Download Unit 2: Structure of Matter Content Outline: History of the Atomic

Unit 2: Structure of Matter
Content Outline: History of the Atomic Model (2.1)
I. Democritus (400 B. C.)
A. He was a Greek philosopher of science.
B. First to use the term “atom” to describe the basic particle of nature.
1. “atom” means “indivisible”
2. Atom – the smallest particle of an element that still retains the chemical properties of
that element.
II. John Dalton (1808)
A. He was an English schoolteacher.
B. He was the first to propose an “Atomic theory” that contains the 5 following statements:
1. All matter is composed of extremely small particles called “atoms”.
2. Atoms of a given element are identical in size, mass, and other properties; atoms of
different elements differ in size, mass, and other properties.
a. This has since been modified based on Isotopes and ions.
3. Atoms cannot be subdivided, created, or destroyed.
a. This has since been modified based upon current studies in quantum physics. Such
examples include muons and quarks.
4. Atoms of different elements combine in simple whole-number ratios to form chemical
compounds.
5. In chemical reactions, atoms are combined, separated, or rearranged.
III. Joseph John Thomson (1897)
A. He was an English Physicist.
B. He worked with glass gas-filled tubes referred to as Cathode-Ray tubes.
1. The glass tubes were filled with a gaseous element under low pressure.
2. He then passed an electrical current through the gas using a battery and wires.
a. The electrical current caused the gas within the tube to intensely glow with a beam
(“ray”).
i.
Magnets could make the “ray” move/deflect in various directions.
ii.
The ray is being deflected by the negative charge of the magnet.
iii.
Negative charge repels/deflects like negative charges.
iv.
The ray is made of a negative charge that Thompson called electrons (since
they were associated with the electrical current.)
b. The electrical current came into the chamber (by a wire) at the cathode end. (The
end where electricity enters the tube.)
c. The electrical current left the tube on the anode end. (The end where the electricity
goes back into the wire.)
d. Hence the term Cathode Ray tubes.
C. Further investigations using different elements in Cathode-Ray tubes confirmed that every
element’s atoms possess electrons.
D. He proposed the “Plum Pudding” model of atoms.
1. It stated that negatively charged electrons are evenly placed inside a positively charged
mass.
IV. Robert A. Milliken (1909)
A. He was an American Physicist.
B. He was the first to measure the charge and mass of an electron.
The symbol for an electron is: e-
1. Electron charge = 1.602 x 10-19 Coulombs.
a. This is an extremely small quantity of energy.
2. Electron mass = 9.11 x 10-31 kg
a. Electrons are 1/1837th the mass of a single proton or neutron.
b. This is a very, very, very small amount and size.
C. Milliken’s experiments allow for 2 inferences (conclusions based upon evidence and
reasoning) to be made:
1. Because atoms, in the natural state, are electrically neutral, they must also contain an
equal amount of positively charged particles.
2. Because electrons have so little mass, atoms must contain other particles with much
greater mass (protons & neutrons).
V. Ernest Rutherford, Hans Geiger, and Ernest Marsden (1911)
A. Geiger and Marsden were students of Rutherford a New Zealand Physicist.
B. They performed the Gold Foil Experiment.
C. They used high-energy alpha particle radiation (2 protons & 2 neutrons ejected from a
decomposing, radioactive element) to bombard a piece of gold foil that was surrounded by a
fluorescent screen.
1. As alpha particles struck the fluorescent screen, they would produce a small detectable
burst of light.
2. As the experiment was running, they detected light burst mainly behind the gold foil,
but also occasionally all around the ring.
a. These bursts of light around the ring were because of the positively charged alpha
particles been deflected by positively charged particles in the atoms of the foil.
b. The particles became known as protons.
i.
Just as with the electrons, positive charges repel/deflect like positive charges.
c. As most of the bursts of light occurred behind the gold foil, they concluded that the
majority of space in an atom is “empty space” that the alpha particles travelled
through and never hit anything.
D. Rutherford proposes the idea of the neutrally charged neutron particle in 1920.
VI. Niels Bohr (1913)
A. He was also a student of Rutherford’s.
B. He proposed the Bohr model of an atom.
1. The electrons move in a circular pattern around the positively charged center. (Much
like the planets revolve around the sun.)
VII. Dmitri Ivanenko & Victor Ambartsumian (1930)
A. These gentlemen were Russian Physicists.
B. They proposed a model of the nucleus of an atom that is composed of positively charged
protons and neutral charged particles (neutrons).
VIII.
James Chadwick (1932)
A. He was an English Physicist.
B. He proved that the nucleus is definitely composed of protons and neutrons through his
experiments with alpha particle radiation.
Content Outline: Basic Atomic Structure and Mass (2.2)
I.
Atom
A. The smallest particle of an element that still retains the chemical properties of that element.
B. Atoms are composed of 3 sub-atomic particles:
1. Electrons (Thomson proposed)
a. Electrons possess negative electrical charges.
b. Electrons are found orbiting the nucleus of an atom, in what is referred to as the electron
cloud. (They move at the speed of light and “create” a cloud-like appearance.)
c. Electrons are 1/1837th the mass of a single proton or single neutron.
2. Protons (Rutherford, Geiger, and Marsden proposed)
a. Protons possess a positive electrical charge.
b. Protons are found clumped together within the nucleus of an atom.
c. Each proton has a mass of 1 atomic mass unit (AMU) or 1 Dalton (Named after John Dalton.)
3. Neutrons (Rutherford proposed)
a. Neutrons possess no electrical charged and are therefore referred to as neutral.
b. Neutrons are also found clumped together within the nucleus of an atom.
c. Each neutron has a mass of 1 AMU or 1 Dalton.
C. Nuclear Forces
1. These are short-range proton-to-neutron OR proton-to-proton OR neutron-to-neutron
attractive forces that help hold together the nucleus of an atom.
2. These forces are greater than the repulsive same charge electrical forces exhibit by protons.
D. Atomic Radii
1. This term refers to the relative size of an individual atom of an element.
2. It is measured from the center of the nucleus to the outermost electron cloud.
3. It is measured in picometers (pm)
a. A picometer is 1.0 x 10-12 meters (so it is very, very small)
4. Charles-Augustin de Coulomb (1785)
a. He proposed Coulomb Forces – attractions that exist between oppositely electrically
charged particles (protons & electrons) within a single atom.
b. The forces directly affect the atomic radii of an atom.
i.
More protons than electrons = radii shrinking (getting smaller) because the positive
charge is greater than the smaller negative charges and pulls them in toward the
nucleus.
ii.
More electrons than protons = radii increases (getting larger) because the electrons
are farther away from the positive nucleus.
iii.
The Natural state of atoms has protons = electrons; so atomic radii are stable (not
changing) for each element.
5. Atomic radii can have an effect on the chemical properties of an element.
II. Atomic Number
A. This term refers to the number of protons found within the nucleus of an atom for that element.
B. Each element has a unique and identifying number of protons.
C. The atomic number for each element led to the creation of the Periodic Table
1. The Periodic Table was originally created by Dmitri Mendeleev in 1869.
a. He was a Russian Chemist.
D. The atomic number is usually written as superscript (above) the Elements Chemical symbol.
1. Some of the symbols use the Latin term, instead of the English word like Iron, its symbol is Fe
for “Ferrum”.
a. Latin is used because it is a “dead” language (will not change over time) and was the original
language of science.
E. The Periodic Table was created based upon increasing Atomic Number.
III. Atomic Mass Units (AMU)
A. Also known as the Mass Number
B. This term refers to the total mass of an atom of that element.
C. It is found by adding the number of protons and neutrons together.
1. Each proton OR each neutron has a mass of 1 AMU or 1 Dalton.
2. The Electrons’ mass is insignificant as they are so small (1/1837th that of protons/neutrons).
D. The Atomic Mass is usually written as a subscript (below) the Element symbol.
E. This was based on Carbon-12 as the standard element of measure. It has 12.0 AMU.
IV. Isotopes
A. This term refers to atoms of an element that have different masses (AMUs) because they have
different numbers of neutrons within the atom; even though it is the same element because they
have the same number of protons. (Remember, protons identify the element.)
1. The isotopes behave relatively the same as the natural atom in terms of chemical properties.
2. Some Isotopes are radioactive (the nucleus is “breaking apart”).
B. To find the number of neutrons:
Start with AMU, subtract the # of protons (atomic number), and that leaves the number neutrons
AMU – # protons = # neutrons
C. How Isotopes are written chemically:
1. Hyphen notation – symbol- number, For example: Carbon-14 OR C-14.
2. Nuclear notation – AMU over Atomic Number symbol, for example 146 C.
D. Nuclide
1. This term is used to refer to the nucleus only (no e- cloud) of an Isotope.
V. Average Atomic Mass
A. As some elements have several isotopes also present in nature, their masses must also be
considered to find the average mass for an element (as seen on the Periodic Table).
B. How to calculate the average Atomic Mass of an element:
Step 1: Multiply the AMU for a single isotope by the % found in nature.
Cu 63 – AMU of 63 = 62.93 AMU; so 62.93 x 69.15% (nature) = 62.93 x .6915 = 43.52 AMU
Cu 65 – AMU of 65 = 64.93AMU; so 64.93 x 30.85% (nature) = 64.93 x .3085 = 20.03 AMU
Step 2: Add the all AMUs together.
43.52 + 20.03 = 63.55 AMU
Step 3: Round to two places after the decimal for each isotope calculation in Step 1.
Content Outline: Nuclear Chemistry Basics (2.3)
I.
Nuclide
A. This term refers to the nucleus of an atom. It is without the electron cloud.
B. Nucleons
1. This term refers to the components of an atom’s nucleus – protons & neutrons.
2. Remember, each proton and neutron has a mass = 1 AMU or Dalton; protons carry a positive
charge and neutrons carry no charge (neutral).
C. Nuclides can be written 2 ways:
1. Hyphen notation – the name followed by the Atomic mass. For example, Radium-228
2. Nuclear notation – Atomic mass over Atomic Number followed by Symbol.
For example, 22888 Ra
D. Nuclear Force
1. Short-range attractive forces that help hold together the nucleus of an atom.
A. Nuclear Binding Energy This term refers to the very small amount of energy that is released when
a nucleus is formed from nucleons.
B. This term can also refer to the energy required to break apart a nucleus.
C. The general trend is: as the nucleus increases in size and number of nucleons, the nuclear binding
energy increases until Iron (Fe). After Iron, the nuclear binding energy decreases due to the
nucleus being so large.
II. Nucleus Stability
A. Nuclear Shell Model
1. Nucleons exist in different energy levels/shells within the nucleus, just like electrons in the cloud.
2. Even numbers tend to be more stable than odd numbers of nucleons.
B. As the number of protons in a nucleus increases, the stability of the nucleus decreases
1. This is because the positive repulsive forces are greater than the Nuclear Force.
2. To reduce this instability, neutrons are needed to increase the distance between protons and
decrease the repulsive forces between protons.
a. This makes the Nuclear Force greater than the repulsive force.
C. Beyond Bismuth (Bi) with Atomic Number of 83, the repulsive forces are so great that there is no
Nuclear force and each atom is unstable and radioactive.
III. Marie and Pierre Curie (1896)
A. The Curie’s were a husband and wife team of French Physicists.
B. Marie Curie was the first woman to win a Nobel in Prize in 1903 for her work on radioactivity. She
later won a second Nobel Prize in 1911. She is the only female to win 2 Nobel Prizes. That is
awesome! Kinda like Alabama winning back-to-back National Championships.
C. Radioactive Decay
1. This is defined as the spontaneous disintegration (falling apart) of a nucleus into a smaller and
lighter (AMUs) by emitting (releasing) particles, electromagnetic radiation, or both.
D. Nuclear Radiation
1. This refers to particles or electromagnetic radiation emitted from the nucleus during radioactive
decay.
E. Radioactive Nuclide
1. Refers to an unstable, radioactive nucleus of an atom.
IV. The types of radioactive decayA. If you change the Atomic Number you change the element. This is known as a transmutation. So you
need a Periodic Table.
B. Alpha Particle Emission (α) or (42He) 1. This is the emission (release) of 2 protons and 2 neutrons as an alpha particle.
2. This causes the Atomic Number to decrease by 2 and the Atomic Mass to decrease by 4 AMUs.
3. This is the weakest penetrating radiation. (A couple sheets of paper can stop it.)
C. Beta Emission (β) or (0-1β)
1. This is when a neutron is converted to a proton and an electron. The electron is emitted from
the nucleus.
2. The Atomic Number increases by 1, but the Atomic Mass remains the same in AMUs.
3. This has greater penetrating strength than Alpha particle emission.
D. Positron Emission (0+1β)
1. This is when a proton is converted to a neutron and a positron. The positron is emitted.
a. Positron is a particle with the same mass as an electron, but possessing a positive charge.
b. The Atomic Number decreases by one, but the atomic mass remains the same in AMUs.
E. Electron Capture
1. An innermost electron is captured by the nucleus and combines with a proton to create a
neutron.
2. The atomic number decreases by 1, but the Atomic Mass remains the same in AMUs.
F. Gamma Emission (γ)
1. This is high energy electromagnetic radiation that is emitted as a nucleus goes from an excited
state to a ground state
a. This is similar to electrons emitting photons as the go from an excited to ground state.
b. The nuclear shell is emitting energy.
2. This is the most penetrating form of radiation. (It can penetrate lead.)
V. Writing and Balancing Nuclear Reactions
Example:
94Be
+ 42 He  126C + 10n Beryllium -9 + Alpha Particle yields Carbon-12 + 1 Neutron
OR
212
84
Po  42 He + 20882 Pb
A.
B.
C.
D.
Step 1: Write in nuclear notation the element that is a reactant.
Step 2: Write in nuclear notation the type of radiation that is a reactant, if adding.
Step 3: Add or subtract the Atomic Masses (top number).
Step 4: Add or subtract the Atomic Number (bottom number) and then use the Periodic Table to find the
new element.
E. Step 5: Write in nuclear notation the new element as a product on the right side of the arrow.
F. Step 6: Write in nuclear notation the other remaining particle on the product side.
Content Outline: Nuclear Energy & Half-Life (2.4)
I.
Decay Series
A. This is a chart showing a series of radioactive nuclides produced by successive radioactive decay
until a lighter, more stable nuclide is reached in the end.
1. The heaviest, most unstable nuclide in a series (at the top) is referred to as the parent nuclide.
2. The nuclides produced (in the downward trail) by the decay of the parent nuclide are referred to
as daughter nuclides.
II. Half-life (t1/2) calculations
A. This is defined as the time required for half of the radioactive atoms in a sample to decay into a
lighter, more stable nuclide.
1. No two radioactive isotopes possess the same rate of decay (half-life).
a. More stable nuclides decay at a slower rate.
b. More unstable nuclides decay at a faster rate.
For example: 3215 P which has a half-life of roughly 14 days.
Starting with 100% 3215 P at TZERO
50% 3215 P remains at 14 days (50% has decayed).
25% 3215 P remains at 28 days (75% has decayed).
12.5% 3215 P remains at 42 days (87.5% has decayed).
6.25% 3215 P remains at 56 days (93.75% has decayed).
3.125% 3215 P remains at 70 days (96.875% has decayed).
B. This calculation can be used in medicine or for telling the age of rocks & fossils.
C. To detect radiation, we use instruments such as Geiger counters, Scintillation counters, and Film
Badges (worn on a person’s uniform).
1. These devices measure Roentgens (R).
2. The Roentgens can be used to calculate REM
a. This is a unit used to measure the dose of any type of radiation that factors into the effect on
human tissue. Radiation can cause cancers by mutating DNA.
III. Nuclear Fusion (“fuse” means “to come together as one”)
A. This is the process where low mass nuclei combine into heavier, more stable nuclei.
B. This releases massive amounts of energy, such as seen in our Sun and stars.
Sun’s example: 4 11 H atoms  42He + 2 01β particles + Energy
IV. Nuclear Fission (“fission” means “to split”)
A. This is the process of a heavy, unstable nuclide, such as U-235, splitting into smaller, more stable
nuclides.
Such as: U-235 splitting into Kr-93 and Ba-140
B. The releases large quantities of energy, such as is carried out in our Nuclear Power Plants.
C. This fission starts a chain reaction (successive chemical reactions) that yields more and more
energy each time a split occurs.
1. The end product is low energy radioactive products called nuclear waste
a. These are hard to dispose of once formed.
b. The largest nuclear waste dump/facility is Yucca Mountain, just outside Las Vegas, Nevada.
D. Our Nuclear Power Plants use Nuclear Reactors to generate electricity.
1. These use the heat energy, from the radioactive chain reaction decay of U-235, to convert liquid
water into steam. The steam is then used to power giant turbines by kinetic movement of the
steam. The turbines turn and generate electricity.
a. The steam is the cooled (condensed) back to liquid water. The water is reused but becomes
radioactive waste over time.
2. Shielding Containment Structures
a. These are usually concrete reactor containment domes.
b. The control the radioactive exposure by absorbing the lose radiation from the reactor.
3. Fuel Rods
a. These are usually long solid rods of highly reactive compact U-235.
b. They eventually become nuclear waste.
4. Control Rods
a. These solid rods are used to control the amount of heat generating radiation by absorbing
free neutrons from the water surrounding the both the control and fuel rods.
i.
If the control rods are lowered down – less energy is produced.
ii.
If the control rods are raised up – more energy is produced.
5. Nuclear Meltdown
A. This is where there is a breach in the containment dome because the fuel rods got to hot
(because the control rods could not be lowered) and melted through the bottom.
B. When this happens massive amounts of radiation are being released into the environment.
1. This recently happened in Fukushima, Japan when a tidal wave hit the plant.
2. It last happened in the U.S. at 3-Mile Island (Pennsylvania) in 1979.
3. A very bad meltdown occurred in Chernobyl, Russia in 1986.
4. All three sites are still highly radioactive locations.
Content Outline: Bohr Model of atoms and Electron Energy (2.5)
I.
Niels Bohr (1913)
A. Danish Physicist
B. Proposed the Bohr Model of Atom structure
1. Electrons travel in set paths around the nucleus called orbits or energy levels.
2. Each orbit corresponds with an energy level.
a. Electrons have a natural tendency to occupy the lowest (most stable) energy level first.
i.
The lowest energy level (ground state) is the closest to the nucleus.
ii.
This is related to Coulomb forces – opposite electrical forces attract.
iii.
The farther away from the nucleus the greater the Potential Energy for that electron.
iv.
The closer to the nucleus the less Potential Energy for that electron.
b. Electrons can absorb energy (absorption) from their surroundings from another energy
source, such as sunlight energy (A.K.A electromagnetic energy).
i.
Electrons that gain energy (absorption) are said to be “excited”.
ii.
Electrons that lose energy (emission) emit light as they return to a more stable (less
energy) grounded state.
iii.
The unit of light energy is referred to as a photon.
iv.
The unit of measurement for the energy lost OR gained by an atom is a quantum.
II. Electrons as Particles & Waves
A. Electrons can move as particles around the nucleus because they have mass (if ever so small).
B. Electrons, as they are moving, move in wave-like fashion (like waves in the Gulf of Mexico… up,
down, up, down)
C. Properties of waves
1. Wavelength (λ)
a. This is defined as the distance between two identical points (such as crest –top or ebbbottom) on adjacent waves.
i.
As it is distance, some unit of measurement of distance (meter) is used for
wavelength, usually nanometer (nm OR 10-9).
2. Frequency (v)
a. This is defined as the number of waves that pass a given point in a specified time, usually
seconds.
b. Frequency is expressed in Hertz (Hz) or waves/sec.
i.
Heinrich Hertz defined 1 wave/sec = 1 Hertz.
3. Speed of light (c)
a. Electrons travel at the speed of light.
b. Waves are measured against the speed of light (electromagnetic radiation).
c. C = λv is the equation for the speed of light.
i.
As light speed never changes, it is considered to be a constant at 3.00 x 108 m/sec.
ii.
The properties of light are inversely proportional.
α. As wavelength decreases, frequency increases.
b. As wavelength increases, frequency decreases.
d. As electrons gain more energy, they travel faster and get farther from the nucleus.
III. Electromagnetic Spectrum (Light Energy)
A. This term refers to the whole spectrum (variations) of electromagnetic radiation.
1. Radiation is used to define the wave-like movement of light particles.
2. Light moves at 3.00 x 108 m/sec.
3. The electromagnetic spectrum includes: x-rays, microwaves, visible (white) light, ultra-violet
light, infrared light, and radio waves.
Content Outline: Photoelectric Effect and Emission Spectrum (2.6)
I.
Photon
A. Albert Einstein proposed this concept in 1905.
B. A Photon is a unit of light energy having no mass and possessing a single quantum of energy.
1. Quantum
a. Minimum quantity of energy that can be lost or gained by an atom.
b. Proposed by German physicist Max Planck in 1900.
c. This sets the field of Quantum Physics (nanoscale physics) in motion.
d. Planck wins the Nobel Prize in 1918 for this work.
II. Photoelectric Effect
A. The emission (ejection) of an electron from a metal surface when light shines on the surface.
1. This shows a direct connection between light and light possessing energy.
2. This light energy is Einstein’s photon.
B. The light has to be of a minimum frequency in order for the effect to take place.
C. Each metal requires a different frequency of light.
III. Planck’s Constant Theory
1. This tries to explain the Photoelectric Effect by proposing a relationship between a quantum of
energy and the frequency of radiation. Remember, light is considered electromagnetic radiation,
so the frequency changes with the various forms of light/radiation. This is called the
Electromagnetic Spectrum.
A. E = hv
1. E = energy for a quantum of radiation.
a. It is measured in Joules (J).
2. v = frequency of the radiation.
a. measured in waves/sec. (s) -1 or Hertz
3. h = Planck’s Constant
a. Defined as 6.626 x 10-34 J * s (* = times)
B. Planck-Einstein Relation
1. Albert Einstein expanded on Planck’s work in 1905.
2. He proposed that light has a combination of wave properties and particle properties.
a. Each particle of light carries 1 quantum of energy
3. E = hv (Planck’s version) then becomes Ephoton = hv (Einstein’s version).
4. Matter can only absorb Electromagnetic radiation (light) in whole number (1, 2, and so forth)
quantities of photons.
5. In order for a single electron to be emitted from the metal surface, the electron must be struck
by a single photon possessing at least the minimum amount of energy to eject the electron.
a. This minimum amount of energy is directly related to the frequency.
i. The greater the frequency – the more possible to emit an electron from the surface.
ii. The smaller the frequency – the less likely to emit an electron from the surface.
b. Different metals require different frequencies for the Photoelectric effect to take place.
IV. Emission Spectrum of Hydrogen
A. This is an expansion of the Cathode-Ray Tube experiment.
1. It uses Hydrogen gas (which glows pink) and a glass prism (triangular shaped piece) placed in
the path of the light ray.
a. The light ray split into 4 different colors (red, green, blue, and purple).
i.
These become known as the light emission spectrum.
ii.
Each color represents a fixed quantity of energy for an excited electron.
iii.
It is later added for other frequencies of light, such Infrared and Ultra-violet.
V. These experiment set the groundwork for the Modern Quantum Atomic Model of atoms.
Content Outline: Modern Quantum Model of Atoms (2.7)
I.
Louis de Broglie (1924)
A. He was a French Scientist.
B. He proposed that electrons (e-) had wave-like properties using Bohr’s Atomic model.
1. He stated that electrons are confined to areas around the nucleus, known as orbits.
C. Using the Planck-Einstein Relation, he reasoned that electrons have specified energies/frequencies.
Diffraction (bending of waves) experiments with electron beams and light beam showed similar
results, proving electrons travel like light… in waves.
D. Interference (waves overlapping) experiments with waves and electron beams showed similar
results, proving electrons have particle like properties associated with energy.
1. Some areas increased in energy as a result of overlap; other areas decreased in energy.
II. Werner Heisenberg (1927)
A. He was a German Physicist.
B. He calculated that electrons and photons have about the same amounts of energy.
1. Photons are used to help detect the presence of electrons.
a. When they collide, the electron is deflected in a random direction.
C. Heisenberg’s Uncertainty Principle
1. This states that it is impossible to determine exactly both position and velocity simultaneously of
an electron in an orbit.
III. Erwin Schrödinger (1926)
A. He was an Austrian Physicist.
B. He helped develop the Quantum Theory of Atoms
1. This Theory tries to describe, by mathematics, the wave-like properties of electrons and other
very small particles.
2. This re-enforces that electrons travel in orbitals.
a. Orbital is defined as a 3- Dimensional region around a nucleus that indicates the probable
location of a single electron within an orbital.
C. Quantum Numbers
1. These are four numbers, written consecutively, that specify the properties of atomic orbitals and
properties of the electrons within an orbital.
For example: n,l,m,s 1,0,0,1/2 or 2,1,1,-1/2
2. The first number: the Principle Quantum number
a. Symbolized as “n”.
b. This states the electron energy level (or shell).
c. Uses only whole numbers 1-7.
d. As “n” increases, the distance from the nucleus also increases; therefore the Potential Energy
of that electron also increases.
3. The second number: the Angular Momentum Quantum number
a. Symbolized as “l”.
b. This states the shape of orbitals or sub-orbitals.
i.
0 = s (sphere shaped); there is only 1 orbital/energy level.
ii.
1 = p (dumbbell shaped); there can be up to 3/energy level.
iii.
2 = d (4 leaf clover shaped with a possible ring); there can be up to 5/energy level.
iv.
3 = f (too complex to describe); there can be up to 7/energy level.
4. The third number: the Magnetic Quantum number
a. Symbolized by “m”.
b. This states the orientation (up/down, left/right, forward/backward) of an orbital around
the nucleus.
i.
s=0
ii.
p = 1, 0, -1 (are possible)
iii.
d = 2, 1, 0, -1, -2 (are possible)
iv.
f = 3, 2, 1, 0, -1, -2, -3 (are possible)
5. The fourth number: the Spin Quantum number
a. Symbolized as “s”
b. As each orbital can only hold two negatively charged electrons, they must be opposites in
spin (motion).
c. We use either +1/2 (up) or -1/2 (down) for this.
IV. Orbital Electron Configuration Models
A. These are used to help represent the electron arrangement of an atom; electrons are represented by
numbers such as 1s1 for Hydrogen or 1s2 2s2 2p5 for Fluorine or as arrows, either 
B. It is important to remember that each orbital can only hold 2 electrons maximum.
C. You “build the model” starting from the ground state (lowest energy level) and work upward,
starting at 1s1 ; the exponent states the number of electrons present at that level… hint look at the
periodic table ROWS. Highest is row 7.
D. This uses 3 rules:
1. Aufbau Principle (“aufbau” is German for “to build”)
a. Electrons occupy the lowest energy level first that can receive the electron. Starting at 1s
and work toward row 7.
2. Pauli Exclusion Principle
a. Contributed by Wolfgang Pauli, an Austrian physicist.
b. No 2 electrons, in the same atom, can have the same 4 Quantum numbers.
3. Hund’s Rule
a. Contributed by Friedrich Hund, a German Physicist.
b. States that orbitals and sub-orbitals of equal energy are each occupied by a solo electron
before a second electron is entered into the orbital or sub-orbital and all electrons in singly
occupied orbitals/sub-orbitals must have the same spin number. (Start with =1/2 ; then go
back and use -1/2 )
V. Noble Gas Electron Configuration
A. This method is used to shorten the traditional model building.
B. Starting at Neon (Ne), you can begin the notation using a Noble gas first.
1. The Noble Gas must be in [ ]. Then start with the next energy row.
Content Outline: Periodic Trends associated with Electrons and Coulombs Forces (2.8)
I.
Coulomb’s Forces
A. These are the attractive forces that exist between two oppositely electrically charged particles, such
as positive protons and negatively charged electrons.
1. Remember, in the natural state, an atom has equal numbers of protons and electrons; therefore,
they are electrically neutral.
2. As the protons are larger in mass, if an atom has more protons than electrons, the electrons are
pulled inward and thereby possess less Potential Energy. See below 1 & 2
3. If the electron has more electrons than protons, the electrons are able to be farther from the
nucleus and thereby possess greater Potential Energy. See below 1 & 2
B. The formula for Coulomb’s Force is:
l q1q2 l
This equation is affected by 1) magnitude (number) of charges (q)
F = ke x r2
q1 would be the number of Protons; q2 number of electrons
2) Distance – r2
9
ke – is Coulomb’s constant = 8.99 x 10
II. Effective Nuclear Charge
A. This is symbolized as ZEFF
B. This is the net positive charge experienced by a single outer shell electron of an atom.
1. The Pull the electron “feels” by the nucleus.
2. Electron shielding
a. This term refers to the inner energy level of electrons helping to reduce (shield) the
outermost electrons from the full pull of the nucleus. Hence the term electron shielding.
i.
This allows the outermost electrons to remain farther away from the nucleus.
ii.
The more electrons an atom has, the greater the distance away from the nucleus
because of greater shielding.
III. The Periodic Trends in the Periodic Table
A. There are trends across a Period (row), referred to as Period Trends. There are trends within a
Family (column), referred to as Group Trends or Family Trends.
B. Atomic Radii Trends
1. Period Trends
a. Atomic Radii decreases as you go across (left to right) a period. This is due to greater
Coulombs forces or greater ZEFF. (Nucleus has greater pull on the electrons.)
2. Group Trends
a. Atomic Radii increases as you gown down a group. This is due to greater numbers of
electrons in higher energy levels being present and greater electron shielding. (Nucleus has
less pull.)
b. Exceptions 1) – Gallium. It is because Gallium has all the Group D protons, whereas Aluminum
does not. Therefore Gallium is small due to greater Coulombs forces or ZEFF. 2) At the hallway
point through the “d” and “f” blocks as each sublevel is half full. It resumes when the next
electron is added into a half filled orbital.
C. Ionization Energy Trends
1. Ion
a. This is a charged particle that is the result of gaining or losing an electron.
b. Cation
i.
These are positively charged particles/atoms that have lost electrons.
ii.
Because of the Law of Conservation of Matter, the electrons must have gone to
another location.
α. The Law states: Matter is neither created nor destroyed; just transferred and
transformed.
c. Anion
i.
These are negatively charged particles/atoms that have gained electrons.
d. Ionization
i.
This is the process of creating ions by gaining or losing electrons. (“tion” means
“process of”).
e. Ionization Energy (IE)
i.
This is the energy required to remove one electron from a neutral atom.
ii.
It is measured in kilojoules/mole (kJ/mol).
2. Period Trends
a. Group 1 Metals have the lowest Ionization Energies. (They lose electrons easily, hence their
being very reactive.)
b. Group 18 (Noble Gases) have the highest Ionization Energies on the whole Periodic Table.
(They do not lose electrons easily, hence their being unreactive.)
c. The Ionization Energies increase as you go across the period due to greater Coulomb Forces
or increased ZEFF. (The nucleus has a greater pull…requires more energy to remove.)
3. Group Trends
a. Ionization Energies generally decrease as you down a group due to lesser Coulomb Forces or
weaker ZEFF. (Electrons are farther away (more shielded) from the nucleus, so less energy is
needed to remove electrons.)
4. To remove more than one electron, even greater amounts of energy are needed than for the first.
This information can be found on an ionization table. With each electron removed, the greater
the Coulombs forces or increased ZEFF become within the atom.
D. Electron Affinity
1. This term refers to the energy change (positive or negative) to add an electron to an atom.
a. If the change is negative, energy has been released. (This is easy to do.)
b. If the change is positive, energy has been invested into the atom. (This is hard to do.)
i.
It requires a greater pushing force be put on the atom to overcome the negative
repulsive forces, thus making the atom very unstable. (This is what happens in your
cell phone… we use the greater electrical force to reverse the electron flow back to a
“charged state”.)
2. It is measured in kilojoules/mole (kJ/mol) too.
3. Period Trends
a. Group 17 (Halogens) has the greatest affinity as shown by large negative numbers.
b. In general, Electron Affinity increases as you go across a period.
c. Exception to the rule: Carbon – adding an electron allows the atom to reach a half filled p sublevel. That makes adding the fourth into a half filled sub-orbital more difficult as seen in N.
4. Group Trends
a. In general, Electron Affinity decreases (become more difficult) as you go down a group
because of greater Atomic radii and increased electron repulsive forces.
5. Adding a second electron is generally more difficult, unless it is a Group 16 element, such as
Oxygen. It needs 2 electrons to fill the valence shell.
E. Ionic Radii
1. Cations have decreased atomic radii due to greater Coulombs Forces or greater ZEFF. (The
electron cloud gets smaller.)
2. Anions have increased atomic radii due to greater electron shielding and no increase in ZEFF.
(The electron cloud gets larger.)
3. Period trends
a. Ionic Radii decreases until Group 15 and then begins to increase.
b. This is because Groups 1 – 14 are more Cationic. See above in F.1
c. Groups 15 – 17 are more Anionic. See above in F.2
4. Group Trends
a. Ionic Radii increases as you go down a group due to greater energy levels of electrons.
F. Electronegativity
1. This can be thought of as the desire to acquire negative electrons from another atom.
2. Fluorine (F) is the most electronegative element. (It has the strongest desire to be “like” a noble
gas…if it can just get that last electron in the valence shell.
3. Francium (Fr) is the least electronegative element. (It has very little desire to acquire an
electron. It would rather get rid of the solo electron in its valence shell and then be “like” the
Noble Gas Radon (Rn).)
4. Period Trend
a. Electronegativity generally increases going across a period.
5. Group Trend
a. Electronegativity decreases as you go down a group because the valence electrons are
farther from the nucleus. (There is more electron shielding, so it is less desirous to acquire
electrons.)