ATOMIC THEORY

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ATOMIC THEORY

• Building blocks of matter

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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

• IN 400 BC, DEMOCRITUS SAID
• All matter is made of tiny particles called
  atomos
• Disputed by Aristotle

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Why?

• Eminent philosophers of the time, Aristotle and
  Plato, had a different idea.

They favored the earth, fire, air and water
theory of matter. They were more popular, so the
atomos idea was buried for approximately 2000
years.
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Atomos

• To Democritus, atoms were small, hard particles
  like marbles with different shapes and sizes.
• Atoms were infinite in number, always moving and
  capable of joining together.

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For the next 2000 years

• Alchemists tried to make gold from other metals.
• UNTIL
• 1808 JOHN DALTONS NEW ATOMIC THEORY

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LAWS FROM ATOMIC THEORY(from last chapter)

• 1. Law of conservation of matterthe mass of the
  reactants before the reaction equals the mass of
  the products after.
• 2. Law of definite proportionsEvery sample of
  the same compound has the same mass ratio of
  component elements.
• 3. Law of multiple proportionsIn a series of
  compounds of the same two elements, the ratio of
  an element in one compound to another is also a
  small, whole number.

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Daltons Model

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

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Daltons atomic theory

• 1. All matter is made up of atoms
• 2. Atoms of the same element are alike.
• 3. Atoms of different elements are different.
• 4. Compounds have a definite composition by
  weight and combine in small whole number ratios.
• 5. Atoms cannot be subdivided.

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.

• This 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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Thomson Model

• His model of the atom is sometimes called the
  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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Thomson Model

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

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• This surprised Thomson, because the atoms of the
  gas were uncharged. Where had the negative
  charges come from?

Where did they come from?
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He 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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Ernest Rutherford

• In 1908, the English physicist Ernest Rutherford
  was hard at work on an experiment that had 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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Rutherford

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

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Rutherford

• The gold atoms in the sheet were mostly empty
  space. Atoms were not a plum pudding.
• Atom has a small, dense, positively charged
  center that repelled the positive bullets.
• He called the center of the atom the nucleus
• The nucleus is tiny compared to the atom as a
  whole.

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• http//chemmovies.unl.edu/ChemAnime/RUTHERFD/RUTHE
  RFD.html

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Atomic Particles

• Electrondiscovered by Thomson in 1890s
• Robert Millikandetermined the charge of an
  electron in 1909 w/ oil drops
• Protondiscovered by Rutherford in 1911
• Neutrondiscovered by James Chadwick in 1932

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Particles and Charge

• Mass Charge Location
• Proton (p)
• Neutron (no)
• Electron (e-)

1 amu 1 nucleus
1 amu 0 nucleus
1/1840 amu -1 electron cloud
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How Atoms Differ

• Dalton said that all atoms of an element are
  alike, but we know that is not completely true.
  So what is alike?
• All atoms of the same element have the same
  number of protons.
• If the atom is neutral, that means they also have
  the same number of electrons.
• The number of neutrons, however, can vary.

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How Atoms Differ

• Atoms with the same atomic number are the same
  element, but they may have different numbers of
  neutrons.
• Atoms of the same element with a different number
  of neutrons are called isotopes.
• Atoms of the same element with a different number
  of electrons than protons are called ions.

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The Chemists Shorthand Atomic Symbols for
Isotopes
p no
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Mass number ?
K
?? Element Symbol
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Atomic number ?
p or e-
Mass - Atomic no
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Atomic Masses

• Elements occur in nature as mixtures of isotopes
• Atomic mass is the weighted average of all
  isotopes for an element.
• Carbon 98.89 12C
• 1.11 13C
• lt0.01 14C
• Carbon atomic mass 12.01 amu

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MASS NUMBER AND AVERAGE ATOMIC MASS

• Atomic masses are based on CARBON. The atomic
  mass unit is 1/12 of the mass of one carbon atom.
  
• How do we calculate average atomic mass?
• Multiply the times the mass for each isotope,
  then add them together.

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Average atomic mass

• Calculate the average mass of isotopes of
  neptunium with
• 50.0 at 238.05 amu
• 29.4 at 235.1 amu
• 20.6 at 237.98 amu
• (.500 x 238.05) (.294 x 235.1) (.206 x
  237.98) 237.17amu

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Another problem

• Calculate the average atomic mass of calcium with
  these isotopes
• 28.4 at 40.06 amu
• 34.1 at 41.02 amu
• 22.8 at 40.89 amu
• 14.7 at 39.98 amu

(.284x40.06)(.341x41.02)(.228x40.89)(.147x39.98
) 40.56
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One more for Arsenic

• 35.1 of 74.9 amu
• 18.6 of 74.2 amu
• 46.3 of 75.02 amu

74.83 amu
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Atomic Mass

• Atomic mass is the weighted average of all of the
  known isotopes of an element, so will always be
  shown as a decimal number.

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Covalent Chemical Bonding

• The forces that hold atoms together in compounds.
  Covalent bonds result from atoms sharing
  electrons between nonmetal atoms.
• Molecule a collection of covalently-bonded
  atoms.
• Atom representative particle for a monatomic
  element

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Ionic Chemical Bonding

• Cation A positive ion
• Mg2, NH4
• Anion A negative ion
• Cl?, SO42?
• Ionic Bonding Force of attraction between
  oppositely charged ions. Smallest particle
  called a formula unit.

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Stupendous Seven
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Periodic Table

• Elements classified by
• properties
• atomic number
• Groups (vertical columns)also called families
• 1A alkali metals
• 2A alkaline earth metals
• 7A halogens
• 8A noble gases
• Periods (horizontal rows)

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Periodic Table

• Antoine Lavoisier , 1790s made first list of
  known elements, 23 total. By 1870, there were
  70!
• John Newlands, 1864Law of Octaves When element
  were placed in order of increasing atomic mass,
  every 8th element repeated properties.
• Lothar Meyer, 1869Periodic table based on
  physical characteristics only and increasing
  atomic mass.
• Dmitri Mendeleev, 1869Periodic table based on
  physical and chemical characteristics and
  increasing atomic mass. Predicted new elements.
• Henry Moseley, 1913Modern periodic law based on
  subatomic particles There is a periodic
  repetition of chemical and physical properties of
  the elements when they are arranged by increasing
  atomic number (protons).

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Periodic Trends
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Periodic Trends

• Groups Periods

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Periodic Trends
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Periodic Trends
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Periodic Trends