Download unit iii - atomic theory 1

UNIT III - ATOMIC THEORY
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I. Atomic Basics
* atom - from Greek atomos, meaning “indivisible”; first coined by Democritus (300-200 BC)
A. Parts of the atom - subatomic particles
Particle
proton (p+)
neutron (no)
electron (e-)
Approximate Mass (amu)
1
1
1/1837
Charge
+1
0
-1
Location
nucleus
nucleus
outside nucleus
* 1 atomic mass unit (u) = 1.66 x 10-27 kg
B. Relationship to charge - found by balance of charged particles (protons and electrons)
- 1 proton + 1 electron = neutral charge (0)
* ion - an atom with an unbalanced number of electrons and protons; a charged particle
S. What is the charge on an atom containing:
____1) 11p+, 12no, 10e-
____2) 16p+, 16no, 18e-
____3) 25p+, 30no, 23e-
C. Atomic Number - defined as the number of protons in an atom of an element
- used to identify elements
S. Using the three examples above, identify what elements they are
____1)
____2)
____3)
D. Mass Number - defined as the total number of protons and neutrons
- Warning!!! : do not confuse this with atomic mass
____1)
S. Using the three examples above, give the mass number
____2)
____3)
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E. Isotope - most elements have two or more different forms with different mass numbers
(therefore, different amounts of neutrons), these different forms are called
isotopes of that element
- when dealing with samples of an element containing two or more different isotopes,
the mass number must be indicated with the element’s name:
- another way to represent different isotopes is to put the mass number on the top-left
of the symbol for that element:
magnesium-24 : mass number = 24, protons =
, neutrons =
Symbol:
magnesium-25:
, neutrons =
Symbol:
mass number = 25, protons =
* Electrons occur in Energy Levels (areas which can hold up to a certain number of electrons)
S. Write the correct symbol and draw a picture of the following atoms:
* Also label each as an atom or ion
1) hydrogen-1
2) carbon-14
3)
4)
S. Give the full symbol for each of the following elements:
1)
2)
3)
4)
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II. History of Atomic Theory - Evolution of the Atomic Model
A. John Dalton (1803) - first resurrected term “atom”
* What was known at this time?
1) Law of Constant Composition – a given compound always contains the same
proportion of elements by mass
* example: water is always composed of 88.9% O & 11.1% H
2) Law of Conservation of Mass
* 4 parts to his theory:
1) All matter is composed of indivisible, indestructible atoms.
2) All atoms of the same element are similar.
3) All atoms of different elements are different.
4) A compound is a chemical combination of two or more atoms.
* Drawing: How did Dalton “see” atoms?
B. J.J. Thomson (1897)
* Crookes (1870’s) – developed the Cathode Ray Tube
* Thomson found cathode rays were bent by electrical and magnetic fields
- called them electrons
* Drawing: Show Thomson’s “plum pudding” model
C. Ernst Rutherford (1909)
* performed experiment with alpha particles - apparatus
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* What were the results of his experiment?
* Why did these results contradict Thomson’s model?
* 2 conclusions:
1) Atoms are mostly empty space.
2) In the center of the atom is a densely-packed,
positively-charged nucleus.
*NOTE: no protons or neutrons yet
* Problem with Rutherford’s Model – Where are the electrons?
D. Niels Bohr (1913) - electrons are in energy levels; based his concept on Quantam Theory and
bright-line spectra for hydrogen
1) Quantam Theory – Max Planck (1900) – an object emits energy in specific quanta
(packets) which correspond to their energy
2) Photoelectric Effect – Einstein (1905); purple light caused release of electrons from
sodium metal, but red did not  light is emitted in quanta – photons
* purple photons have greater energy than red ones, therefore electrons escape
3) Bohr’s Model:
* ground state – when all electrons are in their lowest possible energy state
* excited state – when one or more electrons absorb a quantum of energy, it
“jumps” to a higher energy state; in order to return to the
ground state, it emits a specific amount of energy
(therefore a specific wavelength)
* problem with Bohr model  only works for hydrogen (one electron)
E. Charge-Cloud Model - also called orbital model and quantam-mechanical model
1) Matter as Waves – DeBroglie (1924)
* electrons have wavelengths and frequencies like light
* so energy has particle property (quanta), and matter has wave property
2) Heisenberg’s Uncertainty Principle – (1927)
* the exact position and velocity of an object can not be known simultaneously
3) Quantam-Mechanical Model
* basically says electrons are somewhere outside nucleus, not in neat orbits
- this “area” is known as an electron cloud
* orbital - area around nucleus which contains 2 electrons of opposite spin
* Drawing:
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III. Atomic Mass - weighted average of all isotopes of a certain element
T. Chlorine comes in two isotopes: Chlorine-35 and Chlorine-37. You take a sample of chlorine in
nature and find that it contains 75.53%
chlorine?
and 24.47%
S.
1) Bromine occurs in the following proportions:
Bromine-79
Bromine-80
Bromine-81
25.34%
50.00%
24.66%
What is the atomic mass of bromine?
2) Oxygen occurs in the following proportions:
oxygen-16
oxygen-17
oxygen-18
99.76%
0.04%
0.20%
What is the atomic mass of oxygen?
. What is the atomic mass of
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IV. The Nature of Light - behaves as both a wave and particle - we’ll deal with it as a wave:
wavelength
amplitude
peak
* frequency – number of peaks that pass a fixed point per second; Energy is directly
proportional to frequency (i.e. the higher the frequency of the radiation, the more energy it has)
* electromagnetic radiation
- the energy of a light beam is directly proportional to the frequency
- color of light depends on frequency
- visible light
- What color light has the highest energy? lowest?
* each element releases a bright-line spectra when its atoms are excited
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V. Orbital Diagrams - represent where electrons reside
* Each energy level is assigned a principal quantam number (n)
* Each energy level subdivides into a number of sublevels equivalent to “n”:
Principal
Energy Level (n)
Sublevels Available
1
2
3
4
5
6
1s
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f
5s, 5p, 5d, 5f, 5g
6s, 6p, 6d, 6f, 6g, 6h
* Each sublevel has a corresponding energy value:
4f
4d
n=4
4p
3d
4s
n=3
3p
3s
2p
n=2
2s
n=1
1s
* Each sublevel contains a certain number of orbitals:
Principal
Energy
Level (n)
s
Number of sublevels available
p
d
f
Total
Number of
Orbitals
Total
Number of
Electrons
1
2
3
4
1
1
1
1
-3
3
3
--5
5
---7
1
4
9
16
2
8
18
32
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* OK, let’s build some atoms
* Rules:
1)
2)
3)
4)
5)
T/S.
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Electrons want to be in the lowest possible E level available.
Orbitals hold 2 electrons.
We are only dealing with neutral, ground state atoms.
Pauli Exclusion Principle - 2 electrons in the same orbital must have opposite spin
Hund’s Rule - electrons want to spread out as much as possible within a sublevel.
1s
2s
2p
3s
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VI. Electron Configurations - a shorter way of representing electron locations
1s2
* pronounced: “one - s - two”
* label what each part represents
* easily done directly from the periodic table:
T. Give the electron configuration for:
1) Hydrogen
2) Helium
3) Lithium
4) Carbon
5) Fluorine
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S. Give the electron configuration for:
1) Boron
2) Cobalt
3) Calcium
4) Aluminum
5) Chlorine
VII. Advanced Electron Configurations
- silly to show lower patterns if they are always the same
- example: 1s22s22p6 is Neon. This never changes for any other element of higher number
- use [Ne] instead of 1s22s22p6
T.
1) Aluminum
2) Bromine
3) Potassium
4) Promethium
5) Osmium
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S.
1) Silicon
2) Krypton
3) Vanadium
4) Holmium
5) Mercury
* There is one set of strange ones: columns 6 and 11
- half-filled orbitals
T.
1) Chromium
2) Copper
S.
1) Molybdenum
2) Silver
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VIII. Orbital Shapes
A. s-orbitals
B. p-orbitals
C. d-orbitals
D. f-orbitals  How many lobes would you guess they have? _______