AP Notes Chapter 2

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AP Notes Chapter 2

• Atoms and Elements

History of the atom Summed-up

• Greeks
• Democritus and Leucippus - atomos
• Aristotle- elements.
• Alchemy
• 1660 - Robert Boyle- experimental definition of
  element.
• Lavoisier- Father of modern chemistry.

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Daltons Atomic Theory

1. Elements are made up of atoms
2. Atoms of each element are identical. Atoms of
   different elements are different.
3. Compounds are formed when atoms combine. Each
   compound has a specific number and kinds of atom.
4. Chemical reactions are rearrangement of atoms.
   Atoms are not created or destroyed.

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The Atom

• Dalton (early 1800s)

indivisible
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A Helpful Observation

• Gay-Lussac- under the same conditions of
  temperature and pressure, compounds always react
  in whole number ratios by volume.
• Avagadro- interpreted that to mean
• at the same temperature and pressure, equal
  volumes of gas contain the same number of
  particles.
• (called Avagadros Hypothesis)

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Experiments theories to determine what an atom
was

• John Dalton- atoms indivisible
• J. J. Thomson- Cathode ray tubes, electrons
• Marie Curie- radioactivity
• Robert Millikan- electron mass charge
• Ernest Rutherford- protons
• James Chadwick- neutrons

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Thomsons Experiment

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Thomsons Experiment

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Thomsons Experiment

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• Passing an electric current makes a beam appear
  to move from the negative to the positive end.

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Thomsons Experiment

• By adding an electric field

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Thomsons Experiment

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• By adding an electric field, he found that the
  moving pieces were negative

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The Atom

• Thompson ( 1900)

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cloud of () charge
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electron(-) charge
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plum pudding model
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Thomsoms Model

• Found the electron.
• Couldnt find positive (for a while).
• Said the atom was like plum pudding.
• A bunch of positive stuff, with the electrons
  able to be removed.

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Millikans Experiment
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Millikans Experiment
X-rays
X-rays give some electrons a charge.
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Millikans Experiment

• Some drops would hover

From the mass of the drop and the charge on the
plates, he calculated the mass of an electron
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Radioactivity

• Discovered by accident
• Henri Bequerel photographic plates
• Marie Curie studied named it
• Three types
• alpha- helium nucleus (2 charge, large mass)
• beta- high speed electron
• gamma- high energy light

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James Chadwick

• Neutrons
• Particles from radioactive polonium hit a
  beryllium target and produced particle
• no charge
• slightly greater mass than the proton

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Rutherfords Experiment

• Used uranium to produce alpha particles.
• Aimed alpha particles at gold foil by drilling
  hole in lead block.
• Since the mass is evenly distributed in gold
  atoms alpha particles should go straight through.
• Used gold foil because it could be made atoms
  thin.

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Rutherfords Experiment
Florescent Screen
Lead block
Uranium
Gold Foil
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Rutherfords Experiment
What he expected
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Rutherfords Experiment
Because
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Rutherfords Experiment
Because, he thought the mass was evenly
distributed in the atom.
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Rutherfords Experiment
What he got
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Rutherfords Experiment
How he explained it

• Atom is mostly empty
• Small dense,
• positive pieceat center.
• Alpha particlesare deflected by
• it if they get close
• enough.

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Rutherfords Experiment
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Gold Foil Experiment

• Rutherford (1911) Nuclear Model

heavy central() nucleus
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e- aboutnucleus
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The Atom

• Rutherford ( 1911)

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heavy central() nucleus
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(e-) aboutnucleus
Nuclear Model
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The Atom
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• Bohr ( 1913)

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e- inallowed orbits
central ()nucleus
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n3
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n2
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n1
Planetary Model
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Modern View

• The atom is mostly empty space.
• Two regions
• Nucleus- protons and neutrons.
• Electron cloud- region where you might find an
  electron.

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The Atom

• Heisenberg, de Broglie, Schroedinger (mid 1920s)

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e- in regionsdefined by mathfunctions
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Quantum Mechanical Model
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Sub-atomic Particles

• Z - atomic number number of protons determines
  type of atom.
• A - mass number number of protons neutrons.
• Number of protons number of electrons if
  neutral.

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Nuclear Symbols Notation
A
X
Mass Number?
?Element Symbol
Z
Atomic Number?
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Na
Na
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11
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Isotopes of elements

• Isotopes are forms of an atom that differ by the
  number of neutrons
• Mass number is approximation of exact atomic mass
  of an isotope
• Atomic mass or atomic weight is the average mass
  of the isotopes of atoms
• Isotopic percent abundance or fractional
  abundance is a description of the proportion of
  an isotope in a sample of an element

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Atomic Mass

• Atoms are so small, it is difficult to discuss
  how much they weigh in grams.
• Use atomic mass units.
• an atomic mass unit (amu) is one twelth the mass
  of a carbon-12 atom.
• This gives us a basis for comparison.
• The decimal numbers on the table are atomic
  masses in amu.

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They are not whole numbers

• Because they are based on averages of atoms and
  of isotopes.
• can figure out the average atomic mass from the
  mass of the isotopes and their relative
  abundance.
• add up the percent as decimals times the masses
  of the isotopes.

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Isotopes of Hydrogen
hydrogen deuterium tritium
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Examples

• There are two isotopes of carbon 12C with a mass
  of 12.00000 amu(98.892), and 13C with a mass of
  13.00335 amu (1.108).
• There are two isotopes of nitrogen , one with an
  atomic mass of 14.0031 amu and one with a mass of
  15.0001 amu. What is the percent abundance of
  each?

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Percent Abundance Percent abundance number of
atoms of a given isotope x 100
total number of atoms of all isotopes
Fractional Abundance Fractional abundance
Percent Abundance 100

• Atomic Weight
• (abundance isotope 1)(weight isotope1)
• (abundance isotope 2)(weight isotope2)
• or

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• A portion of an atoms mass of protons, neutrons
  and electrons is converted to energy that holds
  the atom together.
• Einstein gave us ?E (?m)C2
• The loss of this mass as the atom forms is called
  the mass defect. This missing mass is converted
  to binding energy (BE)
• Mass atom BE pro. elec. neu.

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Allotrope Different forms of the same element
that exist in the same physical state under the
same conditions of Temperature Pressure
Carbon

• Diamond
• Graphite

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Graphite Diamonds Buckyballs
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Periodic Table
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Metals

• Conductors
• Lose electrons
• Malleable and ductile

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Nonmetals

• Brittle
• Gain electrons
• Covalent bonds

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Semi-metals or Metalloids
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Alkali Metals
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Alkaline Earth Metals
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Halogens
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Transition metals
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Noble Gases
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Inner Transition Metals
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Periodic Table
1A
8A
Families or Groups
2A
3A 4A 5A 6A 7A
3B 4B 5B 6B 7B 8B 1B 2B
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Periodic Table
Periods
1 2 3 4 5 6 7 8
Lanthanide Series Actinide Series
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Periods and Groups or Families
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Hydrogen

• Shuttle main engines use H2 and O2

The Hindenburg crash, May 1939.
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Group 1A Alkali Metals
Potassium
Reaction of potassium H2O
Cutting sodium metal
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Group 2A Alkaline Earth Metals
Magnesium
Magnesium Ablaze!
Magnesium oxide
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Calcium CarbonateLimestone
Champagne cave carved into chalk in France
The Appian Way, Italy
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Group 3A B, Al, Ga, In, Tl
Aluminum
Boron halides BF3 BI3
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Gems Minerals

• Sapphire Al2O3 with Fe3 or Ti3 impurity gives
  blue whereas V3 gives violet.
• Ruby Al2O3 with Cr3 impurity

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Transition Elements

• Lanthanides and actinides

Iron in air gives iron(III) oxide
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Colors of Transition Metal Compounds
Nickel
Cobalt
Copper
Zinc
Iron
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Group 4A C, Si, Ge, Sn, Pb
Quartz, SiO2
Diamond
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Group 5A N, P, As, Sb, Bi

• White and red phosphorus

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Phosphorus

• Phosphorus first isolated by Brandt from urine,
  1669

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Group 6A O, S, Se, Te, Po

• Sulfuric acid dripping from snot-tite in cave in
  Mexico

Sulfur from a volcano
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Group 7A F, Cl, Br, I, At
Halogen
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Group 8A He, Ne, Ar, Kr, Xe, Rn

• Lighter than air balloons
• Neon signs