Name:___________________________
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 terme d “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.
Throu ghout 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 gO / gN Ratio
N
2
O
NO
NO
2
1
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 n 0 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 + + #n 0
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 e- does 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 n 0 , 20 e-
2. 15 p+, 16 n 0 , 15 e-
3. 26 p+, 30 n 0 , 23 e-
4. 35 p+, 44 n 0 , 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
13 p+ and 10 e- and 14 n =
13
27
13
Al +3
(atomic mass).
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
C-12
C-14
Mg-24 p+
6
6
12 n 0
6
8
12
Mg-25
U-235
12
92
13
143
U-238 92 146
Practice:
Write the AZX symbols for the isotope with:
21 p+ and 24 n 0
Atomic mass
17 p+ and 20 n 0
82 p+ and 125 n 0
41 p+ and 52 n 0
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 n o
• Pb-206 206-82 = 124 n o
• Pb-207 207-82 = 125 n o
• Pb-208 208-82 = 126 n o
Isotope
Pb-204
Pb-206
Natural Abundance
1.4%
24.1%
Pb-207
Pb-208
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?
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 Bo hr’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-.
L 3
L2
L 1
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 m l QN gives the __________________ of the orbitals and can have the values _________________________
The m s
QN gives the ___________ of an electron and has the values of _________________
For N=1 l
= _____ This is an _____ orbital which holds _____ e-
For N=2 l
= _____ This is an _____ orbital which holds _____ e-
And l = _____ These are _____ orbitals which hold _____ e-
For N=3 l
= _____ This is an ______ orbital which holds _____ e-
And l
= _____ These are _____ orbitals which hold _____ e-
And l
= _____ These are _____ orbital which hold _____ e-
For N=4 l
= _____ This is an _____ orbital which holds _____ e-
And l
= _____ These are ______ orbitals which hold ____ e-
And l
= _____ These are ______ orbital which hold _____ e-
And 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:___________________________
Sn
Pm
Ti
Practice: Write the orbital diagram and e- configuration for the following:
O
Si
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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
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10
The underlined is the configuration for Kr so we can write:
[Kr]5s 2 4d 10
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