Investigating Atoms and Atomic Theory

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Investigating Atoms and Atomic Theory

• Students should be able to
• Describe the particle theory of matter.
• Use the Bohr model to differentiate among the
  three basic particles in the atom (proton,
  neutron, and electron) and their charges,
  relative masses, and locations.
• Compare the Bohr atomic model to the electron
  cloud model with respect to their ability to
  represent accurately the structure of the atom.

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Atomos Not to Be Cut

• The History of Atomic Theory

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Atomic Models

• This model of the atom may look familiar to you.
  This is the Bohr model. In this model, the
  nucleus is orbited by electrons, which are in
  different energy levels.
• A model uses familiar ideas to explain unfamiliar
  facts observed in nature.
• A model can be changed as new information is
  collected.

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• The atomic model has changed throughout the
  centuries, starting in 400 BC, when it looked
  like a billiard ball ?

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Who are these men?
In this lesson, well learn about the men whose
quests for knowledge about the fundamental nature
of the universe helped define our views.
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Democritus
400 BC

• This is the Greek philosopher Democritus who
  began the search for a description of matter more
  than 2400 years ago.
• He asked Could matter be divided into smaller
  and smaller pieces forever, or was there a limit
  to the number of times a piece of matter could be
  divided?

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Atomos

• His theory Matter could not be divided into
  smaller and smaller pieces forever, eventually
  the smallest possible piece would be obtained.
• This piece would be indivisible.
• He named the smallest piece of matter atomos,
  meaning not to be cut.

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Atomos

• To Democritus, atoms were small, hard particles
  that were all made of the same material but were
  different shapes and sizes.
• Atoms were infinite in number, always moving and
  capable of joining together.

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• This theory was ignored and forgotten for
  more than 2000 years!

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Why?

• The eminent philosophers of the time, Aristotle
  and Plato, had a more respected, (and ultimately
  wrong) theory.

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Robert Boyle

• Robert Boyle 1660
• Boyle is best known for his quantitative work
  with gases.
• He also was the first to propose the existence of
  elements in the modern sense.
• Boyle considered a substance to be an element
  unless it can be broken down into simpler
  substances.

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Lavoisier 1760

• Marie and Antoine Lavoisier studied chemical
  reactions quantitatively.
• They are credited with being the first to propose
  the law of conservation of matter.

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Lavoisier 1760

• Law of Conservation of Matter - Matter is neither
  created nor destroyed in a chemical reaction.

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Daltons Model

• In the early 1800s, the English Chemist John
  Dalton performed a number of experiments that
  eventually led to the acceptance of the idea of
  atoms.

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Daltons Theory

• All matter is composed of atoms.
• Atoms cannot be made or destroyed.
• All atoms of the same element are identical.
• Different elements have different types of
  atoms.
• Chemical reactions occur when atoms are
  rearranged.
• Compounds are formed from atoms of the
  constituent elements.

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Daltons Theory

• became one of the foundations of modern
  chemistry.

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Thomsons Plum Pudding Model

• In 1897, the English scientist J.J. Thomson
  provided the first hint that an atom is made of
  even smaller particles.

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Thomsons Plum Pudding Model

• Atoms were made from a positively charged
  substance with negatively charged electrons
  scattered about, like raisins in a pudding.

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Thomsons Plum Pudding Model

• Thomson studied the passage of an electric
  current through a gas. As the current passed
  through the gas, it gave off rays of negatively
  charged particles.

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Thomson concluded that the negative charges came
from within the atom. A particle smaller than
an atom had to exist. The atom was divisible!

• Thomson called the negatively charged
  corpuscles, today known as electrons.
• Since the gas was known to be neutral, having no
  charge, he reasoned that there must be positively
  charged particles in the atom.
• But he could never find them.

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Prousts CuCO3 Study

• In 1808 Joseph-Louis Proust, a French chemist,
  studied the chemical composition of the compound
  copper carbonate (CuCO3).

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Prousts CuCO3 Study

• Study proved that the relative quantities of any
  given pure chemical compound's constituent
  elements remain invariant, regardless of the
  compound's source.
• Supports John Dalton's "law of definite
  proportions"
• Law of Definite Proportions
• A given compound always contains the same
  proportion of elements by mass.

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Structure of Atoms

• Scientist began to wonder what an atom was like.
• Was it solid throughout with no internal
  structure or was it made up of smaller, subatomic
  particles?
• It was not until the late 1800s that evidence
  became available that atoms were composed of
  smaller parts.

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Rutherfords Gold Foil Experiment

• In 1908, the English physicist Ernest Rutherford
  was hard at work on an experiment that seemed to
  have little to do with unraveling the mysteries
  of the atomic structure.

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• Rutherfords experiment Involved firing a stream
  of tiny positively charged particles at a thin
  sheet of gold foil (2000 atoms thick)

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Rutherfords Gold Foil Experiment

• Most of the positively charged bullets passed
  right through the gold atoms in the sheet of gold
  foil without changing course at all.
• Some of the positively charged bullets,
  however, did bounce away from the gold sheet as
  if they had hit something solid. He knew that
  positive charges repel positive charges.

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http//chemmovies.unl.edu/ChemAnime/RUTHERFD/RUTHE
RFD.html
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Rutherfords Gold Foil Experiment

• Rutherford reasoned that all of an atoms
  positively charged particles were contained in
  the nucleus. The negatively charged particles
  were scattered outside the nucleus around the
  atoms edge.

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Millikans Oil Drop Experiment

• In 1909, Robert Millikans oil drop experiment
  allowed him to determine the charge on an
  electron.
• This charge can be plugged into Thomsons formula
  the mass of the electron calculated
• Mass electron 9.11 x 10-31 kg
• e 1.60 x 10-19 coulombs

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Millikans Oil Drop Experiment

• Based on balancing forces the gravitational pull
  down on an oil drop and the electric force up on
  ionized particles.
•  
• http//webphysics.davidson.edu/applets/pqp_preview
  /contents/pqp_errata/cd_errata_fixes/section4_5.ht
  ml

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Bohr Model

• In 1913, the Danish scientist Niels Bohr proposed
  an improvement. In his model, he placed each
  electron in a specific energy level.

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Bohr Model

• According to Bohrs atomic model, electrons move
  in definite orbits around the nucleus, much like
  planets circle the sun. These orbits, or energy
  levels, are located at certain distances from the
  nucleus.

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Wave Model
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Schrödingers Wave Equation

• Todays atomic model is based on the principles
  of wave mechanics.
• According to the theory of wave mechanics,
  electrons do not move about an atom in a definite
  path, like the planets around the sun.

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The Wave Model

• In fact, it is impossible to determine the exact
  location of an electron. The probable location of
  an electron is based on how much energy the
  electron has.
• According to the modern atomic model, at atom has
  a small positively charged nucleus surrounded by
  a large region in which there are enough
  electrons to make an atom neutral.

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Electron Cloud

• A space in which electrons are likely to be
  found.
• Electrons whirl about the nucleus billions of
  times in one second
• They are not moving around in random patterns.
• Location of electrons depends upon how much
  energy the electron has.

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Electron Cloud

• Depending on their energy they are locked into a
  certain area in the cloud.
• Electrons with the lowest energy are found in the
  energy level closest to the nucleus
• Electrons with the highest energy are found in
  the outermost energy levels, farther from the
  nucleus.

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Three Chemistry Laws

• Conservation of Mass
• Definite Proportions
• Multiple Proportions

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The Discovery of Atomic Structure
Remember...

• Thompson - Cathode rays.
• Milliken - Oil drops.
• Rutherford - backscattering ?-particles.
• Radioactivity, the spontaneous emission of
  radiation from an atom led to the discovery of
  ?-, ?-, and ?-rays.

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Atomic Theory Of Matter (Dalton's Theory)
and...

• Law of definite proportions led to theory that
  all matter made up of atoms.
• Atoms- basic building blocks and don't change
  when react with other atoms.
• Element- describes matter composed of only one
  type of atom.
• Compound- combination of atoms in specific
  proportions.
• Chemical reaction- atoms exchange partners
  producing other compounds.

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Foundations of Atomic Theory
Law of Conservation of Mass Mass is neither
destroyed nor created during ordinary chemical
reactions.
Law of Definite Proportions The fact that a
chemical compound contains the same elements in
exactly the same proportions by mass regardless
of the size of the sample or source of the
compound.
Law of Multiple Proportions If two or more
different compounds are composed of the same two
elements, then the ratio of the masses of the
second element combined with a certain mass of
the first elements is always a ratio of small
whole numbers.
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Law of Multiple ProportionsJohn Dalton (1766
1844)

• If two elements form more than one compound,
  the ratio of the second element that combines
  with 1 gram of the first element in each is a
  simple whole number.
• e.g. H2O H2O2
• water hydrogen peroxide
• Ratio of oxygen is 12 (an exact ratio) 22

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Law Of Multiple Proportions

• Some elements can form more than one compound
  when they react together
• C O CO and CO2
• N O N2O, NO, NO2,
• Daltons law predicted that the mass proportions
  should be proportional.

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Law of Conservation of MassLavoisier (1743 -
1794)

• The mass of substances produced (products) by a
  chemical reaction is always equal to the mass of
  the reacting substances (reactants).
• 1.00g carbon 5.34g sulfur ? 6.34g carbon
  disulfide

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Conservation of Atoms
2 H2 O2
2 H2O
4 atoms hydrogen 2 atoms oxygen
4 atoms hydrogen 2 atoms oxygen
Dorin, Demmin, Gabel, Chemistry The Study of
Matter , 3rd Edition, 1990, page 204
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Legos are Similar to Atoms
Legos can be taken apart and built into many
different things.
Atoms can be rearranged into different substances.
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Law of Definite ProportionsJoseph Louis Proust
(1754 1826)

• Each compound has a specific ratio of elements.
• It is a ratio by mass
• Water is always 8 grams of oxygen for every one
  gram of hydrogen

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Law of Definite Proportions
Whether synthesized in the laboratory or obtained
from various natural sources, copper carbonate
always has the same composition. Analysis of
this compound led Proust to formulate the law of
definite proportions.
Pictures only
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Law of Definite Proportions

• Water is 8 grams of oxygen per gram of hydrogen.
• Hydrogen peroxide is 16 grams of oxygen per gram
  of hydrogen.
• 16 g to 8 g is a 21 ratio
• True, because you have to add a whole atom, you
  cant add a piece of an atom.