CH 2: Atoms, Molecules, and Ions

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CH 2 Atoms, Molecules, and Ions
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Chapter Outline

• History of chemistry (2.1 2.4)
• Chemical laws
• Path to the atom
• Modern atomic structure (2.5)
• Molecules vs. Ions (2.6)
• Naming molecular and ionic compounds (2.8)
• Introduction to the periodic table (2.7)

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History of Chemistry

• Greek Philosophers 5th Century BCE (BCE before
  the common era - replaces BC)
• The Greek philosophers were the first to reflect
  on the nature of matter. 
• They proposed that all matter is made out of
  first 4 elements -- earth, air , water, fire. 
  Aristotle added a fifth element, plasma (also
  called ether). 

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Greek Philosophers

• Democritus had an alternate view of matter.  He
  proposed that matter was  made up of tiny
  particles called atoms. 
• His "theory" was not well accepted at the time.

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Alchemy

• Alchemy 600-1600's CE
• (CE common era, replaces AD)
• Alchemy developed at about the same time in
  China, India, and Greece. It spread into Europe
  in the 8th century.

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Alchemy

• Alchemists had two pursuits
• Search for a means to convert base metals into
  gold
• Search for the elixir of life
• Substance that would lead to immortality

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Alchemy

• Advances from Alchemy
• Many new substances where identified
• Plaster of Paris, nitric acid.
• New lab techniques and equipment developed
• New medicines identified

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Modern Chemistry, 1600 on

• First chemists/physicists to use scientific
  method
• Boyle - elements
• Lavoisier law of conservation of matter
• Proust - law of definite proportions
• Dalton law of multiple proportions, atomic
  theory
• Avogadro - hypothesis
• Thomson charge to mass ratio for an electron
• Millikan charge on the electron
• Bequerel and the Curies - radioactivity
• Rutherford nuclear atom

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Modern Chemistry

• Robert Boyle 1660
• Proposed a substance to be an element unless it
  can be broken down into simpler substances.
• Proposed one of the gas laws CH 5

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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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• Proust late 1700s
• Law of Definite Composition/Proportions
• A given compound always contains the same
  proportion of elements by mass.

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• John Dalton 1800
• Law of Multiple Proportions
• When two elements form more than one compound,
  the ratios of the masses of the second element
  that combines with one gram of the first element
  can always be reduced to small whole numbers.
  (see page 43)
• Page 70 36

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• Dalton also proposed the first table of atomic
  masses
• Most masses later need revision
• Dalton is best known for proposing Atomic Theory

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Daltons Atomic Theory

• Elements are made up of tiny particles called
  atoms.
• Atoms are indivisible and indestructible
• Atoms of a given element are identical
• Atoms of different elements differ in some
  fundamental way(s)

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Daltons Atomic Theory

• Compounds form when atoms of different elements
  combine with each other.
• A given compound always has the same relative
  number and types of elements.

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Daltons Atomic Theory

• Chemical reactions occur when atoms change how
  they are bound to each other.
• Individual atoms are not changed, just rearranged

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Avogadro 1811

• Avogadro's Hypothesis
• At the same temperature and pressure equal
  volumes of gases contain the same number of
  particles.
• Based on Guy-Lussacs data
• See page 45/46

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From Dalton to Atomic Structure

• Daltons atomic theory lead to much research on
  the nature of the atom.
• This research showed the atom to made up of
  smaller particles.

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J.J. Thompson

• Thomson measured the deflection of a cathode ray
  beam in electrical and magnetic fields of known
  strengths.

Applied electrical field
Cathode ray

() Metal electrode
(-) Metal electrode
-
Cathode ray tube experiment, pg 48
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• Thompson found the cathode rays were attracted by
  the positive charge and repelled by negative
• These findings clearly indicated that the rays
  consisted of negatively charged particles.
• Today we know these particles as electrons.

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• Thompson measured the deflection of the beam in a
  magnetic field and determined the chargemass
  ratio for an electron
• e - 1.76 x 108 C/g
• m
• e charge on the electron in coulombs
• M mass of an electron in grams

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• Thomson also found that the cathode ray particles
  were identical regardless of source.
• Concluded all elements contain these negative
  particles (electrons)

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J.J. Thomson

• Thomson
• identified cathode ray beams as a stream of
  negatively charged particles
• calculated the charge to mass ratio for these
  negatively charged particles
• proposed the existence of positively charged
  particles
• To balance the negative charge of the electrons

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Millikan 1909

• Millikans oil drop experiment allowed him to
  determine the charge on an electron
• This charge can be plugged into Thomsons formula
  and the mass of the electron calculated
• Mass electron 9.11 x 10-31 kg
• Page 49

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Radioactivity

• Becquerel, Marie and Pierre Curie 1896
• Henri Becquerel - observed the natural emission
  of energy/rays by uranium.
• Marie and Pierre Curie studied Becquerel's
  rays.
• The Curies findings suggested that matter was
  composed of smaller particles than atoms.
• The Curies coined the term radioactivity to
  describe the rays emitted.

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Radioactivity

• Three types of radioactivity were identified
• gamma rays - very high energy light
• beta particles - high energy electrons
• alpha particles - He2 particles
• 2 protons and 2 neutrons

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• In the early 1900s the accepted model of the
  atom was called the plum pudding model of the
  atom
• Electrons (tiny and negatively charged) were
  pictured to be dispersed in a cloud of positive
  charge.

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Rutherford and the Nuclear Atom

• In 1911 an experiment conducted in Ernest
  Rutherfords lab showed the plum pudding model
  to be incorrect.
• Experiment was conducted by Geiger and Marsden
  and the findings interpreted by Rutherford.
• See page 49

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Rutherfords Atom

• First to propose a nuclear atom.
• An atom has a dense positive center containing
  all of positive charge and most of the mass of
  the atom the nucleus
• Electrons are scattered in the empty space around
  the nucleus
• Electrons occupy a volume that is huge as
  compared to the size of the nucleus.

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A New Model of the Atom

• Expected based on
• Plum pudding model
• Rutherfords model
• Based on his results

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Modern Atomic Structure

• Rutherford continued to study the atom and the
  positive matter of the atom.
• 1919, particle named the proton
• 1932 James Chadwick proposed the existence of a
  third subatomic particle, the neutron.

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Subatomic Particles
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Mass of Subatomic Particles

• Protons and neutrons have the same mass (in the
  range of 10 -27 kg).
• Mass of each and of individual atoms is often
  expressed in amu rather than grams
• Atomic mass unit (amu) 1/12 the mass of a
  carbon-12 atom

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Mass of Subatomic Particles

• The mass of the electron (10-31 kg) is tiny as
  compared to that of the proton and neutron (10-27
  kg) .
• Therefore, the electrons mass is considered to
  be 0 amu when calculating the mass of an atom.

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Subatomic Particles and the Elements

• Each element has a unique number of protons.
• Number of protons defines the element.

Atomic protons
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C
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Subatomic Particles

• Since atoms are neutral, for every proton there
  is a/n _________.
• When atoms interact to form compounds, it is
  their ___________ that interact.

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Terms

• Mass number sum of the of protons and the
  neutrons in the nucleus of an atom
• FOR MOST ELEMENTS THE MASS NUMBER IF NOT ON THE
  PERIODIC TABLE.
• You will be given enough information to determine
  mass number or number of neutrons.

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Terms

• Isotopes atoms of a given element that differ
  in mass number
• Isotopes have the same number of _____________.
• Isotopes differ in the number of _______.

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Isotopes

• Writing atomic symbols for isotopes
• pg 50

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B
Mass
Symbol for element
5
Atomic
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FAQ - Isotopes

• When is mass number found on the periodic table?
• Whats the atomic mass? Is it the same as the
  mass number?

C
84 (209)
6 12.0107
Po
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Molecules and Ions (2.6)

• Atoms of different elements combine to form
  compounds
• Atoms in compounds are held together by chemical
  bonds.
• Bonds involve interactions of the bonding atoms
  ________

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Bonding

• There are two types of bonds
• Covalent bonds bonding atoms share electrons
• Atoms are always nonmetal atoms
• Covalently bonded atoms form molecules
• Ways to represent molecules
• Chemical formula H2O
• Structural formula
• 
  H
• O O

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Bonding

• Ionic bonds attractive force among oppositely
  charged ions
• Bond formed between metal cations and nonmetal
  anions
• No molecules involved

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Ions - Terms

• Ion charged atom or group of atoms
• Formed when atoms gain or lose electrons
• Cation positively charged ion
• Formed when an atom _______ electrons
• Anion negatively charged ion
• Formed when an atom ______ electrons

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Ions

• Describing ion formation
• Cation example
• Anion example

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Naming Binary Compounds

• Binary compounds compound composed of 2
  elements
• NaCl
• CO
• CO2

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Types of Binary Compounds

• Type I binary ionic compounds
• Metal forms only one ion
• Type II binary ionic compounds
• Metal forms more than one ion
• Use roman numerals to indicate the charge on the
  ion
• Type III binary covalent compounds
• Compound between 2 nonmetals

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Types I Binary Compounds

• Compound between a metal and a nonmetal
• Metal forms only one ion
• Name the cation and then the anion.
• Name of the cation is the name of the element
• Name of the anion is the name of the nonmetal
  with the ending changed to ide

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Monoatomic cations to know
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Monoatomic anions to know
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Practice

• Name ? chemical formula
• Chemical formula ? name