I. Atomic Basics
* atom - from Greek atomos , meaning “indivisible”; first coined by Democritus (300-200 BC)
A. Parts of the atom - subatomic particles
Particle Approximate Mass (amu) Charge proton (p + ) neutron (n o )
1
1
+1
0 electron (e ) 1/1837 -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) sodium ion (Na +1 ) fluorine atom (F) fluoride ion (F -1 )
11 p +
12 n o loses one electron
11 p +
12 n o
9 p
+
10 n o gains one electron
9 p
+
10 n o magnesium atom (Mg) magnesium ion (Mg +2 ) oxygen atom (O) oxide ion (O
-2
)
12 p +
12 n o loses two electrons
12 p
+
12 n o
8 p +
8 n o gains two electrons
8 p +
8 n o
S . What is the charge on an atom containing:
+1
1) 11p + , 12n o , 10e -
-2
2) 16p + , 16n o , 18e -
+2
3) 25p + , 30n o , 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)

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 Symbol:
24
12
Mg
* the “-24” is the mass number magnesium-25: mass number = 25, protons = 12 , neutrons = 13 Symbol:
25
12
Mg
* 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 2) carbon-14 3)
18
8
O

2
1p, 0n, 1e 6p, 8n, 6e 8p, 10n, 10e
4)
7
3
Li

1
3p, 4n, 2e
1 p
+
0 n o
1) atom
4
9
Be
6 p
+
8 n o
8 p
+
10 n o
3 p
+
4 n o atom ion (anion)
S . Give the full symbol for each of the following elements: ion (cation)
2)
11
5
B 3)
2
4
He 4)
24
Mg

2
12
4 p
+
5 n o
5 p
+
6 n o
2 p
+
2 n o
12 p
+
12 n o

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) as solid, uncharged spheres:
* 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 radioactive source alpha particles:
4  + note that most particles went straight through the foil, but a few were deflected zinc sulfide screen

* 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
(packets) which correspond to their energy quanta
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:
Nucleus
1 st Energy Level
2 nd Energy Level
3 rd 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:

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% respective mass number; total the result
35
17
Cl and 24.47%
37
17
Cl . What is the atomic mass of chlorine?
 simply make the percentages into decimals, and multiply each percentage by its
(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

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 radia- tion, 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 10 14 – 7.5 x 10 14 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)

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”:
1 1s
2
3
4
5
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f
5s, 5p, 5d, 5f, 5g
6 6s, 6p, 6d, 6f, 6g, 6h
* Each sublevel has a corresponding energy value: n = 4 n = 3
4f
4d
4p
3d
4s
3p
3s n = 2
2p
2s
n = 1
* Each sublevel contains a certain number of orbitals:
1
2
3
4 s
1
1
1
1
p
--
3
3
3 d
--
--
5
5 f
--
--
--
7
1s
1
4
9
16
2
8
18
32

hydrogen
* 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.
1s 2s 2p 3s helium

note the
Pauli exclusion principle lithium
beryllium
boron
carbon nitrogen

note
Hund’s
Rule here oxygen
fluorine
neon

VI. Electron Configurations - a shorter way of representing electron locations
1s 2
* 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
3s
4s
5s
6s
7s
3d
4d
5d
6d
2p
3p
4p
5p
6p
7p
4f
5f
T . Give the electron configuration for:
1) Hydrogen
1s 1 –
1 st square in the 1s section
2) Helium
1s 2 –
2 nd square in the 1s section
3) Lithium
1s 2 2s 1 –
2 squares in the 1s section + 1 square in the 2s section
4) Carbon
1s 2 2s 2 2p 2
5) Fluorine
1s 2 2s 2 2p 5

S . Give the electron configuration for:
1) Boron
1s 2 2s 2 2p 1
2) Cobalt
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7
3) Calcium
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2
4) Aluminum
1s 2 2s 2 2p 6 3s 2 3p 1
5) Chlorine
1s 2 2s 2 2p 6 3s 2 3p 5
VII. Advanced Electron Configurations
- silly to show lower patterns if they are always the same
- example: 1s 2 2s 2 2p 6 is Neon. This never changes for any other element of higher number
- use [Ne] instead of 1s 2 2s 2 2p 6
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]3s 2 3p 1
2) Bromine
[Ar]4s 2 3d 10 4p 5
3) Potassium
[Ar]4s 1
4) Promethium
[Xe]6s 2 4f 5
5) Osmium
[Xe]6s 2 4f 14 5d 6

S .
1) Silicon
[Ne]3s 2 3p 2
2) Krypton
[Ar]4s 2 3d 10 4p 6
[Ar]4s 2 3d 3
3) Vanadium
[Xe]6s 2 4f 11
4) Holmium
5) Mercury
[Xe]6s 2 4f 14 5d 10
* There is one set of strange ones: columns 6 and 11
T .
- 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
S .
1) Chromium normally:
[Ar]4s 2 3d 4 becomes:
[Ar]4s 1 3d 5
2) Copper
[Ar]4s 1 3d 10
[Kr]5s 1 4d 5
1) Molybdenum
2) Silver
[Kr]5s 1 4d 10
* NOTE: Gold (Au) does this pattern also, but tungsten (W) does not

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
C. d-orbitals p x p y p z d xy d xz dyz d x2 - y2
D. f-orbitals

How many lobes would you guess they have?
8 d z2