Transcript Atomic structure

Chapter 1
Atomic Structure,
Radioactivity and Relative
Masses
1.1
The Atomic Nature of Matter
1.2
The Experimental Evidence of
Atomic Structure
1.3
1
Sub-atomic Particles
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2
1.4
Atomic Number, Mass Number and
Isotopes
1.5
Radioactivity
1.6
Nuclear Reactions
1.7
Uses of Radioactive Isotopes
1.8
Mass Spectrometer
1.9
Relative Isotopic, Atomic and
Molecular Masses
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1.1 The atomic nature of matter (SB p.2)
What is “atom”?
The Greek philosopher Democritus
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1.1 The atomic nature of matter (SB p.2)
These are iron
atoms!!
Continuous
division
Iron
Continuous
division
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1.1 The atomic nature of matter (SB p.2)
Dalton’s atomic theory
John Dalton proposed
his Dalton’s atomic
theory.
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1.1 The atomic nature of matter (SB p.2)
Main points of Dalton’s atomic theory
1.
All elements are made up of atoms.
2.
Atoms can neither be created nor destroyed.
3.
Atoms of the same element are identical. They have the same
mass and chemical properties.
4.
Atoms of different elements are different. They have different
masses and chemical properties.
5. Atoms of different elements combine to form a compound. The
numbers of various atoms combined bear a simple whole
number ratio to each other.
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1.2 The experimental evidence of atomic structure (SB p.3)
Discovery of electrons
- A beam of rays came out from the cathode and
hit the anode.
- He called the beam cathode rays.
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1.2 The experimental evidence of atomic structure (SB p.4)
The beam was composed of negatively charged fastDeflected in the
Deflected
in
the
moving particles.
magnetic field
electric field
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1.2 The experimental evidence of atomic structure (SB p.4)
Measure the mass to
charge ratio (m/e) of the
particles produced
The particles were
constituents of all
atoms!!
9
Independent of the nature
of the gas inside the
discharge tube
He called the particles
‘electrons’.
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1.2 The experimental evidence of atomic structure (SB p.4)
Thomson’s atomic model
An atom is electrically
neutral
Atom
10
No. of
positively
charged
particles
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=
No. of
negatively
charged
particles
1.2 The experimental evidence of atomic structure (SB p.4)
How are the particles distributed in an atom?
- An atom was a positively
charged sphere
+
+
+
+
+
+
Positive
charge
11
- Negatively charged
electrons embedded in it
like a ‘raisin pudding’
Electron
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1.2 The experimental evidence of atomic structure (SB p.4)
Gold foil scattering experiment
- performed by Ernest
Rutherford
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1.2 The experimental evidence of atomic structure (SB p.4)
He bombarded a thin gold foil with a beam of fastmoving -particles (+ve charged)
Observation:
-most -particles passed
through the foil without
deflection
-very few -particles
were scattered or
rebounded back
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1.2 The experimental evidence of atomic structure (SB p.5)
Interpretation of the experimental results
- The condensed core is called ‘nucleus’.
- The positively charged particle is called ‘proton’.
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1.2 The experimental evidence of atomic structure (SB p.5)
Rutherford’s atomic model
Expectation:
Mass of atom
= Total mass of protons
Mass of atom
Total mass of protons
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>
1.2 The experimental evidence of atomic structure (SB p.5)
Chadwick’s atomic model
- the presence of neutrons
- proved by James
Chadwick
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1.2 The experimental evidence of atomic structure (SB p.5)
Chadwick’s atomic model
Proton
Electron
Neutron
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1.3 Sub-atomic particles (SB p.6)
Sub-atomic particles
3 kinds of sub-atomic particles:
- protons
- neutrons
Inside the condensed
nucleus
- electrons
Moving around the
nucleus
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1.3 Sub-atomic particles (SB p.7)
A carbon-12 atom
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1.3 Sub-atomic particles (SB p.6)
Characteristics of sub-atomic particles
Sub-atomic
particle
Symbol
Proton
p or 1H
1
Neutron
n or 1 n
0
Electron
e- or 0 e
-1
Location in atom
Nucleus
Nucleus
Surrounding the
nucleus
Actual charge (C)
1.6 x 10-9
0
1.6 x 10-9
Relative charge
+1
0
-1
Actual mass (g)
1.7 x 10-24
1.7 x 10-24
9.1 x 10-28
1
1
0
Approximate
relative mass
(a.m.u.)
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1.3 Sub-atomic particles (SB p.6)
Relative size of the atom and the nucleus
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1.4 Atomic number, mass number and isotopes (SB p.7)
Atomic number
The atomic number (Z) of an element is the number
of protons contained in the nucleus of the atom.
Atomic
number
=
Number of
protons
=
Reason: Atoms are
electrically neutral.
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Number of
electrons
1.4 Atomic number, mass number and isotopes (SB p.8)
Mass number
The mass number (A) of an atom is the sum of the
number of protons and neutrons in the nucleus.
Mass
number
23
=
Number of
protons
+
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Number of
neutrons
1.4 Atomic number, mass number and isotopes (SB p.8)
Atomic numbers and mass numbers
of some common atoms
Atom
No. of
protons
No. of
electrons
No. of
neutrons
Atomic
number
Hydrogen
1
1
0
1
(1 + 0)
=1
Oxygen
8
8
8
8
(8 + 8)
= 16
Argon
18
18
22
18
(18+22)
= 40
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Mass
number
1.4 Atomic number, mass number and isotopes (SB p.8)
Isotopes
Isotopes are atoms of the same element with the same
number of protons but different numbers of neutrons.
Representation:
Mass
number
Symbol of the
element
A
X
Atomic
number
25
Z
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1.4 Atomic number, mass number and isotopes (SB p.8)
e.g. the two isotopes of chlorine are written as:
35
Cl
17
37
Cl
17
OR labelled as Cl-35 and Cl-37.
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1.4 Atomic number, mass number and isotopes (SB p.9)
Isotopes of some common elements
Element
Hydrogen
Carbon
27
Isotope
Atomic
number
No. of
protons
No. of
neutrons
Natural
abundanc
e (%)
1
1H
2
1H
1
1
0
99.8
1
1
1
0.02
12
6C
13
6C
14
6C
6
6
6
98.89
6
6
7
1.11
6
6
8
trace
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1.5 Radioactivity (SB p.10)
Nuclear stability
binding force
p
n
n
p
Strong binding force
The atom is stable
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1.5 Radioactivity (SB p.10)
The stability of an isotope depends on its neutron to
proton ratio.
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1.5 Radioactivity (SB p.10)
What will happen to the unstable nuclei?
- split and divide to give smaller nuclei
smaller
nuclei
unstable
nuclei
smaller
nuclei
This process is called nuclear fission.
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1.5 Radioactivity (SB p.10)
High energy radiation and smaller particles may be
emitted.
smaller
particles
high energy
radiation
unstable
nuclei
This phenomenon is called radioactivity.
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1.5 Radioactivity (SB p.10)
Radiation:
The energy and particles emitted
when a nucleus splits.
high energy
radiation
smaller
particles
Radioactive isotopes:
Isotopes with unstable nuclei
unstable
nuclei
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1.5 Radioactivity (SB p.11)
Symbols for some common nuclides
in nuclear reactions
Particle
Symbol
Proton
p or
1
1
Neutron
n or
1
0
Electron (beta particles)
e- or 0
H
n
e
-1
Helium nucleus (alpha
particle)
33
4
2
He or 
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or 
1.5 Radioactivity (SB p.11)
Discovery of radioactivity
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1.5 Radioactivity (SB p.11)
- Ernest Rutherford passed a beam of radiation from a
radioactive source through electrically charged plates
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1.5 Radioactivity (SB p.12)
-particles
- Helium nucleus
- Mass: 4 a.m.u.
+
- Charge: +2
+
helium
nucleus
36
- Deflected by both electric and
magnetic fields.
- Weak penetrating power
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1.5 Radioactivity (SB p.12)
e.g.
238
92 U

+
+
He
Th
90
2
234
1. Atomic number
2. Mass number
37
4
2
4
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1.5 Radioactivity (SB p.13)
 -particles
- Fast-moving electron
- Negligible mass
- Charge: -1
-
- Deflected much more readily
by an electric field than particles
fast-moving
electron
38
- Moderate penetrating power
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1.5 Radioactivity (SB p.13)
e.g.
234
92 Th

+
+
e
Pa
91
-1
0
234
1. Atomic number
2
2. Mass number is not affected
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1.5 Radioactivity (SB p.13)
 -radiation
- Electromagnetic radiation of
short wavelength
- Negligible mass
- No charge
- NOT deflected by both electric
and magnetic fields.
- Strong penetrating power
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1.5 Radioactivity (SB p.13)
Relative penetrating power of -particles,
 -particles and  -radiation
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1.5 Radioactivity (SB p.13)
Tracks of -particles and  -particles
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1.6 Nuclear reactions (SB p.14)
Chemical reactions
- Rearrangement of electrons
- Numbers of protons and neutrons in the nuclei
remain unchanged
- NO new elements are formed
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1.6 Nuclear reactions (SB p.14)
Nuclear reactions
- Rearrangement of
protons and neutrons
- New elements are
formed
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1.6 Nuclear reactions (SB p.14)
Naturally occurring radioactive decay
1. From tritium to helium
- tritium undergoes  -decay to form a stable
isotope of helium 3
2
3H
1
46
3 He
2
He
+ 0e
-1
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+ 
1.6 Nuclear reactions (SB p.14)
1. Mass number of the atom is not affected
2. Atomic number
47
1
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1.6 Nuclear reactions (SB p.14)
2. From carbon to nitrogen
- carbon undergoes  -decay to form a stable
isotope of nitrogen
14C
6
48
14N
7
+ 0e
-1
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+ 
1.6 Nuclear reactions (SB p.14)
1. Mass number of the atom is not affected
2. Atomic number
49
1
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1.6 Nuclear reactions (SB p.15)
3. From uranium to lead
- uranium-238 undergoes  -decay to form
thorium-234
238U
92
50
234Th + 4 He
90
2
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+ 
1.6 Nuclear reactions (SB p.15)
238
Decay series from 92 U to
51
206
82
Pb
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1.6 Nuclear reactions (SB p.15)
Artificial nuclear reactions
-bombarding certain stable nuclei with 1
particles, -particles, neutrons, protons ( H)
1
2
and deuterons (
).
1
e.g.
9
H
1
4 Be + 1 H
14
2
7N + 1H
23
1
11 Na + 0 n
52

+
B
5
10

+
+
n
O
8
0
1
15
23
10
Ne + 1 H + 
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1
1.6 Nuclear reactions (SB p.15)
Nuclear equations
- changes of mass numbers and atomic numbers
e.g.
9
4
4 Be + 2 He
Sum of mass numbers and
sum of atomic numbers on =
the L.H.S.
53

n
+
+
C
0
6
12
1
Sum of mass numbers and
sum of atomic numbers on
the R.H.S.
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1.6 Nuclear reactions (SB p.16)
Half-life of radioactive isotopes
Half-life of a radioactive isotope is the time taken for
its radioactivity to drop to half of its initial value.
- Not affected by:
1) no. of radioactive nuclei,
2) chemical conditions,
3) temperature
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1.6 Nuclear reactions (SB p.16)
Decay of a 16 g sample of phosphorus-32
Each passage of a half-life
causes one half of sample
remains!
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1.6 Nuclear reactions (SB p.16)
Half-lives of some radioactive isotopes
Radioactive isotope
Oxygen-13
Radon-222
Iodine-131
Phosphorus-32
Cobalt-60
Hydrogen-3
Carbon-14
Plutonium-239
Uranium-238
56
Half-life
8.7 x 10-3 seconds
3.8 days
8.06 days
14.3 days
5.32 years
12.3 years
5730 years
24 400 years
A wide
4.5 billion years
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range!
1.6 Nuclear reactions (SB p.16)
-kt
N = N0 e
N: amount left
N0: initial amount
k: constant
t: time taken for the decay
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1.6 Nuclear reactions (SB p.16)
Check Point 1-1
(a) Balance the following equation:
(i)
(ii)
(iii)
(iv)
(v)
58
14
7
4
He
2
N +
7
Li
3
16
8
+
1
H
1
O +
238
U
92
2
1
n
0
+
U
92
238
17
1
H
1
O +
8
4
He
2
+ 
13
C +
6
239
1
n
0
92
238
Np +
93
+ 
U
4
He
2
+ 
+ 
0
-1 e
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+ 
Answer
1.6 Nuclear reactions (SB p.16)
Check Point 1-1 (cont’d)
(b) A series of radioactive decays can be represented
by the following equation.
232
90
 -decay
-decay
 -decay
Th
Y
Z
X
Write symbols for the elements X, Y and Z.
X:
Y:
Z:
59
228
88
228
89
228
Ra
Ac
Th
90
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Answer
1.6 Nuclear reactions (SB p.16)
Check Point 1-1 (cont’d)
(c) Give three differences between -particle, -particle
and -radiation.
-particle -particle -radiation
Charge
Mass
Relative
penetrating
power
60
+2
-1
0
4 a.m.u.
0 a.m.u.
0 a.m.u.
low
moderate
high
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Answer
1.7 Uses of radioactive isotopes (SB p.17)
Uses of radioactive isotopes
Five main uses of radioactive isotopes:
1. Leak detection
2. As tracers
3. Radiotherapy
4. Carbon-14 dating
5. Nuclear power
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1.7 Uses of radioactive isotopes (SB p.17)
Leak detection
- short-lived radioactive source is introduced into
storage tanks and underground pipelines
- located with the Geiger-Muller counter
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1.7 Uses of radioactive isotopes (SB p.17)
As tracers
- detection of the metabolic pathway of an element in a
living organism
- e.g. I-131, P-32
I-131 is used for
diagnosing thyroid disease
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1.7 Uses of radioactive isotopes (SB p.18)
Radiotherapy
For those cancer cells located
deep inside the body,
 -radiation (from Co-60 and
Ce-137) is used.
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1.7 Uses of radioactive isotopes (SB p.18)
Carbon-14 dating
It is estimated that the bowl
How
can
the
age
of
this
was made >5000 years ago!
bowl be estimated?
Half-life of C-14: 5730
yrs
Changes to a stable
isotope, N-14
Measure the radioactivity
of C-14 in the bowl
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1.7 Uses of radioactive isotopes (SB p.18)
Nuclear power
- nuclear fission can be used as a source of energy
- e.g. in the disintegration of
235
92
235
92
66
U  n  Sr 
1
0
90
38
U
Xe 3 n
143
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1
0
1.7 Uses of radioactive isotopes (SB p.18)
A chain reaction
Initial
neutron
235
92
67
U
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1.8 Mass spectrometer (SB p.20)
Cl has 2 isotopes: Cl-35 and Cl-37
Isotopes
Relative
abundance
Cl-35
75%
Cl-37
25%
How can we
know?
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1.8 Mass spectrometer (SB p.20)
Mass spectrometer
A highly accurate
instrument!
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1.8 Mass spectrometer (SB p.20)
Mass spectrometer consists of 6 parts:
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1.8 Mass spectrometer (SB p.21)
Mass spectrum of Cl2:
m/e ratio
71
35
Corresponding
ion
35Cl+
37
37Cl+
70
35Cl─35Cl+
72
35Cl
─ 37Cl+
74
37Cl
─37Cl+
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1.8 Mass spectrometer (SB p.21)
Mass spectrum of CH3Cl:
m/e ratio
35
Corresponding
ion
35Cl+
37
37Cl+
50
72
12CH
37Cl+
─
3
51
13CH
52
12CH
New Way Chemistry for Hong Kong A-Level Book 1
35Cl+
─
3
3
─37Cl+
1.9 Relative isotopic, atomic and molecular masses (SB p.22)
Relative isotopic mass
The relative isotopic mass of a particular isotope of an
element is the relative mass of one atom of that isotope
on the carbon-12 scale.
e.g.
relative isotopic mass of Cl-35 = 35
relative isotopic mass of Cl-37 = 37
73
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1.9 Relative isotopic, atomic and molecular masses (SB p.22)
What is carbon-12 scale?
- use carbon-12 as the reference standard
Mg has the same
mass as two C-12
atoms
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1.9 Relative isotopic, atomic and molecular masses (SB p.23)
Relative atomic mass
The relative atomic mass of an element is the weighted
average of the relative isotopic masses of its natural
isotopes on the carbon-12 scale.
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1.9 Relative isotopic, atomic and molecular masses (SB p.23)
What is the relative atomic mass of Cl?
The relative abundances of Cl35 and Cl-37 are 75.77 and
24.23 respectively
Relative atomic mass of Cl
(35  75.77)  (37  24.23)
=
(75.55  24.23)
= 35.48
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1.9 Relative isotopic, atomic and molecular masses (SB p.23)
Relative molecular mass
The relative molecular mass is the relative mass of a
molecule on the carbon-12 scale.
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1.9 Relative isotopic, atomic and molecular masses (SB p.23)
What is the relative molecular mass of CH3Cl?
Relative molecular mass of
CH3Cl
(50 123)  (51  2)  (52  40)
=
(123  2  40)
= 35.48
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1.9 Relative isotopic, atomic and molecular masses (SB p.23)
Example 1-1
The mass spectrum of
neon is given on the right.
Determine the relative
atomic mass of neon.
Solution:
Relative atomic mass of
neon
=
(20  114)  (21  0.2)  (22  11.2)
(114  0.2  11.2)
= 20.18
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Answer
1.9 Relative isotopic, atomic and molecular masses (SB p.24)
Check Point 1-2
(a) The mass spectrum of
lead is given on the
right.
the abundance of the peak
Let xGiven
be thethat
relative
relative
atomic mass
at m/e 208.
of lead is 207.242,
(204  1.5
206  23.6 + 207  22.6 +
calculate
the+relative
208x)  (1.5
+ 23.6 + 22.6 + x) = 207.242
abundance
of the
peak
 at m/e 208.
x = 52.3
 The relative abundance of the peak at
m/e 208 is 52.3.
Answer
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1.9 Relative isotopic, atomic and molecular masses (SB p.24)
Check Point 1-2 (cont’d)
(b)The mass spectrum of
dichloromethane
is given
on
Let y be the relative
molecular
mass of
thedichloromethane.
right. Calculate the
relative molecular mass of
y = (84  94 + 85  3.0 + 86  59 + 87  2.2 + 88
dichloromethane.
 13 + 89  2.5 + 90  0.8 )  (90 + 3.0 + 59
+ 2.2 + 13 + 2.5 + 0.8)
= 85.128
 The relative molecular mass of
dichloromethane is 85.128.
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Answer
The END
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