Molecular Bonds

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Molecular Bonds

• (Putting Elements Together)

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Molar Mass

• Each atom has an atomic mass
• Molar mass is the atomic mass of all the atoms in
  the molecule summed together
• For Example
• H2O 2 x Atomic Mass of H
• 1 x Atomic Mass of O

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Counting Atoms in a Molecule

• In the example, NH3, the subscript 3 only applies
  to the hydrogen.
• Therefore there is 1 N and 3 H in ammonia
• In the example, 3Ca3(PO4)2, the number of atoms
  changes due to the Coefficient in front of the
  molecule
• The 3 is multiplied to the Ca, P and O
• The subscript 2, multiplies the P and O
• 3Ca3(PO4)2

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3 Ca3 ( P O4 ) 2

• This means that
• there are 3 x 3 Ca,
• 3 x 2 P and 3 x (4 x 2) O

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Bonds. . .
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No, not that kind bonds between atoms to form
molecules

• It all depends upon the atoms valence (outer
  shell) electrons
• These are the e- in the last Energy Level (n
  1 through 7)
• Figure these out using the Periodic Chart and/or
  Lewis Dot Diagrams

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The Roman Numerals Tell You How Many Valence
Electrons for the Primary or Representative
Elements The Valence Electrons for the
Transition Elements Vary

• I II III IV V
  VI VII VIII

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• Group I is monovalent II is divalent III is
  trivalent IV is tetravalent V is back to being
  trivalent (since three e- openings) VI is
  divalent VII is monovalent and VIII has a
  complete octet, so these seldom react or bond

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• Bond Types (In General)
• Pure or Non-Polar Covalent
• ? difference 0 to 0.5 on the Pauling EN Scale
• The pair of e- shared are done so equally
• Two nonmetals bonded together
• Polar Covalent
• A shared pair of e-, but not equally
• ? difference 0.5 to 1.6
• Molecule has Partial and Charges
• Ionic Bonds
• ? difference 1.7 or higher to the maximum of
  4.0
• Metal bonded with a nonmetal
• Metallic Bonds are similar to Ionic Bonds

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Metallic Bonds

• Two or more metals mixed are called alloys
• Two major formats
• Interstitial and Substitutional
• These bonds permit the
• roaming of e- which creates
• a sea of dissociated e-
• Called the Electron Sea
• Model

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Ionic Bonds

• These are the bonds between a metal and a
  nonmetal
• The metal Ion is positively charged and called a
  cation
• The nonmetal Ion is negatively charged and called
  an anion
• The bonded molecule should be neutrally charged
  when finished

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Knowing where the metals and nonmetals are on the
table will make your life easier
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Lets take a moment to discuss polyatomic ions. .
.

• This is a molecule that acts as a cation or anion
• For example
• NH4 ammonium N3- azide
• ClO4- perchlorate CN- cyanide
• HCO3- bicarbonate OH- hydroxide
• CrzO7-2 chromate NO3- nitrate
• ClO3- chlorate C2H3O2- acetate
• Dont PANIC I gave a list to you!

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In an Ionic Bond one or more electrons are
lost or gained by the atoms involvedThis allows
the atoms to have a complete valence shell
following the octet rule
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• In an Ionic Compound balance the molecule using
  the criss-cross rule
• Mg 2 Cl-1
• Mg Cl2 The one is understood.
• This applies even if using a polyatomic ion

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• NH4 O-2
• (NH4)2O The parentheses are used to
  keep
• 
  the polyatomic together
• Pb4 CO3-2
• Pb2 (CO3)4 and this can be
  simplified by reducing the subscripts
  to
• Pb(CO3)2

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• Naming Ionic Compounds is really simple
• 1. Name the cation (metal) using its proper
  name if it is a polyatomic, do the same
• 2. Then, using the stem of the anion
  (nonmetal), simply add the suffix ide
• Zinc Chlorine Zinc Chloride
• Iron Oxygen Iron Oxide
• Lithium Cyanide Lithium Cyanide
• Ammonium Fluorine Ammonium Fluoride
• Cobalt Phosphorous Cobalt Phosphide

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• Transition Metals present an issue for balancing
  and naming molecules since they can have varying
  oxidation states
• For example
• Manganese can be a 2 or 3
• Iron can be a 2 or 3
• Lead can be a 2, or even a 4
• Copper is a 1 or 2
• Gold is usually a 1 or 3
• And Hydrogen is a 1 or a -1!

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Transition Metals

• To determine the correct Roman Numeral to place
  after the metal
• Roman Numeral - (Charge
  anion)(anions)
  
  ( cations)
• This is needed because, for example,
• iron chloride can be either FeCl2 or FeCl3
• or iron (II) chloride or iron (III) chloride

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Therefore Ionic Bonds are

• Metal Nonmetal
• ion - ion
• cation anion
• monatomic monatomic or
• (except NH4) polyatomic
• left of steps right of steps
• Reactions are Exothermic
• Form Crystal Lattice Structures

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Covalent Compounds

• These can be monatomic or polyatomic compounds
• It is a bond between two nonmetals
• They share a pair of electrons
• They can be subgrouped into polar or nonpolar
• If a binary compound (2 atoms) use the same
  naming rules as in Ionic Compounds

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• If it has more than two atoms need to use the
  prefixes
• Number Prefix Number Prefix
• 1 Mono 7 Hepta
• 2 Di 8 Octa
• 3 Tri 9 Nona
• 4 Tetra 10 Deca
• 5 Penta 11 Undeca
• 6 Hexa 12 Dodeca

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Naming Covalent Compounds

• Process
• Prefix Indicating full name of first
• nonmetal
• Prefix Indicating root name of second
  nonmetal suffix ide
• Watch for polyatomics and use their proper names

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For Example

• P4S10 becomes Tetraphosphorous Decasulfide
• P2O5 Becomes Diphosphorous Pentaoxide
• SF6 becomes Sulfur Hexafluoride
• SiBr4 becomes Silicon Tetrabromide

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Covalent Bonds can be Polar or Nonpolar

• A nonpolar has no discernable
• negative or positively charged sides
• (EN difference is 0)
• A polar covalent bond means one
• side is negative and the other positive

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• Electronegativity Percent Ionic Bond
• Difference Character Type
• 0.2 1 Non-polar
• 0.4 4 Covalent
• 0.5
• --------------------------------------------------
  ------------------------
• 0.6 9
• 0.8 15
• 1.0 22 Polar
• 1.2 30 Covalent
• 1.4 39
• --------------------------------------------------
  ------------------------
• 1.6 47 Ionic if metal/nonmetal
• 1.8 55 Polar Cov. if non/nonmetal
• 2.0 63
• --------------------------------------------------
  ------------------------
• 2.2 70
• 2.4 76 Pure Ionic
• 2.6 82

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• Some elements are able to form more than one
  oxyanion (polyatomic ions that contain oxygen),
  each containing a different number of oxygen
  atoms.
• For example, chlorine can combine with oxygen in
  four ways to form four different oxyanions
  ClO4-, ClO3-, ClO2-, and ClO- (Note that in a
  family of oxyanions, the charge remains the same
  only the number of oxygen atoms varies.)
• The most common of the chlorine oxyanions is
  chlorate, ClO3-. In fact, you will generally find
  that the most common of an elements oxyanions
  has a name with the form (root)ate.

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• The anion with one more oxygen atom than the
  (root)ate anion is named by putting per- at the
  beginning of the root and -ate at the end. For
  example, ClO4- is perchlorate.
• The anion with one fewer oxygen atom than the
  (root)ate anion is named with -ite on the end of
  the root. ClO2- is chlorite.
• The anion with two less oxygen atoms than the
  (root)ate anion is named by putting hypo- at the
  beginning of the root and -ite at the end. ClO-
  is hypochlorite.

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Oxyanion Example

• ClO- Hypochlorite
• ClO2- Chlorite
• ClO3- Chlorate
• ClO4- Perchlorate
•  

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• Some compounds have common names as well as their
  scientific names you should learn these and
  others!
• NO nitrogen monoxide nitric oxide
• H2O dihydrogen monoxide water
• NH3 nitrogen trihydride ammonia
• CH4 carbon tetrahydride methane
• C4H10 tetracarbon decahydride butane

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Some atoms are Diatomic KNOW THESE!

• H2 N2 O2 F2 Cl2 Br2 and I2
• and P is usually found as P4
• while Sulfur is found as S8
• Other elements will bond beyond the octet rule
  like PCl5, and the noble gas Xe bonds with F in
  XeF6, XeF2, XeF4, and XeO4 and this is due to a
  thing called hypervalence or expanded octet

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Molecular Geometry

• The 3-Dimensional Shapes of Molecules depend upon
  the valence e-s of the atoms involved
• Valence Bond Theory and VSEPR Model both use the
  same shapes
• Basically they focus on covalent bonds with the
  shared bonding pairs of electrons (BP)
• The assumption is made that the molecule will
  adopt a geometry to minimize the repulsion
  between e-s

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• The General Shapes

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Basic Geometry Bond Angles

• Linear 180o
• Trigonal Planar 120o
• Tetrahedral 109.5o
• Trigonal Bipyramidal 90o and
• 120o
• Octahedral 90o

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Molecular Orbital Theory

• MOT uses atomic orbitals (AO), e- ?s and e-
  density regions to examine bonds

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This is the end of Part I

• Next
• Van der Waals and London Dispersion Forces
• Polarity
• Intermolecular Forces
• Lewis Dot Diagrams with Covalent Bonds
• Determining Molecular Structure
• Resonance Structures