Chapter 2: Atoms, Molecules and Ions

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Chapter 2 Atoms, Molecules and Ions
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Early Models of Atoms

• Democritus (460-400B.C.) first suggested the
  existence of these particles, which he called
  atoms for the Greek word for uncuttable. They
  lacked experimental support due to the lack of
  scientific testing at the time.
• Plato and Aristotle formulated the notion that
  there can be no ultimately indivisible particles,
  so the atomic view faded for a number of years.
  
• John Dalton (1766-1844) performed experiments to
  study the ratios in which elements combine in
  chemical reactions. He formulated hypotheses and
  theories to explain his observations, which
  became Daltons Atomic Theory.
• All elements are composed of tiny indivisible
  particles called atoms.
• Atoms of the same element are identical. The
  atoms of any one element are different from those
  of any other element.
• Chemical reactions occur when atoms are
  separated, joined or rearranged. Atoms of one
  element, however, are never changed into atoms of
  another element as a result of a chemical
  reaction.
• Atoms of different elements can physically mix
  together or combine in simple, whole number
  ratios to form compounds.

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Daltons Atomic Theory

• According to Daltons atomic theory atoms are the
  smallest particles of an element that retain the
  chemical identity of the element. His theory
  explains several simpler laws of chemical
  combination from his time.
• Law of constant composition In a given compound,
  the relative numbers and kinds of atoms are
  constant. (4)
• Law of conservation of mass The total mass of
  materials present after a chemical reaction tis
  the same as the total mass present before the
  reaction (3)
• Law of multiple proportions If two elements A
  and B combine to form more than one compound, the
  masses of B that can combine with a given mass of
  A are in the ratio of small whole numbers.
• Example CO2 and CO H2O2 and H2O

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Discovery of Atomic Structure

• As scientists began to develop methods for more
  detailed probing of the nature of matter, we
  discovered more. Atoms are now known to be
  divisible as they can be broken down to even
  smaller particles by atom smashers.
• J.J. Thomson (1856-1940) discovered electrons
  using cathode ray tubes. Another CRT
• Robert Millikan (1868-1953) carried out
  experiments to determine the charge of an
  electron (-). He also determined the ratio of the
  charge to the mass of an electron.
• In 1886, E. Goldstein observed a cathode ray
  tube and found rays traveling in the opposite
  direction to that of the cathode rays. He called
  these rays canal rays and concluded that they
  must be positive particles, which are now called
  protons.
• In 1932, James Chadwick confirmed the existence
  of yet another subatomic particle the neutron.
  Neutrons are subatomic particles with no charge
  but with a mass nearly equal to that of a proton.
  See simulation

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• After discovering these subatomic particles,
  scientists wondered how they were put together.
• JJ Thompson thought since the electrons
  contributed such a small fraction of the atoms
  mass, they were probably an equal fraction of it
  size so it was like Plum Pudding.
• In 1911, Ernest Rutherford and his coworkers
  performed the Gold Foil Experiment to further
  study the phenomenon.
• Concluded that most of the mass of each atom and
  all of its positive charge reside in a very
  small, extremely dense region which is called the
  nucleus. The rest of the atom is mostly empty
  space.

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Modern View of Atomic Structure

• Since the time of Rutherford, physicists have
  learned much about the nucleus. Although many
  other parts have been discovered, chemists tend
  to only work with three main particles since they
  determine chemical behavior Electron, Neutron
  and Proton
• Electron has a charge of -1.602 X 10-19 C and a
  proton has a charge of 1.602 X 10-19 C so this
  quantity of Coulombs is known as one electronic
  charge and atomic and subatomic particles usually
  have a charge that is multiples of this. Neutrons
  have no charge and are electrically neutral.
• Atoms have extremely small masses so instead of
  using the real numbers, atomic mass units (amus)
  are used. Protons and neutrons are very similar
  in mass but it would take 1836 electrons to equal
  1 proton so most of an atoms mass is in the
  nucleus.
• Atoms are also extremely small with diameters
  between 1 X 10-10 and 5 X 10-10 so they are
  usually expressed with angstroms, which is 10-10.
  

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Illustrating the Size of an Atom

• The diameter of a US penny is 19 mm. The diameter
  of a silver atom, by comparison is only 2.88 A.
  How many silver atoms could be arranged side by
  side in a straight line across the diameter of a
  penny?
• 19 mm 1X10-3m 1A 1 Ag atom 6.6 X107 Ag
  atoms
• 1mm 1X10-10m 2.88 A
• This is over 66 million silver atoms could sit
  side by side across a penny!

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Atomic Number

• The number of protons in the nucleus of an atom
  of that element, which is the primary difference
  that distinguishes each element.
• For an atom with no charge, this is also the
  number of electrons since the positive charge of
  the protons cancels the negative charge of the
  electrons.

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Mass Number

• Most of the mass of an atom is found in the
  nucleus so the total number of protons and
  neutrons equals the mass number.
• If you know the atomic number and mass number you
  can determine the composition of that atom.
• The composition can be represented by the
  shorthand notation using the element symbol,
  atomic number and mass number.
• For gold, Au is the symbol for the element and
  the atomic number is subscript and mass number is
  superscript on the left side.

Au
197
79

• Do Sample Exercises 2.3 and Practice Exercises
  in that box on pg 46

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Isotopes

• Atoms that have the same number of protons but
  different number of neutrons.
• Affects the shorthand notation of the element.
• Do Sample Exercises 2.3 and Practice Exercises in
  that box on pg 46

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Atomic Mass

• Today we can determine the masses of individual
  atoms with a relative high degree of accuracy but
  since they are so small atomic mass units are
  used with hydrogen being 1 amu.
• The average atomic mass for an element due to the
  different isotopes, the mass of those isotopes
  and the natural percent abundance. It is also
  known as atomic weight.
• Add up the different atomic mass of each atom and
  then divide by the number of atoms.
• Or, multiply mass by and then determine average
  mass.
• Sample Exercise 2.4 and Practice Exercise on pg 47

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Mass Spectrometer

• The most direct and accurate means for
  determining atomic and molecular weights. See pg
  48

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Periodic Table

• The arrangement of elements in order of
  increasing atomic number, with elements having
  similar properties placed in vertical column.
• Atomic number, symbol, name, atomic weight are
  found in each square for each element. Some
  tables have additional information as well.
  Example
• Can be arranged according to metals, non-metals
  and metalloids, solid liquid and gases, and by
  family. Example

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Molecules and Molecular Compounds

• Even though the atom is the smallest
  representative sample of an element, only the
  noble gas elements are normally found in nature
  as isolated atoms. All others form either
  molecules or ions.
• A molecule is an assembly of two or more atoms
  tightly bound together by a covalent bond created
  by two atoms sharing electrons.
• Diatomic atoms form diatomic molecules (remember
  7 start at 7 form a 7 and hydrogen).
• Compounds that are composed of molecules that
  contain more than one type of element are
  molecular compounds.
• Most molecules are composed of nonmetals.
• Chemical formulas that indicate actual number and
  types of atoms in a molecules are called
  molecular formulas. Such as H2O, C6H12O6, and
  C2H4.
• Empirical formulas give only the relative number
  of atoms, they are basically the reduced formula.
  Such as H2O, CH2O, and CH2.
• Do Sample Exercise 2.6 and Practice Exercise on
  pg 53.

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Picturing Molecules

• The molecular formula of a substance describes
  the composition but doesnt show how they come
  together.
• Structural formula shows which atoms are
  attached to which.
• Atoms are represented by their symbol and the
  bonds are represented by lines.
• Perspective Drawing shows actual geometry to
  give some sense of three-dimensional shape.
• Ball-and-stick Models shows atoms a spheres and
  bond as sticks. Accurately represents the angles
  at which the atoms are attached to one another
  within the molecules.
• Space-filling Model shows what the molecule
  would look like if the atoms were scaled up to
  size.

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Ions

• Some atoms can gain or lose electrons to try and
  get the same number of electrons as the nearest
  noble gas, when an electron is gained or lost
  from a neutral atom a charged particle occurs
  called an ion.
• An ion with a positive charge (lost an electron)
  is called a cation, where as an ion with a
  negative charge (gained an electron) is called an
  anion.
• In general, metals atoms tend to lose electrons
  to form cations and nonmetals tend to gain
  electrons to form anions.
• In addition to simple single atom ions, there are
  polyatomic ions, which consist of atoms joined as
  a molecule but they have a net positive or
  negative charge.
• Ionic charge can be predicted by determining how
  many electrons an atom has to lose to become like
  the nearest stably arranged noble gas.
• Do Sample Exercises 2.7 and 2.8 and Practice
  Problems on pg 55.

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Ionic Compounds

• When a positive ion such as Na comes close to a
  negative ion such as Cl, their opposite charges
  are attracted and form an ionic compound
  connected by a ionic bond.
• Generally, they are combinations of metals and
  nonmetals such as Na and Cl.
• Ions in ionic compounds are arranged in
  three-dimensional structures.
• The formula for an ionic compound is always an
  empirical formula (most reduced form) because
  there is no discrete molecule of NaCl.
• Chemical compounds are always electrically
  neutral, so the empirical formula shows the ratio
  of the ions for this to be true.
• For example, Mg2 and N3- would have to be Mg3N2.
  
• Sample Exercise 2.10 and Practice Exercises on pg
  58

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Naming Inorganic Compounds
Lithium Fluoride

• To obtain information about a particular
  substance you must know its chemical name and
  formula, the system used for this is chemical
  nomenclature. Some compounds also have common
  names in addition to their chemical nomenclature
  such as water.
• The rules for naming a compound is based on
  divisions of substances into categories. The
  major division is between inorganic and organic.
• Among the inorganic compounds the three basic
  divisions are ionic compounds, molecular
  compounds and acids.

Sodium Nitrate
Potassium Oxide
Aluminum chloride
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Naming Positive Ions (Cations)

• Cations formed from metals atoms have the same
  name as the metal found on the periodic table.
  These are monatomic ions.
• Mg2?Magnesium ion and K ?Potassium ion
• If a metal can form different cations, the
  positive charge is indicated by a roman numeral
  in parentheses following the name of the metal.
  These are usually transition metals.
• Cu2 ?Copper (II) ion and Cu ?Copper (I) ion
• Cations formed from nonmetal atoms have name that
  end in ium. These are polyatomic ions.
• NH4- ?Ammonium ion and H3O ? Hydronium ion

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Naming Negative Ions (Anions)

• The names of monatomic anions are formed by
  replacing the ending of the name of the element
  with ide.
• O2- ?oxide ion and N3- ?nitride ion
• Polyatomic anions containing oxygen have names
  ending in -ate or -ite.
• NO3- ? nitrate ion and NO2- ?nitrite ion
• Anions derived by adding H to an oxyanion are
  name by adding hydrogen or dihydrogen as a prefix
  as appropriate.
• HCO3- ?hydrogen carbonate ion and H2PO4-
  ?dihydrogen phosphate ion

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Naming Ionic Compounds

• Names of ionic compounds consist of the cation
  name followed by the anion name.
• Do Sample Exercises 2.12 and 2.13 and Practice
  Exercises on pg 63

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Naming Acids

• You know a molecule is an acid because its cation
  is hydrogen.
• Acids containing anions whose names end in
  -ide are named by changing the -ide ending to -ic
  adding the prefix hydro- to this anion name and
  then following with the word acid.
• HCl?Hydrochloric Acid
• Acids containing anions whose name end in -ate or
  -ite are named by changing the -ate to -ic and
  -ite to -ous and then adding the word acid.
• HNO3?Nitric Acid HClO2?Chlorous Acid
• Do Sample Exercise 2.14 and Practice Exercise on
  pg 64-65.

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Naming Binary Molecular Compounds

• The name of the element farther to the left in
  the periodic table is usually written first.
  Except Oxygen is written last with all except
  Flourine.
• If both elements are in the same group in the
  periodic table, the one having the higher atomic
  number is named first.
• The name of the second element is given and -ide
  ending.
• Greek prefixes are used to indicate the number of
  atoms of each element. Although mono is never
  used with the first element.
• mono, di, tri, tetra, penta, hexa, hepta, octa,
  nona, deca
• N2O5 ? Dinitrogen pentoxide
• Sample Exercise 2.15 and Practice Exercise on pg
  65