Chapter 3 Atomic Structure

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Chapter 3Atomic Structure

• Chemistry
• Mr. Bass

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Atomic Structure

• 3-1 How are elements organized?
• 3-2 What is the basic structure of an atom?
• 3-3 How do the structures of atoms differ?

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3.1 How are elements organized?

• Objectives
• Periodic Table
• Basic Components of an Atom
• Basic Definitions

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3.2 What is the basic structure of an atom?

• Creating Atomic Models by Inference
• 3 Laws of Nature
• John Daltons Atomic Theory
• J.J. Thompsons Atomic Theory
• Ernest Rutherfords Atomic Theory
• The Physics of Energy
• Niels Bohrs Atomic Theory.

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3.2 What is the basic structure of an atom?

• Objectives (SWBAT)
• infer the existence of atoms from the laws of
  definite composition, conservation of mass, and
  multiple proportions.
• list the five basic principles of Daltons atomic
  theory.
• describe Daltons, Rutherfords, and Bohrs
  atomic models.

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3.2 What is the basic structure of an atom?

• Objectives (continued) SWBAT
• compare and contrast the properties of electrons,
  protons, and neutrons.
• explain the particle-wave nature of electrons.
• describe the quantum model of the atom.

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Creating Atomic Models by Inference
Indicator 1.4.8

• Study the patterns of nature.
• Develop models that fit the information.
• Test/Refine the models.
• Wind It cant be seen, but its force can be
  felt.
• The evidence of the wind is indisputable. This
  is inference.
• Straw Men Develop a model and then try to
  destroy it.
• Models should reflect the properties of nature.

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3 Known Laws of Nature

• Law of Definite Composition
• Law of Conservation of Mass
• Law of Multiple Proportions

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Known Laws of Nature

1. Law of definite composition a compound contains
   the same elements in exactly the same proportions
   by mass regardless of the size of the sample or
   source of the compound.

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Law of Definite Composition
Sugar (Sucrose) has the same composition
regardless of the size of the sample or source of
the sample. 100 g of sugar 42.1 carbon, 51.4
oxygen, 6.5 hydrogen. 100 Mg of sugar 42.1
carbon, 51.4 oxygen, 6.5 hydrogen. Sugar
from Sugar Beats 42.1 carbon, 51.4 oxygen,
6.5 hydrogen. Sugar from Sugar Cane 42.1
carbon, 51.4 oxygen, 6.5 hydrogen.
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Law of conservation of mass

• In a chemical reaction, the mass of the reactants
  is equal to the mass of the products.
• Restated In a chemical reaction, mass is
  neither created nor destroyed.

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Law of Conservation of Mass

• Combination of Atoms
• 32 g of S and 32 g of O2 ? 64 g of SO2
• Separation of Atoms
• 434 g of HgO ? 402 g of Hg and 32 g of O2
• Rearrangement of Atoms
• 62 g of H2CO3 ? 18 g H2O and 44 g CO2
• In every case, the mass of the products is equal
  to the mass of the reactants!

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Law of Multiple Proportions

• The mass ratio for one of the elements that
  combines with a fixed mass of the other element
  can be expressed in small whole numbers.
• This compares two substances made of the same
  elements. For example, water and hydrogen
  peroxide (both composed of hydrogen and oxygen).

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Law of Multiple Proportions

• 16g O2 2g H2 ? 18g Water
• 32g O2 2g H2 ? 34g Hydrogen Peroxide
• Ratio of Oxygen in two compounds is
• 32 g O2 21 ratio of oxygen
• 16 g O2

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How do these 3 laws infer the existence of atoms?

• Law of definite composition infers that there
  must be small units (atoms) because compounds
  always have the same percent composition.
• In order for compounds to always have the same
  percent composition, there must be a smallest
  unit or particle of an element that makes up the
  compound.

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Atomic Inference

• Law of conservation of mass infers that there
  must be atoms because the sum of the reactants
  mass is equal to the sum of the mass of the
  products.
• This law is always true no matter how many (or
  few) units of mass are used.
• This infers that there must be a basic small unit
  of nature (atom) that is being swapped around
  during a chemical reaction.

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Atomic Inference

• Law of Multiple Proportions infers that there
  must be a small unit of nature (atoms) because
  every compound composed of the same elements can
  be reduced to a simple mass ratio.
• That simple ratio is caused by the fact that
  compounds are made of atoms, and when atoms
  combine in small whole number ratios it is
  reflected in the mass ratio.

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Atomic Inference

• Law of Multiple Proportions (Continued)
• This also infers that atoms of different elements
  have different masses.
• This also infers that different atoms of the same
  element have the same mass.

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John Dalton 1776 - 1844
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John Dalton

• Together the 3 experimental laws represented much
  of the quantitative data that chemists had in the
  1700s.
• They implied the existence of what was to become
  known as the atom.
• Atom is from greek atomos meaning indivisible.
• John Dalton was the first one to put all the
  pieces together in 1805.

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John Daltons Atomic Theory

1. All matter is made of indivisible and
   indestructible atoms.
2. Atoms of a given element are identical in their
   physical and chemical properties.
3. Atoms of different elements have different
   physical and chemical properties.
4. Atoms of different elements combine in simple,
   whole-number ratios to form chemical compounds.

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John Daltons Atomic Theory

• Atoms cannot be subdivided, created, or destroyed
  when they are combined, separated, or rearranged
  in chemical reactions.
• Daltons theory brought much attention from other
  scientists who tested it.
• While some exceptions were found, Daltons Atomic
  Theory has not been discarded, just modified and
  expanded.

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Definition of Element
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Testing the atomic theory

• During the next 200 years the atomic theory was
  tested over and over again.
• One of the main ways of testing the atomic theory
  was a device used during the 1800s called a
  Cathode Ray tube.

• The Cathode ray tube was essentially a low
  pressure tube that had electricity put through
  it. It produced a light called a cathode ray.

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Cathode Ray Tube
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Cathode Ray Tube

• Cathode a negative electrode through which
  current flows.
• Anode a positive electrode through which
  current flows.
• The reason for calling the light a cathode ray
  was because it appeared to start at the cathode
  and go to the anode.
• The cathode ray is the basis for todays TVs and
  monitors.

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Cathode Ray Tube Experiments
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The search for smaller particles

• Basic Discoveries made with the cathode ray tube.
• The ray originated at the cathode and goes to the
  anode.
• This infers that the ray is negative in charge.

• Principles of Nature
• The electric charge of matter is normally
  neutral.
• Opposite Charges Attract
• Like Charges Repel

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The search for smaller particles

• When a paddle wheel was placed in front of the
  cathode ray it moved toward the anode.
• This infers that the cathode ray must be composed
  of small, individual particles that could push
  the paddle wheel down the cathode-ray tube.
• Late in the 19th century G. Johnstone Stoney
  named the small, negatively charged particles
  electrons.

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J.J. Thomson

• English physicist.
• Discovered the electron in 1897.
• Nobel Prize Winner 1906.
• Negative in charge.
• Almost no mass.

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J. J. ThomsonIn His Own Words
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J. J. Thomson

• When a magnet or charged plates were placed above
  the cathode ray, the ray was deflected.
• This implied that the mass of an electron was
  small.
• Also that an electron had a negative charge.

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J. J. Thomson
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J. J. Thomsons Atomic Theory

• Based upon his experiments and observations, J.
  J. believed that the atom was a solid ball with
  electrons located on the outer skin of the ball.
• This was called the plum pudding model because of
  the appearance of raisins in pudding are like the
  electrons in the theory.

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Earnest Rutherford

• A student of J. J. Thomson.
• From Australia originally.
• Determined that atoms are composed mainly of
  space, with a small dense center.
• Discovered the proton.
• Won the Nobel Prize in 1908.

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Ernest Rutherford

• Coined the names of many atomic particles.
• alpha, beta, and gamma rays
• proton, neutron
• half life, daughter atoms
• Many influential scientists studied under him.
• Neils Bohr
• James Chadwick
• Robert Oppenheimer

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Rutherfords Gold Foil Apparatus
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Gold Foil Apparatus

• The atom is composed primarily of space.
• All of the positive charges are in a dense
  center.
• The center of the atom contains the vast majority
  of the mass of an atom.
• Discovered the nucleus.
• Nucleus from Latin word meaning little nut

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Ernest Rutherfords Atomic Theory

• Electrons travel in the space surrounding the
  nucleus in a way similar to the motion of the
  planets around the sun.
• Called the planetary model.
• There is a total of 7 Primary Shells.
• The difference between these shells is the
  distance from the nucleus.

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Ernest Rutherfords Atomic Theory
Electrons orbit the nucleus like the planets
orbit the sun. 7 primary orbits (or shells) More
than one electron can be in an orbit.
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James Chadwick

• In 1932 a British scientist discovered the
  neutron.
• He recognized that these particles had the same
  properties as those proposed by Ernest
  Rutherford.
• Neutral particles that have a mass equal to that
  of protons.

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The Physics of Energy

• As scientists came to have a more complete
  understanding of the atom, they began to study
  because they thought it might hold the answer to
  the structure of an electron.
• Waves a characteristic pattern of energy.
• Dual Nature of Light Light behaves both as a
  mass particle (called a photon) and as energy
  (called a wave).

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Electromagnetic Spectrum
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Electromagnetic Spectrum

• Electromagnetic Spectrum all wavelengths of
  light.
• Visible Spectrum only visible wavelengths of
  light.
• Prism separates light into the different
  wavelengths.
• Diffraction Grating separates light into the
  different wavelengths.

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Electromagnetic Spectrum

• Continuous Spectrum All wavelengths of light
  are seen.

ROY G BIV Order of colors. Red, Orange, Yellow,
Green, Blue, Indigo, Violet.
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Electromagnetic Spectrum
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Electromagnetic Spectrum

• Absorption Spectrum Wavelengths of light are
  absorbed by a substance.
• Produced by passing light through cool gases.
• White Absorption No light absorbed, all light
  reflected.
• Black Absorption All light absorbed. No light
  reflected.
• Absorption Movie

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Electromagnetic Spectrum

• Emission Spectrum Only certain wavelengths of
  light are seen.
• Produced by passing electricity through
  hot/excited gases.
• Color seen is characteristic for element.
• Color seen is a blend of the specific wavelengths
  of light.

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Electricity and Magnetism

• If the blue portion is the wave created by
  electricity, the green portion is the magnetic
  field generated at a 90o angle.

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Standing Wave

• Standing Wave A wave with nodes that do not
  move.

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Wave Terms

• Wavelength (?) The distance between two
  identical portions of a wave.
• Frequency (f) The number of times a wave passes
  a particular spot in a set period of time. The
  unit for this is hertz (Hz 1/s)
• Axis The midline of a wave.
•  

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Wave Terms

• Node Where the wave crosses the axis.
• Peak The top of the wave.
• Trough The bottom of the wave
• Amplitude The measurement from axis of a
  
  wave to either the peak or the trough.

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Relationship between Frequency and Wavelength

• Wavelength is inversely proportional to
  frequency. f ? 1/?

Math Principle Whenever two factors are
proportional, they can be made equal by
multiplying by a constant.

• The constant for this relationship is velocity
  (speed) so f v/?
• For light f c/?, where c speed of light
• c 3.00E8 m/s

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Max Planck 1858-1947

• Nobel Prize Award 1918
• Discovered the value of the constant that relates
  energy to wavelength.
• Discovered the fundamental concept of quanta,
  which lead to the quantum theory.

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Planks Constant

• Energy was discovered to be proportional to
  frequency. E ? f
• Plank discovered the constant that made it equal,
  and the constant was named after him. E hf
  where h Planks Constant
• Since E hf, and f c/?, then Ehc/?
• This is the main math we will use in the bright
  line experiment (Exp 3C)

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The deBroglie Hypothesis

• From Einstein we know Emc2
• From Planck we know Ehc/?
• Therefore mc2hc/?
• Cancel c and mch/?
• Now solve for ?
• ? h/mc
• This implies that moving particles have
  wavelength. Interesting, eh???

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Heisenberg Uncertainty Principle

• The problem with the wave-particle duality of
  nature both natures cannot be tested at the
  same time with the same experiment.
• Heisenberg stated it this way It is impossible
  to know the exact position and the exact momentum
  at the same time.

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Niels Bohr 1885 - 1962

• 1913 proposed that electrons could only reside in
  certain energy levels (quanta).
• Danish physicist.
• Proposed the step ladder analogy.
• Took all of the available information and
  synthesized it into the proposed quantum theory.
• Nobel Prize 1922

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Niels Bohr

• What he saw

• 6 bright lines of light for the atomic emission
  of hydrogen.
• Using Planks equation he related the lines of
  light to specific quantities of energy.

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Hydrogen Emission Spectrum

• Bohr stated that the electron usually stays at
  the lowest energy level possible.
• He called this the ground state.

• Principles of Nature
• The sum of electric charges are usually equal in
  nature.
• Opposites Attract Likes Repel
• Entropy Things tend to go to the lowest energy
  level possible.

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Hydrogen Emission Spectrum

• Bohr stated that when energy is passed through a
  sample of gas most of the energy is not absorbed.
• Electrons only absorb the energy when it is
  exactly the same amount as the next higher energy
  level for the electron.
• Excited State When an electron has absorbed
  energy to go to a higher energy level.
• Electrons stay in the excited state for only
  moments before falling back to ground state.

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Niels Bohr
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Hydrogen Emission Spectrum
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Neils Bohr Hydrogen Emission Spectrum
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Hydrogen Emission Spectrum

• Quantum A specific amount of energy.
• Bohr stated that electrons could only absorb
  specific quanta of energy. When this energy is
  absorbed the electron is in an excited state for
  a moment. It quickly returns to the ground state
  because of its instability.
• Entropy Things tend to go to the lowest energy
  level possible.

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Hydrogen Emission Spectrum

• Bohr studied a series of 6 lines called the
  Balmer series. Because there was 6 lines he
  hypothesized that there must be 7 principle
  shells.
• Later other series of lights confirmed what Bohr
  had hypothesized.
• Still later it was hypothesized that there must
  be another energy level called sub-shells.

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Confirmation
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Bohrs Atomic Theory

• Electrons can only exist in certain energy
  levels.
• Primary Shell Distinguished by the distance
  from the nucleus (7 primary shells) same as
  Rutherfords model.
• Sub-shell Distinguished by the shape of the
  electron cloud.
• Four sub-shells, s, p, d and f

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Electron Cloud

• Area of high probability of finding an electron.
• Each of the sub-shells are different types of
  electron clouds

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Electron Orbitals

• Each orbital can contain up to two different
  electrons (in a normal situation).

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s Sub-Shell
s Sub-Shell round ball shape region.
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p Sub-Shell
p Sub-Shell Pear shaped pairs of lobes on all
three axis.
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d Sub-Shell
d Sub-Shell Each sub-shell has four pairs of
pear-shaped lobes.
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f Sub-Shell
f Sub-Shell Seven sets of four pear-shaped
lobes around the different planes of the atom.
This is extremely difficult to draw.
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Sub-Shells Chart

s
p
d
f
of e- Possible (in Sub-Shell) 2 6 10 14
Principle Shell 1 2 3 4
of Orbitals 1 3 5 7
Multiply by 2 X 2 X 2 X 2 X 2
Sub-Shells Some People Do Fine
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3.3 How do the structures of atoms differ?

• Objectives
• Shorthand Notation
• Valence Notation
• Electron Configuration (Orbital Diagram)
• Quantum Numbers

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Objectives (SWBAT)

• write the shorthand notation for any element.
• write the valence notation for any element.
• write the electron configuration for any element.
• write the quantum numbers for any electron.
• discuss the total possible number of electrons
  for any principle or sub-shell.

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Shorthand Notation

• Shorthand Notation describes the sub-shell
  location of every electron of an element.

Sub-shell electrons are found in
1S2
Total number of electrons in sub-shell
Primary Shell that sub-shell is located in
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Locating Sub-Shells on Periodic Table
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Primary/Sub-Shells
1 2 3 4 5 6 7
2 3 4 5 6 7
Primary Shell 1st Seen S 1 P 2
Consecutive s D 3 F 4
3 4 5 6
4 5
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Shorthand Notation Guidelines

1. Treat the atomic number as the electron .
2. Always count electrons in numerical order.
3. The superscript is always the number of electrons
   in that sub-shell and that principle shell.
4. The sum of the superscripts should add up to the
   atomic number of the element.
5. Electrons are listed in order of increasing
   energy. This is called the Aufbau Principle
   (means building up in German).

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Shorthand

• H 1s1
• He 1s2 Noble Gas
• Li 1s2 2s1
• Be 1s2 2s2
• B 1s2 2s2 2p1
• C 1s2 2s2 2p2
• N 1s2 2s2 2p3

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Shorthand

• O 1s2 2s2 2p4
• F 1s2 2s2 2p5
• Ne 1s2 2s2 2p6 -Noble Gas
• Na 1s2 2s2 2p6 3s1
• Mg 1s2 2s2 2p6 3s2
• Al 1s2 2s2 2p6 3s2 3p1
• Si 1s2 2s2 2p6 3s2 3p2

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Shorthand

• P 1s2 2s2 2p6 3s2 3p3
• S 1s2 2s2 2p6 3s2 3p4
• Cl 1s2 2s2 2p6 3s2 3p5
• Ar 1s2 2s2 2p6 3s2 3p6 Noble Gas
• K 1s2 2s2 2p6 3s2 3p6 4s1
• Ca 1s2 2s2 2p6 3s2 3p6 4s2
• Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1

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Shorthand Alternative Method
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Valence Notation

• Valence Electrons Electrons in the outermost
  shell of the atom.
• Noble Gases The outermost shell of the atom is
  always full for a Noble Gas.
• Valence Notation Start with the last Noble Gas
  before the element and do the shorthand notation
  from there.

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Valence Notation

• H 1s1
• He He
• Li He 2s1
• Be He 2s2
• B He 2s2 2p1
• C He 2s2 2p2
• N He 2s2 2p3

• H 1s1
• He 1s2 Noble Gas
• Li 1s2 2s1
• Be 1s2 2s2
• B 1s2 2s2 2p1
• C 1s2 2s2 2p2
• N 1s2 2s2 2p3

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Valence Notation

• O 1s2 2s2 2p4
• F 1s2 2s2 2p5
• Ne 1s2 2s2 2p6 -Noble Gas
• Na 1s2 2s2 2p6 3s1
• Mg 1s2 2s2 2p6 3s2
• Al 1s2 2s2 2p6 3s2 3p1
• Si 1s2 2s2 2p6 3s2 3p2

• O He 2s2 2p4
• F He 2s2 2p5
• Ne Ne
• Na Ne 3s1
• Mg Ne 3s2
• Al Ne 3s2 3p1
• Si Ne 3s2 3p2

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Valence Notation

• P Ne 3s2 3p3
• S Ne 3s2 3p4
• Cl Ne 3s2 3p5
• Ar Ar
• K Ar 4s1
• Ca Ar 4s2
• Sc Ar 4s2 3d1

• P 1s2 2s2 2p6 3s2 3p3
• S 1s2 2s2 2p6 3s2 3p4
• Cl 1s2 2s2 2p6 3s2 3p5
• Ar 1s2 2s2 2p6 3s2 3p6-Noble Gas
• K 1s2 2s2 2p6 3s2 3p6 4s1
• Ca 1s2 2s2 2p6 3s2 3p6 4s2
• Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1

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Quantum Mechanics

• As the Quantum Theory of Atoms advanced it soon
  became apparent that locating an electron was
  going to be impossible (see Heisenbergs
  Uncertainty Principle and Electron Cloud).
• Soon a branch of math called Quantum Mechanics
  was developed to help identify areas of high
  probability of finding an electron.

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Quantum Mechanics

• As the quantum mechanics developed it was soon
  apparent that there were four energy levels
  instead of two
• Principle Shell (1-7)
• Sub-Shell (s, p, d, f)
• Orbital different lobes of the sub-shells
• Spin clockwise and counter-clockwise

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Electron Configuration

• A new way developed to describe all the energy
  levels of electrons in an atom.
• Electron Configuration Location of all the
  different energy levels for all the atoms of an
  element. (Also called Orbital Diagram)
• Orbital Different lobes of the sub-shells.
• Spin either clockwise ( ) or counter-clockwise
  ( )
• Clockwise is lower energy, always goes 1st.

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Electron Configuration Conventions

• There are two basic ways that electron
  configurations are drawn.
• Each student may choose which to use.
• There is one box (line) per orbital of the
  sub-shell
• Using boxes as orbitals
• Using lines as orbitals
• Arrows are electrons

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Electron Configuration Rules

1. Treat the periodic table as a table of electrons
   for this exercise. Follow the same order as the
   electron configuration.
2. Each orbital can only hold two electrons.
3. Electrons are added in order of increasing
   energy.
4. Pauli Exclusion Principle no two electrons can
   have the same energy levels.

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Electron Configuration Rules

• Hunds Rule the most stable arrangement of
  electrons is that with the maximum number of
  unpaired electrons, all with the same spin
  direction.
• Within a sub-shell each orbital must have one
  electron before any can have a second electron.
• Exception Group 6 and Group 11 elements. These
  steal an electron from the s sub-shell to have 5
  electrons in the d sub-shell.
• Apparently the d half-filled and filled
  sub-shells are a lot lower in energy level.

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Electron Configuration

• H
• He
• Li

• Be
• B
• C

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Electron Configuration

• N
• O
• F

• Ne
• Na
• Mg

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Electron Configuration

• Ar
• Kr

Exception!!!!
Exception!!!!
99
Orbital Notation
100
Quantum Numbers

• It had been known for some time that it was only
  the outermost (valence) electrons that affected
  the behavior of the atom.
• A shortcut called quantum numbers was developed
  to give the energy level of individual electrons.
• Primary Shell (n) There are 7 primary shells
  and they are numbered 1-7.
• Sub-Shell (l) 4 sub-shells, 0-3
• s 0
• p 1
• d 2
• f 3

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Quantum Theory

• Orbital (ml) This is like a number-line, the
  center orbital is always numbered 0.
• s 0
• p -1 0 1
• d -2 1 0 1 2
• f -3 -2 -1 0 1 2 3
• This relates to the electron configuration
  diagrams

-3 -2 -1 0 1 2 3
-2 -1 0 1 2
102
Quantum Numbers

• Spin (ms)
• Use 1/2 for clockwise ( ) lowest energy
• Use 1/2 for counter-clockwise ( ) higher energy

103
(n) Principle Quantum Number and Periodic Table
Some 1 People 2 Do 3 Fine 4







1 2 3 4 5 6 7






2 3 4 5 6 7




3 4 5 6


4 5
104
Sub-Shell Quantum Number
Some s 0 People p 1 Do d 2 Fine f 3

















0
1
2
3


105
Orbital ml Quantum Number
Some 0 People -1 0 1 Do -2 -1 0 1
2 Fine -3 2 1 0 1 2 3
0 0







-1 0 1 1 0 1










-2 1 0 1 2 2 1 0 1 2


-3 2 1 0 1 2 3 3 2 1 0 1 2 3
106
Spin Quantum Number Periodic Table
1/2 -1/2
1/2 1/2













1/2
-1/2






1/2
-1/2
107
Quantum Number Examples
Electron n l ml ms







1
1 0 0 1/2
5
2 1 -1 1/2
2 1 1 -1/2
10
21
3 2 -2 1/2
81
6 1 -1 1/2
5 3 2 1/2
95
118
7 2 -1 -1/2
108
Maximum Number of Electrons
Principle energy level Sublevels available Number of orbitals in sublevel (2 l 1) Number of e- possible in sublevel 2 (2 l 1) Total e- possible for energy level (2n2)
1 s s-1 2 2
2 s, p s-1, p-3 s-2, p-6 8
3 s, p, d s-1, p-3, d-5 s-2, p-6, d-10 18
4 s, p, d, f s-1, p-3, d-5, f-7 s-2, p-6, d-10, f-14 32
5 s, p, d, f, g s-1, p-3, d-5, f-7, g-9 s-2, p-6, d-10, f-14, g-18 50
6 s, p, d, f, g, h s-1, p-3, d-5, f-7, g-9, h-11 s-2, p-6, d-10, f-14, g-18, h-22 72
7 s, p, d, f, g, h, I s-1, p-3, d-5, d-7, g-9, h-11, i-13 s-2, p-6, d-10, f-14, g-18, h-22, i-26 98 Total Poss 280 e-