Download General Chemistry Chapter 3 Note Packet

Name:___________________________
General Chemistry Chapter 3 Note Packet
Learning objectives
1. Discuss the history of the current atomic model, including
contributions of:
Dalton, Thomson, Rutherford, Moseley, Bohr.
2. Explain the laws of conservation of mass, definite proportions,
and multiple proportions.
3. Describe the location, mass and charge of the three
components of an atom.
4. Explain and calculate the atomic mass, atomic number and
charge for any given atom.
5. Explain the concept of isotopes and calculate the average
atomic mass for an element given natural abundances.
6. Write the AZX symbol for any given isotope.
7. Discuss the development of the quantum mechanical model of
the atom.
8. Define the 4 quantum numbers.
9. Write the electron configuration and orbital diagram for any
given element or ion.
10. Relate electron configuration to the arrangement of the
periodic table.
The history of the atomic model
Ancient times- Around 400 BC, Greek philosophers
developed the idea that all ________ must be composed of
tiny, ___________ particles which were termed “atoms”.
They suggested that atoms of different substances were
_________________ and that atoms “hooked together” to
form large scale matter.
Democritus is usually credited with the development of this
idea.
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Name:___________________________
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 _____________________________.
Although much study was done during this time, little of it
was rigorous and much was given mystical explanations –
This was _____________________.
~ 1800 – More rigorous ___________________ 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 ______________________.
Total mass of reactants = Total mass of products
The law of definite proportions - A compound will always
be composed of the _____________________ of each
element by weight.
Examples:
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Name:___________________________
The law of multiple proportions
If two compounds contain the same ______________, 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
_____________________.
Example:
Compound % O
%N
N2O
gO / gN
Ratio
1
NO
NO2
John Dalton used the 3 laws to develop his atomic theory in
_________. This was the first theory that described the
composition of matter based on ______________________.
1. Matter is made of _________________ atoms.
2. All atoms of an element are _________________.
3. Atoms of different elements have different ____________.
4. Atoms of different ____________ combine with each
other in whole number ratios.
5. Chemical reactions are ___________________ of atoms.
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Name:___________________________
J.J. Thomson- In a series of experiments with a _________
ray tube (CRT) in 1897, discovered that _____________
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 ____________.
Drawing of Thomson’s experiment:
Drawing of Thomson’s model (Plum pudding model)
Ernest Rutherford – 1911 – Proposed a new model of the
atom based on the results of the ____________ experiment.
Rutherford suggested that the atom is composed of a very
small, __________, positively charged nucleus surrounded
by an area of __________________ containing the atom’s
electrons.
Drawing of the gold foil experiment:
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Name:___________________________
Henry Moseley – 1913 – Discovered that the number of
________________ 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 ______________.
Niels Bohr – 1913 – Considered that the Rutherford model
of the atom was______________ 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
____________________ (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 ______________.
e- can ______________ specific amount of energy to jump
to a higher energy level or ___________ energy to drop to a
lower level – quantum jumps.
e- in ____________ energy level = ground state
e- jumps to __________ energy level = excited state
Specific numbers of e- can reside in each energy level.
Drawing of Bohr Model:
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Name:___________________________
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 - ________________.
Name
Symbol
Mass
Charge
Location
Proton
Neutron
Electron
The mass of a neutron is actually very slightly _________
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
(_______________________).
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 _________ than the
other particles that we considered it to be __________ in
calculating the mass of an atom.
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Name:___________________________
So the mass of an atom, in amu’s, is simply the number of
protons plus the number of neutrons. This is sometimes
called “__________________”
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 _________
magnitude and _____________ 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
_________________ charge.
• If the ion has less e- than p+ it will have a total
________________ charge.
Ex: O ion has 8 p+ and 10 e- . 2 more electrons than protons
so the ion will have a charge of ________, (O-2)
Mg ion has 12 p+ and 10 e- . 2 more protons than electrons
so the ion will have a charge of ________, (Mg+2)
Br ion has 35 p+ and 36 e- . 1 more electrons than protons
so the ion will have a charge of _________, (Br-)
La ion has 57 p+ and 54 e- . 3 more protons than electrons
so the ion will have a charge of __________, (La+3)
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Name:___________________________
Ions can only form by the loss or gain of _____________.
Since protons determine the identity of the element, ions are
never formed by losing or gaining p+.
Writing symbols for atoms – the AZX method
• X – the _________________ for the element, H, O, Ca
• A – the _________________ (atomic mass)
• Z – the _________________ (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 e3. 26 p+, 30 n0, 23 e4. 35 p+, 44 n0, 36 e-
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Name:___________________________
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).
272
13 p+ and 10 e- and 14 n013 =
27
13
Al+3
The number of ____________ that atoms of a given element
contains can vary.
Differing numbers of neutrons do not change the _________
of the element, however they will change the _________ of
the atom.
Isotopes are versions of the same element with different
numbers of ___________ and therefore different
__________________.
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Name:___________________________
Isotope
p+
n0
C-12
6
6
C-14
6
8
Mg-24
12
12
Mg-25
12
13
U-235
92
143
U-238
92
146
Atomic mass
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
_______________ is the weighted average of the masses of
all the naturally occurring ______________ for that element.
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Name:___________________________
The average takes into account both the mass of the isotope
and it’s __________________ (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
Pb-204
Pb-206
Pb-207
Pb-208
Natural Abundance
1.4%
24.1%
22.1%
52.4%
To calculate Average Atomic Mass: A weighted average!
 multiply the isotope mass by it’s natural abundance
 make sure to move the decimal point 2 places
 do this for each of the isotopes
 add the results
204 x 0.014 =
206 x 0.241 =
207 x 0.221 =
208 x 0.524 =
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Name:___________________________
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.
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Name:___________________________
Louis DeBroglie : In 1924 proposed a new model to explain
the problems with Bohr’s model. Suggested that _________
could be considered as particles or waves – just as Einstein
had proposed about ______________
Only certain “orbits” are possible because
these correspond to multiples of the
wavelength for the electron. The orbit can be
thought of as a _________________ wave.
Erwin Schrodinger: In 1926, derived an equation that
described the position of an electron as a ______________.
The equation can be used to produce a set of four quantum
numbers for each electron that describes its
probable _____________________ in an
atom.
Wolfgang Pauli – 1925 – The Pauli
exclusion principle states that no two _____________ in an
atom can have the same set of ______ quantum numbers.
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Name:___________________________
Werner Heisenberg – 1927. The Heisenberg uncertainty
principle states that the _______________ of the position
and momentum of an ___________ 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 ___________
Rule 2: Cats do not like ______________
5th
4th
3rd
2nd
1st
A set of 4 ____________________ 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 _____________.
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Name:___________________________
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 of QN’s.
The principal QN is is identified by the letter n and gives the
___________________ of the electron.
The principal QN can have the values ________________
The second QN is identified by the letter l and gives the
__________________________________.
The l QN can have the values ______________________
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Name:___________________________
The ml QN gives the __________________ of the orbitals
and can have the values _________________________
The ms QN gives the ___________ of an electron and has
the values of _________________
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 _____ eAnd l = _____ These are _____ orbitals which hold _____ eAnd l = _____ These are _____ orbital which hold _____ eFor N=4 l = _____ This is an _____ orbital which holds _____ eAnd l = _____ These are ______ orbitals which hold ____ eAnd l = _____ These are ______ orbital which hold _____ eAnd l = _____ These are ______ orbitals which hold ____ e-
1s
2s
2p
3s
3p
Translates to electron configurations
1s 2s 2p 3s 3p 4s 3d 4p 5s
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Name:___________________________
To write orbital diagrams and electron configurations:
1. Decide how many electrons the atom has. For neutral
atoms this is the same as the ________________.
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.
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Name:___________________________
Practice: Write the orbital diagram and e- configuration for
the following:
O
Si
Ti
Sn
Pm
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Name:___________________________
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
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