Topic 2.1: Atomic
Structure
Honors Chemistry 2014-15
Mrs. Peters
1
Atomic Structure
2.1: The nuclear atom
EI: The mass of the atoms is concentrated in its minute, positively
charged nucleus.
NOS:
1.Evidence and improvements in instrumentation – alpha particles
were used in the development of the nuclear model of the atom
that was first proposed by Rutherford. (1.8)
2.Paradigm shifts- the subatomic particle theory of matter
represents a paradigm shift in science that occurred in the late
1800s (2.3)
2
Atomic Structure
2.1: The nuclear atom
Understandings:
1. Atoms contain a positively charged dense nucleus composed
of protons and neutrons (nucleons)
2. Negatively charged electrons occupy the space outside the
nucleus
3. The mass spectrometer is used to determine the relative
atomic mass of an element from its isotopic composition.
3
Atomic Structure
2.1: The nuclear atom
Applications and Skills:
1. Use of the nuclear symbol notation AZX to deduce the number
of protons, neutrons, and electrons in atoms and ions.
2. Calculations involving non-integer relative atomic masses and
abundance of isotopes from given data, including mass
spectra.
4
NOS: Paradigm Shift
History behind Atomic Theory
• Democritus (420 BCE) first proposed the idea that
matter may be made up of small, indivisible
particles called atoms.
• Aristotle (384-322 BCE)
Greek philosopher; matter composed of earth, air,
fire, water. This view dominated thought until 17th
century
5
NOS: Paradigm Shift
History behind Atomic Theory
• Atomism developed in Chinese & Arabic cultures
during the Dark Ages in Europe.
• John Dalton (1766-1844) was the first to base
atomic theory on scientific evidence.
6
NOS: Paradigm Shift
Dalton’s Atomic Theory
• Elements are made of tiny particles called atoms.
• All atoms of a given element are identical. The
atoms of a given element are different from those
of any other element.
7
NOS: Paradigm Shift
Dalton’s Atomic Theory
• Atoms of one element can combine with atoms of
other elements to form compounds. A given
compound always has the same relative number of
types of atoms.
• Atoms cannot be created, nor divided into smaller
particles, nor destroyed in the chemical process. A
chemical reaction simply changes the way atoms
are grouped together.
8
NOS: Paradigm Shift
Evidence for sub-atomic particles
1897: J.J. Thomsen: Cathode Ray Tube
Evidence for electrons: Bent a stream of rays
originating from the negative electrode
(cathode). Stream of particles with mass &
negative charge.
9
NOS: Paradigm Shift
Evidence for sub-atomic particles
1909: Ernest Rutherford: Gold Foil
Evidence for protons & nucleus: Alpha
particles deflected passing through gold foil
10
NOS: Paradigm Shift
Evidence for sub-atomic particles
1932: James Chadwick: Beryllium
Evidence for neutrons: Alpha particles caused
beryllium to emit rays that could pass
through lead but not be deflected,
11
U1. and U2. Atomic Structure
Sub-Atomic Particles:
Proton:
Located in the nucleus
Relative charge of +1
Relative mass of 1 amu
Neutron:
Located in the nucleus
Relative charge of 0
Relative mass of 1 amu
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12
U1. and U2. Atomic Structure
Sub-Atomic Particles
Electron:
Located in cloud surrounding
the nucleus
Relative charge of –1
Relative mass of 0.0005 amu
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13
U1. and U2. Atomic Structure
Nucleus consists of protons and neutrons with the
electrons surrounding the nucleus.
In a neutral atom, the #protons = # electrons.
14
A1. Nuclear Symbol Notation
Atomic Number (Z)
The atomic number is the number of protons in the
nucleus. It determines the identity of an atom.
•
•
All oxygen atoms have 8 protons in the nucleus
All lead atoms have 82 protons in the nucleus
15
A1. Nuclear Symbol Notation
Atomic Number (Z)
It also tells us the number of electrons in a neutral
atom
•
A neutral sodium atom contains 11 protons and 11
electrons
•
A neutral bromine atom contains 35 protons and 35
electrons
16
A1. Nuclear Symbol Notation
Mass Number (A)
It is not practical to measure the masses of atoms in grams
due to their small size. Scientists devised a measurement
called atomic mass units (amu).
17
A1. Nuclear Symbol Notation
Mass Number (A)
•
•
•
Protons have a mass of 1 amu
Neutrons have mass of 1 amu
Electrons have mass of 0 amu.
Mass Number of atoms = # protons + # neutrons
18
A1. Nuclear Symbol Notation
Mass Number (A)
Mass Number of atoms = # protons + # neutrons
**Round the Relative Atomic Mass to a whole number to
find the Mass Number
Ex:
• Lithium = 6.94 Mass Number is 7
• Magnesium = 24.31 Mass Number is 24
19
How to read the Periodic
Table
Lithium
3
Li
6.94
Element Name
Atomic Number
Element Symbol
Relative Atomic Mass
20
A1. Nuclear Symbol Notation
Atomic Name: Element Name - A (mass number)
Ex: Carbon-12
Nuclear Symbol:
Mass Number (Protons + Neutrons)
A
X
Z
Element Symbol
Atomic Number
21
Important Terms
Isotope
Atoms of the same element can have different numbers of
neutrons, thus they will have different atomic masses.
These are called isotopes of the element.
These are the same element, just different numbers of
neutrons and mass.
22
Important Terms
Isotope Example
There are three isotopes of hydrogen:
•
Hydrogen-1 has 1 proton, 1 electron, 0 neutrons
•
Hydrogen-2 has 1 proton, 1 electron, 1 neutron
•
Hydrogen-3 has 1 proton, 1 electron, 2 neutrons
23
A1. Deduce the symbol given its mass
number and atomic number
•
Consider an atom that has an atomic number of
29 and a mass number of 63. What is its name
and symbol?
Name:
Symbol:
24
A. Deduce the symbol given its mass number
and atomic number
•
Consider an atom that has an atomic number of
29 and a mass number of 63. What is its name
and symbol?
atomic number of 29 identifies it as copper
Name: Copper-63
Symbol: 63 Cu
29
25
A1. Deduce the symbol given its mass
number and atomic number
•
Consider an atom that has A=32 and Z=16. What
is its name and symbol?
Name:
Symbol:
26
A1. Deduce the symbol given its mass
number and atomic number
•
Consider an atom that has A=32 and Z=16.
What is its name and symbol?
Z=16 identifies it as sulfur
Name: Sulfur-32
Symbol: 32 S
16
27
A1. Deduce the symbol given its mass number
and atomic number
•
Consider an atom that has an atomic number of 74 and a mass
number of 185. What is its name and symbol?
•
Consider an atom that has A=127 and Z=53. What is its name
and symbol?
28
A1. Deduce the symbol given its mass number
and atomic number
•
Consider an atom that has an atomic number
of 74 and a mass number of 185. What is its
name and symbol?
atomic number of 74 identifies it as tungsten
Name: Tungsten-185
Symbol: 185 W
74
•
Consider an atom that has A=127 and Z=53.
What is its name and symbol?
Z=53 identifies it as iodine
Name: Iodine-127
Symbol: 127 I
53
29
A1. Deduce protons, neutrons, and electrons in
atoms and ions from the A, Z, and charge
•
Consider the neutral carbon-12 atom. Find the A,
Z, protons, neutrons, electrons, and symbol
Name is Carbon-12
Atomic mass (A) = 12
Atomic number (Z) = 6
Protons = 6 (atomic number)
Neutrons = 6 (mass – protons)
Electrons = 6 (neutral atom so same as protons)
Symbol is
12
6
C
30
A1. Deduce protons, neutrons, and electrons in
atoms and ions from the A, Z, and charge
•
Consider an atom that has 9 protons, 9
electrons, and 10 neutrons. What is its atomic
number, atomic mass, name, and symbol?
Z=9 (atomic number = # protons)
A=19 (atomic mass = protons + neutrons)
Fluorine-19 (name and mass)
19
F (neutral because protons = electrons)
9
31
A1. Deduce protons, neutrons, and electrons in
atoms and ions from the A, Z, and charge
•
Consider a neutral atom with A=75 and Z=33.
How many protons, neutrons, and electrons are
in the atom. What is the name and symbol?
•
Consider a neutral atom with A=77 and Z=33.
How many protons, neutrons, and electrons are
in the atom. What is the name and symbol?
32
A1. Deduce protons, neutrons, and electrons in
atoms and ions from A, Z, and charge
•
Consider a neutral atom with A=75 and Z=33. How many protons,
neutrons, and electrons are in the atom. What is the name and symbol?
Protons = 33
Neutrons = 42
Name: Arsenic-75
Symbol: 75As
Electrons = 33
33
•
Consider a neutral atom with A=77 and Z=33. How many protons,
neutrons, and electrons are in the atom. What is the name and symbol?
Protons = 33
Neutrons = 44
Name: Arsenic-77
Symbol: 77As
Electrons = 33
33
33
Important Terms
Ions are charged particles formed when atoms gain or lose
electrons resulting in unequal numbers of protons and electrons
Cations: Atoms that lose electrons become positively charged
Anions: Atoms that gain electrons become negatively charged
34
A1. Use of Nuclear Symbol
Notation
Nuclear Symbol:
Mass Number (Protons + Neutrons)
A
Z
X
+
Charge (+, - or nothing)
Determined by electrons
Element Symbol
Atomic Number
35
Important Terms
•
How many protons, neutrons, and electrons are in an ion of K-39
that has lost one electron? What is the charge of the ion? What
is its symbol?
Protons = 19
Neutrons = 20
Electrons = 18
Charge = 1+ or +1
Symbol is 39K1+
19
36
A1. Deduce from nuclear symbol notation
•
The symbol of an anion is 31P 3- . Calculate the number of
15
protons, neutrons, and electrons. What is Z and what is A?
•
What is the symbol of a species containing 26 protons, 30
neutrons, and 23 electrons?
•
What is the symbol of a species with A=56, Z=26, and 24
electrons?
37
A1. Deduce the number of protons, neutrons,
and electrons in atoms and ions from A, Z, and
charge
•
The symbol of an anion is 31P 3- . Calculate the number
15
protons, neutrons, and electrons. What is Z and what is A?
#P = 15; #N = 16; #E = 18; Z= 15; A = 31
•
What is the symbol of a species containing 26 protons, 30
neutrons, and 23 electrons?
56Fe 3+
26
•
What is the symbol of a species with A=56, Z=26, and 24
electrons?
56Fe 2+
26
38
A2. Discuss the use of radioisotopes.
Radioisotopes: isotopes of elements that have become
radioactive because the nucleus is unstable and breaks
down spontaneously emitting radiation.
Radioisotopes can occur naturally or be created artificially
Examples: Carbon-14; Iodine-125; Strontium-90; Cobalt-60, Iodine-131
39
A2. Discuss the use of radioisotopes.
Uses of Radioisotopes
•
•
•
•
•
•
Nuclear power generation
Sterilization of surgical instruments
Crime detection
Food preservation
Dating artifacts
Treating and diagnosing disease
40
U3. The Mass Spectrometer
Mass Spectrometers:
Instruments that measure charge-to-mass ratio of charged particles.
Used to measure masses of isotopes as well as isotopic abundance
41
U3. The Mass Spectrometer
How a Mass Spec Works:
1. Vaporization: sample is heated to gas
state
2. Ionization: sample gas is turned into
ions by blasting free electrons to knock
electrons off from the gas atoms,
creating positive ions
3. Acceleration: increases the speed of
particles, using an electric field
42
U3. The Mass Spectrometer
How a Mass Spec Works:
4. Deflection: using an electromagnet to
create a magnetic field, amount of
deflection depends on mass and charge
of the ion (think of cars going around a
corner)
5. Detection: measures both mass and
relative amounts (abundance) of all the
ions present
43
U3. The Mass Spectrometer
• Mass Spectrometer Video
• http://www.youtube.com/watch?feature=fv
wp&v=lxAfw1rftIA&NR=1
44
U3. Relative Atomic Mass
Relative Atomic Mass
Mass numbers (atomic mass) on the
periodic table are weighted
averages of the isotopes.
Based on 12C.
Has 6 protons, 6 neutrons, and 6
electrons
Has a relative atomic mass of exactly
12.000
45
U3. Relative Atomic Mass
Relative Atomic Mass
One amu is exactly 1/12 of the mass of
a carbon-12 atom.
All other isotopes are measured
compared to this value.
46
U3. Relative Atomic Mass
Average relative atomic mass:
the weighted average for all of the
isotopes of a given element, based
on the percent abundance of each
47
U3. Relative Atomic Mass
To determine Average Relative
Atomic Mass:
• Need masses of each isotopes
• Need abundance (percentage)
of each isotope
o The mass spec is used to determine
these values
• This is the value shown on the
periodic table
48
U3. Relative Atomic Mass
• A sample of neon is
placed in the mass
spectrometer
49
U3. Relative Atomic Mass
• A sample of neon is
placed in the mass
spectrometer
• The results show the
abundance for each
isotope of an element
o 90.92% is neon-20
o 0.26% is neon-21
o 8.82% is neon-22
50
A2. Calculate non-integer relative atomic masses
and abundance of isotopes from given data.
How to Determine Relative
Atomic Mass
Example
1. Convert the percent
1. 95.5% = .955 .5%=0.005
abundance for each
isotope into decimal
2. 24.5 x .955= 23.4 and
2. Multiply the mass for
23.7 x .005 = .119
each isotope by the
abundance
3. 23.4 + .119 = 23.519
3. Add all product values
4. 23.5 amu
from step 2.
4. Include amu for the units
of the value.
51
A2. Calculate non-integer relative atomic
masses and abundance of isotopes from given
data.
Three isotopes of magnesium occur in nature.
Their abundances and masses, determined by
mass spectrometry, are listed in the table on
the right. Use this information to calculate the
atomic weight of magnesium.
•
•
•
•
Three isotopes: 24, 25, 26
Percentage of each isotope: Given
Multiply the percent of each isotope by its
mass
23.98504 x .7899 = 18.95 amu
24.98584 x .1000 = 2.499 amu
25.98259 x .1101 = 2.861 amu
Add these values = 24.31 amu
Isotope
% Abundance
Mass (amu)
24Mg
78.99
23.98504
25Mg
10.00
24.98584
26Mg
11.01
25.98259
52
A2. Calculate non-integer relative atomic
masses and abundance of isotopes from given
data.
Calculate the atomic weight of chromium using the following data for the
percent natural abundance and mass of each isotope:
4.35% 50Cr (49.9461 amu); 83.79% 52Cr (51.9405 amu);
9.50% 53Cr (52.9406 amu); 2.36% 54Cr (53.9389 amu)
53
A2. Calculate non-integer relative atomic
masses and abundance of isotopes from given
data.
Calculate the atomic weight of chromium using the following data for the
percent natural abundance and mass of each isotope:
4.35% 50Cr (49.9461 amu); 83.79% 52Cr (51.9405 amu);
9.50% 53Cr (52.9406 amu); 2.36% 54Cr (53.9389 amu)
49.9461
51.9405
52.9406
53.9389
x
x
x
x
.0435
.8379
.0950
.0236
= 2.17 amu
= 43.52 amu
= 5.03 amu
=+ 1.27 amu
51.99 amu
54
A2. Calculate non-integer relative atomic
masses and abundance of isotopes from given
data.
Determine the atomic weight of lead using
the data from the mass spectrum of lead
•
•
•
•
Four isotopes: 204, 206, 207, 208
Percentage of each isotope:
Total # isotopes is 10 (1+2+2+5)
204: 1/10 = 10% 206: 2/10 = 20%
207: 2/10 = 20% 208: 5/10 = 50 %
Multiply the percent of each isotope by its
mass
204 x .1 = 20.4 206 x .2 = 41.2
207 x .2 = 41.4 208 x .5 = 104
Add these values
20.4 + 41.2 + 41.4 + 104 = 207
Mass Spectrum of Lead
55
A2. Calculate non-integer relative atomic
masses and abundance of isotopes from given
data.
The atomic weight of gallium is 69.72 amu. The masses of the naturally
occurring isotopes are 68.9257 amu for 69Ga and 70.9249 amu for 71Ga.
Calculate the percent abundance of each isotope.
•
Let x = % abundance of
69Ga.
Then 1-x = % abundance of
•
68.9257x + 70.9249(1-x) = 69.72 amu
68.9257x + 70.9249 – 70.9249x = 69.72
-1.9992x = -1.20
x = 0.600 = decimal value of 69Ga so
60.0% 69Ga
1-x = 0.400 = decimal value 71Ga so
40.0 % 71Ga
71Ga.
56
A2. Calculate non-integer relative atomic
masses and abundance of isotopes from given
data.
The atomic weight of copper is 63.546 amu. The masses of the two naturally
occurring isotopes are 62.9298 amu for 63Cu and 64.9278 amu for 65Cu.
Calculate the percent of 63Cu in naturally occurring copper.
57
A2. Calculate non-integer relative atomic
masses and abundance of isotopes from given
data.
The atomic weight of copper is 63.546 amu. The masses of the two naturally
occurring isotopes are 62.9298 amu for 63Cu and 64.9278 amu for 65Cu.
Calculate the percent of 63Cu in naturally occurring copper.
•
Let x = % abundance of
63Cu.
Then 1-x = % abundance of
•
62.9298x + 64.9278(1-x) = 63.546 amu
62.9298x + 64.9278 – 64.9278x = 63.546
-1.998x = -1.382
x = 0.6917 = decimal of 63Cu so
69.17%
65Cu.
63Cu
58