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Atomic Structure
GChem Chapter 3
Learning objectives
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Discuss the history of the current atomic model, including contributions of:
Dalton, Thomson, Rutherford, Moseley, Bohr.
Explain the laws of conservation of mass, definite proportions, and
multiple proportions.
Describe the location, mass and charge of the three components of an
atom.
Explain and calculate the atomic mass, atomic number and charge for any
given atom.
Explain the concept of isotopes and calculate the average atomic mass for
an element given natural abundances.
Write the AZX symbol for any given isotope.
Discuss the development of the quantum mechanical model of the atom.
Define the 4 quantum numbers.
Write the electron configuration and orbital diagram for any given element
or ion.
Relate electron configuration to the arrangement of the periodic table.
The story of mankind’s search for the
basic structure of matter stretches
throughout the entire history of the
human race.
People have always wondered about
what things are made of and what
makes one substance different from
another.
The history of the atomic model
Ancient times- Around 400 BC, Greek
philosophers developed the idea that all
matter must be composed of tiny, indivisible
particles which were termed “atoms”. They
suggested that atoms of different substances
were different and that atoms “hooked
together” to form large scale matter.
Democritus is usually credited with the
development of this idea.
The next 2208 years – Not much
changed.
Throughout most of recorded
human history, the Greek’s idea
of “atoms” as the basis of
matter was accepted but
without any real experimental
evidence.
Although much study was done
during this time, little of it was
rigorous and much was given
mystical explanations – This was
Alchemy. (the emerald tablet)
~ 1800 – More rigorous experimentation was done by
various scientists. The results of these experiments
were formulated into 3 laws:
The law of conservation of mass - Mass is not gained
or lost in chemical reactions.
Total mass of reactants = Total mass
of products
The law of definite proportions - A compound will
always be composed of the same percentage of
each element by weight.
Examples:
100g of NaCl → 39.4g Na and 60.6g Cl
10g of NaCl → 3.94g Na and 6.06g Cl
CO2 will always be composed of 72.7% O and 27.3% C.
The law of multiple proportions
If two compounds contain the same elements, a
comparison of the mass of one element that
reacts with a fixed mass of the other element
will give a factor of a small whole number.
Huhh??
Examples:
H2O → 1g H and 7.94g O
H2O2 → 1g H and 15.88g O
A 2 to 1 ratio!!!!
Compound
%O
%N
gO / gN
Ratio
N2O
36.35
63.65
0.571
1
NO
53.32
46.68
1.142
2
NO2
69.56
30.44
2.284
4
The conclusion was that elements exist in
“chunks” of mass. That for a given element, all
chunks have the same mass.
John Dalton used the 3 laws to develop his
atomic theory in 1808. This was the first theory
that described the composition of matter based
on experimental evidence.
Dalton’s Theory
1. Matter is made of indestructible atoms.
2. All atoms of an element are identical.
3. Atoms of different elements have different
properties.
4. Atoms of different elements combine with
each other in whole number ratios.
5. Chemical reactions are re-arrangements of
atoms.
Dalton’s Model – Billiard Ball Model
• J.J. Thomson- In a series of experiments with
a cathode ray tube (CRT) in 1897, discovered
that negatively charged particles of matter
could be removed from atoms.
• This indicated that atoms were not indivisible
but were composed of even smaller particles.
• The discovered particle was the
electron.
The Plum Pudding Model - JJ Thomson 1904
Ernest Rutherford – 1911 – Proposed a new
model of the atom based on the results of the
Gold Foil experiment.
http://www.mhhe.com/physsci/chemistry/essen
tialchemistry/flash/ruther14.swf
Rutherford Model of the atom
• Rutherford suggested that the atom is
composed of a very small, dense, positively
charged nucleus surrounded by an area of
empty space containing the atom’s electrons.
Henry Moseley – 1913 – Discovered that the number
of positive charges in an atom is equal to the
element number. This indicated that there was a
particle in the nucleus that was the source of +
charge. In 1920, Rutherford named the particle
“proton”.
Niels Bohr – 1913 – Considered that the Rutherford
model of the atom was unstable and the spectra of
atoms (discrete bands of light absorbed and
emitted by atoms) to propose a new model with the
electrons confined to specific energy levels
(sometimes called shells or orbits).
The Bohr Model
• e- exist in specific energy levels in atoms and cannot
exist between energy levels – The levels are
quantized.
• e- can absorb specific amount of energy to jump to
a higher energy level or release energy to drop to a
lower level – quantum jumps.
• e- in lowest energy level = ground state
e- jumps to higher energy level = excited state
• Specific numbers of e- can reside in each energy
level.
• http://science.sbcc.edu/physics/flash/sili
consolarcell/bohratom.swf
• James Chadwick – 1932 – Explained the
difference between the observed mass of
atomic nuclei and the number of + charges
(also considering spin) by proposing the
presence of particles with masses similar to
those of protons but with no charge the neutron.
We now have a workable model of the atom: It is composed
of three particles, proton, neutron and electron.
Name
Symbol
Mass
Charge Location
Proton
p+
1.67x10-27 kg
+
nucleus
Neutron
n0
1.67x10-27 kg
None
Nucleus
Electron
e-
9.1x10-31 kg
-
Energy
levels
The mass of a neutron is actually very slightly more
than that of a proton however, in chemistry we
generally consider them to be the same.
We use a convenient unit to express this mass, the
amu (atomic mass unit).
An amu is defined as 1/12 the mass of a carbon-12
atom.
We generally consider the masses of both p+ and n0 to
be 1 amu.
The mass of an e- is so much less than the other
particles that we considered it to be zero in calculating
the mass of an atom.
So the mass of an atom, in amu’s, is simply the
number of protons plus the number of
neutrons. This is sometimes called “mass
number”
atomic mass = #p+ + #n0
The total charge on an atom is determined by
the number of p+ and e-. Since these particles
have charges of equal magnitude and
opposite sign, their charges cancel. When an
atom has the same number of p+ and e- , the
total charge must be zero – a neutral atom.
• An ION is a form of an atom where the number of edoes not match the number of p+.
• If the ion has more e- than p+ it will have a total
negative charge.
• If the ion has less e- than p+ it will have a total
positive charge.
Ex: O ion has 8 p+ and 10 e- . 2 more electrons than
protons so the ion will have a charge of -2, (O-2)
Mg ion has 12 p+ and 10 e- . 2 more protons than
electrons so the ion will have a charge of +2, (Mg+2)
Br ion has 35 p+ and 36 e- . 1 more electrons than
protons so the ion will have a charge of -1, (Br-)
La ion has 57 p+ and 54 e- . 3 more protons than
electrons so the ion will have a charge of +3, (La+3)
Ions can only form by the loss or gain of
electrons.
Since protons determine the identity of the
element, ions are never formed by losing or
gaining p+.
This should make sense – Electrons are flying
around the outer part of the atom – easy to
gain or lose. Protons are locked into the
nucleus at the center – cannot gain or lose.
•
•
•
•
Writing symbols for atoms – the AZX method
X – the chemical symbol for the element, H, O, Ca
A – the mass number (atomic mass)
Z – the atomic number (number of p+)
•
Additionally, the charge on an ion can be written
in the upper right!
Examples: Write the AZX notation for the following:
1. 20 p+, 20 n0, 20 e2. 15 p+, 16 n0, 15 e-
3. 26 p+, 30 n0, 23 e4. 35 p+, 44 n0, 36 e-
• Summary:
1. the number of protons is the atomic number and tells
you what element you have.
13 p+ = Al
2. The number of electrons can be compared to number of
protons to tell you if you have a neutral atom or an ion.
13 p+ and 10 e- = Al+3
3. The number of neutrons is added to number of protons
to give the mass number (atomic mass).
27 +3
+
0
13 p and 10 e and 14 n = 13 Al
• The number of neutrons that atoms of a given element
contains can vary.
• Differing numbers of neutrons do not change the identity
of the element, however they will change the mass of the
atom.
• Isotopes are versions of the same element with different
numbers of neutrons and therefore different atomic
masses.
Isotope
p+
n0
Atomic mass
C-12
6
6
12
C-14
6
8
14
Mg-24
12
12
24
Mg-25
12
13
25
U-235
92
143
235
U-238
92
146
238
Isotopes of hydrogen
Practice:
Write the AZX symbols for the isotope with:
21 p+ and 24 n0
17 p+ and 20 n0
82 p+ and 125 n0
41 p+ and 52 n0
The atomic mass for an element given on the periodic
table is the weighted average of the masses of all
the naturally occurring isotopes for that element.
The average takes into account both the mass of the
isotope and it’s natural abundance (what
percentage of it occurs in nature).
• Pb has 4 naturally occurring isotopes:
• Atomic Number = 82 so 82 p+ in each
• Pb-204 204-82 = 122 no
• Pb-206 206-82 = 124 no
• Pb-207 207-82 = 125 no
• Pb-208 208-82 = 126 no
Isotope
Natural Abundance
Pb-204
1.4%
Pb-206
24.1%
Pb-207
22.1%
Pb-208
52.4%
To calculate Avg At Mass: A weighted average!
204 x 0.014 = 2.856
multiply the isotope mass
206 x 0.241 = 49.65
by it’s natural abundance
207 x 0.221 = 45.75
(move the decimal point 2
208 x 0.524 =108.99
places)
207.25
add the results!
• Average atomic mass practice:
Uranium has three common isotopes. If the
abundance of 234U is 0.01%, the abundance of 235U
is 0.71%, and the abundance of 238U is 99.28%,
what is the average atomic mass of uranium?
Titanium has five common isotopes: 46Ti (8.0%), 47Ti
(7.8%), 48Ti (73.4%), 49Ti (5.5%), 50Ti (5.3%). What is
the average atomic mass of titanium?
Electron Configuration
Quantum Mechanical View
The Bohr model is very useful but has
real limitations and in some ways does
not agree with observations.
Louis DeBroglie : In 1924 proposed a new model to
explain the problems with Bohr’s model.
Suggested that electrons could be
considered as particles or waves –
just as Einstein had proposed about
light.
Only certain “orbits” are possible
because these correspond to
multiples of the wavelength for the
electron. The orbit can be thought of
as a standing wave.
Erwin Schrodinger: In 1926, derived an equation
that described the position of an electron as a
probability. The equation can be used to
produce a set of four quantum numbers for
each electron that describes its probable
location in an atom.
Wolfgang Pauli – 1925 – The Pauli exclusion
principle states that no two electrons in an
atom can have the same set of 4 quantum
numbers.
• Werner Heisenberg – 1927. The Heisenberg
uncertainty principle states that the
uncertainties of the position and momentum
of an electron are inversely proportional.
The more accurately that one is known, the
less accurately the other can be known.
The Cat Hotel
Rule 1: Cats are lazy
Rule 2: Cats do not like
each other
5th
4th
3rd
2nd
1st
• A set of 4 quantum numbers for an electron in
an atom gives the “address” for that electron.
• Where an e- resides in an atom is called an
orbital. Each energy level contains 1 or more
orbitals.
• The quantum numbers describe the energy
level, the shape and orientation of the orbital
and the spin of the e-.
L3
L2
L1
• Quantum Numbers are used to describe the location of an
electron in an atom.
• Four quantum numbers are needed for each electron and
no electrons in an atom can have the same set on QN’s.
• The principal QN is is identified by the letter n and gives the
main energy level of the electron.
• The principal QN can have the values 1,2,3,4,5……
• The second QN is identified by the letter l and gives the
shape of the electron orbital.
• The l QN can have the values 0,1,2,3…n-1
• The m l QN gives the orientation of the orbitals and can
have the values - l …-3,-2,-1,0,1,2,3…+ l
• The ms QN gives the spin of an electron and has the values
of +1/2 or -1/2.
For N=1 l = _____ This is an _____ orbital which holds _____ eFor N=2 l = _____ This is an _____ orbital which holds _____ eAnd l = _____ These are _____ orbitals which hold _____ eFor N=3 l = _____ This is an ______ orbital which holds _____ e-
And l = _____ These are _____ orbitals which hold _____ eAnd l = _____ These are _____ orbital which hold _____ eFor N=4 l = _____ This is an _____ orbital which holds _____ e-
And l = _____ These are ______ orbitals which hold ____ eAnd l = _____ These are ______ orbital which hold _____ eAnd l = _____ These are ______ orbitals which hold _____ e-
• Orbital diagrams:
1s
2s
2p
• Translates to electron
configurations
1s 2s 2p 3s 3p 4s 3d 4p 5s
3s
3p
•
To write orbital diagrams and electron
configurations:
1. Decide how many electrons the atom has. For
neutral atoms this is the same as the atomic
number.
For + ions, 1 e- is lost for each + charge.
For – ions 1 e- is gained for each - charge.
2. Starting with the 1st energy level (n=1) add 2 e- to
each orbital until all are used up. Remember, for
sets of orbitals (p,d,f) e- do not double up until
they have to.
3. Use a filling diagram to fill in the correct order.
Write the orbital diagram and e- configuration for the
following:
O
Si
Ti
Sn
Pm
Ions
Ground state versus excited state
Shortcut for writing e- configurations
Find the noble gas that has a lower atomic number
that is nearest to the element.
Write the symbol for the noble gas in brackets.
Write the remainder of the e- configuration between
the noble gas and the element:
Ex: Cd
1s22s22p63s23p64s23d104p65s24d10
The underlined is the configuration for Kr so we can
write:
[Kr]5s24d10
Additional Content
• Quarks!