Download I - Holland Public Schools

UNIT III - ATOMIC THEORY
1
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)
* if you have more protons, the ion is +, if you have less, the ion is - 1 proton + 1 electron = neutral charge (0)
* ion - an atom with an unbalanced number of electrons and protons; a charged particle
sodium atom (Na)
11 p+
12 n
o
sodium ion (Na+1)
12 p+
12 no
fluoride ion (F-1)
11 p+
loses one
9 p+
12 no
electron
magnesium atom (Mg)
fluorine atom (F)
10 n
magnesium ion (Mg+2)
loses two
electrons
9 p+
gains one
o
10 no
electron
oxygen atom (O)
oxide ion (O-2)
12 p+
12 no
8 p+
gains two
8 no
electrons
8 p+
8 no
S. What is the charge on an atom containing:
+1 1)
11p+, 12no, 10e-2 2) 16p+, 16no, 18e+2 3) 25p+, 30no, 23e* 11p – 10e = +1 charge
* 16p – 18e = -2 charge
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
Na 1)  sodium’s atomic
S 2)
number is 11 (11p)
D. Mass Number - defined as the total number of protons and neutrons
- Warning!!! : do not confuse this with atomic mass
Mn 3)
S. Using the three examples above, give the mass number
23 1)
 11p + 12n = 23
32 2)
55 3)
UNIT III - ATOMIC THEORY
2
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 = 12 , neutrons = 12
* the “-24” is the mass number
Symbol:
24
12
Mg
magnesium-25:
Symbol:
25
12
Mg
mass number = 25, protons = 12 , neutrons = 13
* Electrons occur in Energy Levels (areas which can hold up to a certain number of electrons)
Energy Level # of electrons
1
2
2
8
3
8
S. Write the correct symbol and draw a picture of the following atoms:
* Also label each as an atom or ion
1) hydrogen-1
1p, 0n, 1e
2) carbon-14
6p, 8n, 6e
1 p+
0n
1)
3)
18
8
O 2
4)
8p, 10n, 10e
6 p+
8 p+
8 no
10 no
7
3
Li 1
3p, 4n, 2e
3 p+
o
4 no
atom
atom
ion (anion)
S. Give the full symbol for each of the following elements:
9
4
2)
Be
11
5
B
4 p+
5 p+
5 no
6 no
3)
4
2
He
4)
24
12
ion (cation)
Mg 2
12 p+
2 p+
2n
o
12 no
UNIT III - ATOMIC THEORY
3
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? as solid, uncharged spheres:
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
the electrons always had their
paths bent toward the positive
charge; therefore they must be
negatively-charged particles
Positively-charged plate
+
+
-
-
+
+
Negatively-charged plate
-
* Drawing: Show Thomson’s “plum pudding” model
C. Ernst Rutherford (1909)
* performed experiment with alpha particles - apparatus
gold foil
note that most
particles
went
straight through
the foil, but a few
were deflected
radioactive source
alpha particles: 4+
-
zinc sulfide screen
+
UNIT III - ATOMIC THEORY
4
* What were the results of his experiment?
- most particles went straight through (see left)
- a few were deflected
* Why did these results contradict Thomson’s model?
- if the solid atom model were correct, the alpha
particles would have all bounced back off
the foil (atoms would be like gray circles at
left)
* 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?
He didn’t have a good answer for this
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:
1st Energy Level
Nucleus
2nd Energy Level
3rd Energy Level
* 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:
UNIT III - ATOMIC THEORY
5
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
35
37
nature and find that it contains 75.53% 17 Cl and 24.47% 17 Cl . What is the atomic mass of
chlorine?  simply make the percentages into decimals, and multiply each percentage by its
respective mass number; total the result
(0.7553)(35) + (0.2447)(37) = 35.4894
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?
(0.2534)(79) + (0.5000)(80) + (0.2466)(81) = 79.9932
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?
(0.9976)(16) + (0.0004)(17) + (0.0020)(18) = 16.0044
UNIT III - ATOMIC THEORY
6
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 – energy produced by the motion of any magnetic or
charged particle
- the energy of a light beam is directly proportional to the frequency
- color of light depends on frequency
- visible light – ranges from 4.3 x 1014 – 7.5 x 1014 Hz
- What color light has the highest energy? lowest? Highest E – violet;
Lowest E - red
* each element releases a bright-line spectra when its atoms are excited
* the colored lines produced by a bright-line spectra directly correspond to energy “jumps”
for electrons in a hydrogen atom as they travel from an excited state back to a ground state
* the color of the light translates to a calculable amount of energy (in Joules)
* this phenomenon is explained by Planck’s Quantum Theory which led to Bohr’s model
(see p. 4)
UNIT III - ATOMIC THEORY
7
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
n=4
4d
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
UNIT III - ATOMIC THEORY
8
* OK, let’s build some atoms – these orbital diagrams below are a visual way of representing
the energy states of electrons in an atom. Each circle represents an orbital, each
slash mark represents an electron
* Rules:
1) Electrons want to be in the lowest possible E level available.
2) Orbitals hold 2 electrons.
3) We are only dealing with neutral, ground state atoms. – this means that the number
of electrons is the same as the atomic number (number of protons)
4) Pauli Exclusion Principle - 2 electrons in the same orbital must have opposite spin
5) Hund’s Rule - electrons want to spread out as much as possible within a sublevel.
Electrons
hydrogen
1
helium
2
lithium
3
beryllium
4
boron
5
carbon
6
nitrogen
7
oxygen
8
fluorine
9
neon
10
1s
2s
2p
3s
 note the
Pauli
exclusion
principle
 note
Hund’s
Rule here
UNIT III - ATOMIC THEORY
9
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: Method: beginning at Hydrogen (#1), you count
up in atomic number order until you arrive at the square belonging to the desired element, keeping track
of all the areas you go through on the way. Each square represents one electron in that section
He
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6s
5d
6p
7s
6d
7p
4f
5f
T. Give the electron configuration for:
1) Hydrogen
1s1 – 1st square in the 1s section
2) Helium
1s2 – 2nd square in the 1s section
3) Lithium
1s22s1 – 2 squares in the 1s section + 1 square in the 2s section
4) Carbon
1s22s22p2
5) Fluorine
1s22s22p5
UNIT III - ATOMIC THEORY
10
S. Give the electron configuration for:
1) Boron
1s22s22p1
2) Cobalt
1s22s22p63s23p64s23d7
3) Calcium
1s22s22p63s23p64s2
4) Aluminum
1s22s22p63s23p1
5) Chlorine
1s22s22p63s23p5
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
Method: work backward on the periodic table from the desired element to a noble gas; put the symbol
of the noble gas in brackets, then do the configuration up from there
T.
1) Aluminum
[Ne]3s23p1
2) Bromine
[Ar]4s23d104p5
3) Potassium
[Ar]4s1
4) Promethium
[Xe]6s24f5
5) Osmium
[Xe]6s24f145d6
UNIT III - ATOMIC THEORY
11
S.
1) Silicon
[Ne]3s23p2
2) Krypton
[Ar]4s23d104p6
3) Vanadium
[Ar]4s23d3
4) Holmium
[Xe]6s24f11
5) Mercury
[Xe]6s24f145d10
* There is one set of strange ones: columns 6 and 11
- half-filled orbitals an arrangement like:
which is much more favorable than having a “lopsided”
arrangement, so the elements below actually “steal” an
electron from an s orbital (thus making that half-full), and
placing it in the adjacent d orbital
T.
1) Chromium
2) Copper
normally:
[Ar]4s23d4
becomes:
[Ar]4s13d5
[Ar]4s13d10
S.
1) Molybdenum
[Kr]5s14d5
2) Silver
[Kr]5s14d10
* NOTE: Gold (Au) does this pattern also, but tungsten (W) does not
UNIT III - ATOMIC THEORY
12
VIII. Orbital Shapes – derived from Schrödinger’s Equations; represent 3-dimensional areas in which
you are 90-95% likely to find an electron of a particular energy state
A. s-orbitals
1s
2s
3s
B. p-orbitals
pz
py
px
C. d-orbitals
dx2 - y2
dxy
dxz
D. f-orbitals  How many lobes would you guess they have?
dyz
8
dz2