Chapter 3 Atoms and Bonding

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					Chapter 3:
Atoms and
Bonding
Overview of Notes
Emphasis on Ionic and
Covalent Bonding
The Periodic Table of Elements
 In1869, Russian chemist Mendeleev
  proposed that elements had a
  “periodicity of properties.”
 Mendeleev originally tried to organize
  elements by atomic weight alone
  believing that elements would gradually
  and systematically change as weight
  increased.
Mendeleev Continued
 Mendeleev    actually found that his theory
  was only true to an extent.
 Physical and chemical properties would
  change gradually then suddenly at
  distinct steps.
 These steps came to be referred to as
  periods.
Elements in the same group
have the same number of
valence electrons.
Mendeleev and the Table
 The Rows on the Periodic Table are
  referred to as the periods.
 The columns are referred to as groups.
 Elements change gradually along the
  period and make more significant
  changes between groups.
 Elements in groups share many similar
  chemical and physical properties.
The Periodic Table of Elements
In Conclusion (Mendeleev)
 Mendeleev   ascertains that the “periodic”
  nature of chemical properties is related to
  electron configuration.
 The way that electrons are around the
  nucleus affect the properties.
Electron Shells
 Bohr’s theory tells us that electrons are not
  randomly located.
 Electrons are actually arranged in a very
  specific pattern called electron shells.
 Each electron shell has a limited capacity.
 The formula for determining the number
  of electrons in a shell is 2n2 where n refers
  to the number of the electron shell.
There is a specific way to draw
electron dot diagrams.
Electron Shells
 The  further the electron shell is away from
  the nucleus, the larger its capacity.
 Valence electrons refer to electrons found
  in the outermost shell regardless of the
  total number of electron shells.
 The row or period number that an
  element resides in on the table is equal to
  the number of shells that contain
  electrons.
Ionic Compounds
 Ionic compounds are composed of
  charged ions that are held together by
  electrostatic force.
 In an ionic bond, one element will give up
  an electron to another element to create
  the ideal number of valence electrons for
  the outer shell.
 Almost all ionic compounds are made up
  of a metal with a non-metal.
Characteristics of Ionic
Compounds
 Binary Compounds are made up of two
  elements composed of a metallic positive
  ion (cations) and non-metal negative ions
  (anions).
 Ternary Compounds contain three
  elements composed of a monoatomic
  ion and a polyatomic ion.
 For the purpose of this class we will mainly
  focus on Binary Compounds.
Important Note about Ionic
Compounds
 Ionic  Compounds do not show
  how the compound actually
  exists in nature. It shows the ratio
  by which the ions combine.
 For Example: Calcium Chloride
  (CaCl2) does not mean that
  there is one calcium atom and 2
  chlorine atoms.
 This does mean that there will be
  twice as many chlorine atoms as
  calcium.
Molecular Compounds
 Molecular   Compounds are created by
  covalent bonds.
 In this type of bond, electron pairs are
  shared between elements.
 Covalent bonds are usually not good
  conductors of electricity and will have a
  lower melting and boiling point.
Oxidation Number
 Whether   or not a compound is made up
  of ions, each atom in the compound will
  have an apparent charge.
 This charge is referred to as the oxidation
  number.
 This number represents the charge that an
  atom would have if electrons were
  transferred completely to the atom with
  the greater attraction.
Oxidation Numbers
 Oxidation  numbers are used to predict
  the rates by which atoms will combine
  when they form compounds.
 Oxidation numbers are written as
  superscripts in chemical formulas.
Oxidation Numbers
                                 Table 5-2a - Predicting Oxidation Numbers
 1. In free elements (that is, in uncombined state), each atom has an oxidation number of zero. Ex. In O2, the
 oxidation number of each oxygen atom is zero.
 2. For ions composed of only one atom, the oxidation number is equal to the charge on the ion. Ex. The
 oxidation number of Ca2+ is +2.
 3. All alkali metals (elements in column 1of the periodic table, with the exception of hydrogen) have an
 oxidation number of +1. Ex. The oxidation numbers of Li, K, and Na will always be +1.
 4. All alkaline earth metals (elements in column 2 of the periodic table) have an oxidation number of +2. Ex.
 The oxidation number of Ba is +2.
 5. The oxidation number of Aluminum (Al) is always +3.
 6. The oxidation number of oxygen in most compounds (such as H2O and CO2) is -2. In hydrogen peroxide
 (H2O2) and peroxide (O22-) oxygen shows a -1 oxidation number.
 7. The oxidation number of hydrogen is +1, except when in is bonded to a metal as a negative ion, in which
 case it is -1. Ex. H2O shows hydrogen as +1. NaH shows hydrogen as -1.
 8. When halogens (elements in column 17 on the periodic table) form negative ions, they will have an
 oxidation number of -1. Ex. NaCl and CaCl2 both show chlorine with a -1 oxidation number.
 9. In a neutral molecule, the sum of the oxidation numbers of all of the atoms must be zero. Ex. In H2O, each
 hydrogen is +1 and the oxygen is -2. So, (2 x +1) + (-2) = 0.
 10. In a polyatomic ion, the sum of oxidation numbers of all the elements in the ion must be equal to the net
 charge of the ion. Ex. In the polyatomic ion known as hydroxide (OH-), the oxygen is -2 and the hydrogen is
 +1. So, (-2) + (+1) = -1, the same as the charge on the hydroxide ion (OH-)

				
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posted:9/22/2012
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