
masses can be made to enter the detector at
the end of the instrument.
• On the detector, ions are collected on a
metal plate and the current flows
through the metal plate to neutralise the
ions and this current is recorded.
• In this way, the relative abundances of
ions of different masses in the sample
can be determined and put into a graph
called “mass spectrum.”
F: The mass spectrometer must be at a
high vacuum for its correct operation
and its correct operation depends on
particles being able to pass through it
w/o colliding with any other particles.
• A: vapourised sample introduced
• B: ionization by electron bombardment
• C: Positive ions accelerated by electrical
field
• D: ions deflected by a magnetic field
• E: detector records ions of a particular
mass
• F: vacuum prevents molecules colliding
• Mass spectrometer is used to identify the
molecular structure of a compound or
analyze mixtures of substances.
• When a molecule loses an electron, it falls
apart, forming fragments. Mass
spectrometer measures the mass of these
fragments, producing a chemical
“fingerprint” of the molecule and providing
clues about how the atoms were connected
together in the molecule.
Atomic weight measurements
How was the atomic weight measured?
• By mass spectrometry
– This also measures
% natural abundance
for a given isotope
- The graph is called as
“mass spectrum.”
• P.53check out the graph.
Atomic weight calculation
There are three naturally occuring isotopes
of neon (Ne):
20Ne
21Ne
22Ne
isotopic mass = 19.99244018 amu
isotopic mass = 20.9938467 amu
isotopic mass = 21.9913855 amu
the atomic weight is reported in text as:
20.1797 amu
Learning Check 1
Naturally occurring carbon consists of three
isotopes, 12C, 13C, and 14C. State the number of
protons, neutrons, and electrons in each of
these carbon atoms.
12C
13C
14C
6
6
6
#P _______
_______
_______
#N _______
_______
_______
#E _______
_______
_______
Solution
12C
6
13C
14C
6
6
#P __6___
_ 6___
___6___
#N __6___
_ _7___
___8___
#E __6___
_ 6___
___6___
Learning Check 2
An atom of zinc has a mass number of 65.
A. Number of protons in the zinc atom
1) 30
2) 35
3) 65
B. Number of neutrons in the zinc atom
1) 30
2) 35
3) 65
C. What is the mass number of a zinc isotope
with 37 neutrons?
1) 37
2) 65
3) 67
Solution
An atom of zinc has a mass number of 65.
A. Number of protons in the zinc atom
1) 30
B. Number of neutrons in the zinc atom
2) 35
C. What is the mass number of a zinc isotope
with 37 neutrons?
3) 67
Learning Check 3
Write the atomic symbols for atoms with
the following:
A. 8 p+, 8 n, 8 e-
___________
B. 17p+, 20n, 17e-
___________
C. 47p+, 60 n, 47 e-
___________
Solution
16O
A. 8 p+, 8 n, 8 eB. 17p+, 20n, 17e-
8
37Cl
17
C. 47p+, 60 n, 47 e-
107Ag
47
Learning Check 4
An atom has 14 protons and 20 neutrons.
A. Its atomic number is
1) 14
2) 16
3) 34
B. Its mass number is
1) 14
2) 16
3) 34
C. The element is
1) Si
2) Ca
3) Se
D. Another isotope of this element is
1)
34X
16
2)
34X
14
3)
36X
14
Solution
An atom has 14 protons and 20 neutrons.
A. It has atomic number
1) 14
B. It has a mass number of
3) 34
C. The element is
1) Si
D. Another isotope of this element would be
3) 36X
14
Masses of Atoms
 A scale designed for atoms gives their small
atomic masses in atomic mass units (amu)
 An atom of 12C was assigned an exact mass of
12.00 amu
 Relative masses of all other atoms was
determined by comparing each to the mass of
12C
 An atom twice as heavy has a mass of 24.00
amu. An atom half as heavy is 6.00 amu.
Average atomic mass(atomic
weight)
• “atomic weight or mass” = average
mass of an atom calculated from the
masses and natural abundances of all
isotopes
(use atomic weights to calculate the
molecular weights of compounds from
their constituent elements!)
Atomic Mass
Na
22.99
 Average atomic mass is based on all the
isotopes and their abundance %
Atomic mass is not a whole number
Calculating Atomic Weight or
Mass
 Percent(%) abundance of isotopes
 Mass of each isotope of that element
 Weighted average =
mass isotope1(%) + mass isotope2(%) + …
100
100
• Naturally occurring C is composed of 98.93
% 12C and 1.07 % 13C. The masses of these
nuclides are 12 amu (exactly) and 13.00335
amu, respectively.
(98.93x12amu )  (1.07 x13.00335amu )
 12.01amu
100
Average atomic
mass(atomic mass) or
atomic weight
Atomic Mass of Magnesium
Isotopes
Mass of Isotope
Abundance
24Mg
=
24.0 amu
78.70%
25Mg
=
25.0 amu
10.13%
26Mg
=
26.0 amu
11.17%
Atomic mass (average mass) Mg = 24.3 amu
Mg
24.3
Atomic mass calculation
How was the atomic mass calculated?
• multiply each isotopic mass by the
reported natural abundance for the
isotope, then:
• add these individual contributions for
each isotope to get the average atomic
mass for the element
Atomic mass calculation
There are three naturally occuring isotopes of
neon (Ne):
20Ne mass # = 19.99244018 amu (90.51%)
21Ne mass # = 20.9938467 amu
(0.27%)
22Ne mass # = 21.9913855 amu
(9.22%)
the atomic mass is reported in text as:
20.1797 amu
18.10 + 0.057 + 2.03 = 20.19 amu
Learning Check 5
Gallium is a metallic element found in
small lasers used in compact disc players.
In a sample of gallium, there is 60.2% of
gallium-69 (68.9 amu) atoms and 39.8% of
gallium-71 (70.9 amu) atoms. What is the
atomic mass of gallium?
Solution
Ga-69
68.9 amu x
60.2
=
41.5 amu for
69Ga
28.2 amu for
71Ga
100
Ga-71 (%/100)
70.9 amu x 39.8
=
100
Atomic mass Ga =
69.7 amu
Finding An Isotopic Mass
A sample of boron consists of 10B (mass
10.0 amu) and 11B (mass 11.0 amu). If
the average atomic mass of B is 10.8
amu, what is the % abundance of each
boron isotope?
Assign X and Y values:
X = % 10B
Y = % 11B
Determine Y in terms of X
X
+
Y
= 100
Y = 100 - X
Solve for X:
X (10.0) + (100 - X )(11.0)
100
100
= 10.8
Multiply through by 100
10.0 X + 1100 - 11.0X = 1080
Collect X terms
10.0 X - 11.0 X
=
1080 - 1100
- 1.0 X = -20
X
=
-20
- 1.0
=
Y = 100 - X
% 11B = 100 - 20% =
20 %
10B
80% 11B
Learning Check 6
Copper has two isotopes 63Cu (62.9 amu)
and 65Cu (64.9 amu). What is the %
abundance of each isotope? (Hint: Check
Zumdahl or any other chemistry text for
atomic mass)
1) 30%
2) 70%
3) 100%
Solution
2) 70%
Solution
62.9X + 6490 = 64.9X = 6350
-2.0 X = -140
X = 70%
Atomic Masses
13C
12C
13.00335 amu (1.11%)
12.0000 amu (98.89%)
atomic weight of C = 12.01115 amu WHY?
Calculating masses of
atoms relative to 12C
(mass of 12C atom) * 1.58320 = mass of F
atom
= 18.99840
reported atomic weight of F = 18.9984
Charged Atoms: Ions
Left to their own devices, atoms are electrically neutral.
That means that they have an equal number of
protons and electrons.
During the course of most natural events,
protons are not gained or lost, but electrons may be.
Atoms with more or fewer electrons than protons are
electrically charged. They are called ions:
an atom that loses electrons takes on a positive charge
(cation);
an atom that gains electrons takes on a negative charge
(anion).
An ISOTOPE is one of a set of nuclides with the same Z and
consequently different A. (ie isotopes are the same chemical
element but different masses). e.g.
12
6C
13
6C
14
6C
An ISOTONE is one of a set of nuclides with the same N
and consequently different A. e.g.
39
40
41
A
,
K
,
18 21 19 21 20 A 21
An ISOBAR is one of a set of nuclides with the same A but
different N and Z.
e.g
14
6C
14
7N
14
8O
• More on atomic notation, which is based on the nuclear
structure:
– Isotope: same Z, different A and N
– Isobar: same A, different Z and N
– Isotone: same N, different Z and A
Example: From the following list of atoms, which are isotopes, isobars, and isotones?
131
54
Component
Atom
Xe
I
Cs
I
Xe
130
53
A
I
132
55
Cs
Z
131
53
I
N