Unit 7 Introduction to Bonding and Chemical Formulas

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Unit 7 Introduction to Bonding and Chemical
Formulas

• Addison Wesley Ch 6, 15, and 16

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Practice Forming Ions

• Atoms lose or gain electrons to become stable
• Do dot diagram first to determine what will
  happen
• Do what is easiest to get 8 electrons
• Cation lithium
• Anion - oxygen

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Ionic Bond

• Electrons are everywhere static is a good
  example
• Positive ion is attracted to a negative ion in an
  ionic bond
• What kind of elements?

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Ionic Compound

• Made up of ions
• Electrically neutral
• Charges must equal each other

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Properties of Ionic Compounds

• Table salt is a good example
• High melting point ionic bonds are strong!
• Dissolve in water
• Solution conducts a current
• Solid doesnt conduct, molten does

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Crystal is Brittle

• Ionic solids shatter along a plane
• This is because like charges align as soon as
  they are hit

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Monatomic Ions

• Single atom loses or gains electrons to get a
  stable configuration (octet)
• Go over charges across table write on your
  periodic table
• Handout

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Monatomic Cations

• Positive 1, 2, or 3
• Transitions can vary
• Use element name
• Use a Roman numeral with any transition that
  varies
• Example Copper(I) is Cu1

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Monatomic Anions

• Can be negative 1,2, or 3
• Never vary
• Change element name to ide ending
• Example chloride

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Polyatomic Ions

• Two or more atoms that are bonded together and
  carry a single charge
• Names are on the handout
• Most are negative with one positive
• Usually end in ate or ite
• Example NO3- is nitrate

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Formulas for Binary Compounds

• Contain a monatomic cation (metal) and a
  monatomic anion (nonmetal)
• Metal is first
• Charges must add to 0
• Use subscripts to get the value to 0
• Why is sodium chloride NaCl?
• Try some

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Shortcut

• Place charge above each ion
• Crisscross
• Reduce if necessary (empirical formula is
  smallest ratio of ions)
• Try some

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Formulas with Polyatomic Ions

• Compounds are tertiary if they have 3 or more
  elements
• Treat polyatomic as a single unit
• Put in parentheses if you have to
• Practice

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Naming Ionic Compounds

• Use NaCl as a good example
• Metal first, then nonmetal
• Ends in ide if binary
• Use polyatomic name if tertiary
• Use Roman numeral if necessary
• How do you know?
• Suspect every transition metal
• Try some

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Covalent Bonding

• Look around
• Most of what you see is covalently bonded
• Definition formed by a pair of electrons that
  are shared between atoms.
• This would occur between nonmetals

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Molecule

• Atoms joined together by covalent bonds
• Substance is molecular if it is made up of
  molecules
• Examples CO2, H2O C6H12O6

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Examples
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Empirical Formula

• This is the simplest whole number ratio of atoms
  in a substance (reduced)
• Ionic formulas are always empirical
• Molecular formulas are sometimes empirical (H2O)
  
• and sometimes not (H2O2)

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Structural Formulas

• Different formulas may have the same formula, so
  structural formulas are often drawn
• Lewis structure is one example
• Based upon the Lewis dot diagram of elements
• Try bromine, magnesium

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Procedure

• Draw the Lewis dot structure for F2
• First, show the dot diagram for each atom
  involved in the molecule
• Share electrons in order for each atom to have an
  octet around it (8 electrons)
• Try HCl

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Try single bonds first!

• Try O2
• Double up if you come up with a shortage of one
  electron.

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Triple bonds

• Try N2
• Triple up if you end up with a shortage of 2
  electrons.
• Practice
• CO2 H2O H2CO HCN

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Exceptions

• Some elements are satisfied with fewer than 8
  electrons (6 or 4)
• Some structures can only be drawn with 7
  electrons
• These substances can be short-lived and reactive
  called free radicals
• Topic of college chemistry

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Properties of Molecular Substances

• Lets look at these substances
• Can be solids, liquids, or gases
• Much weaker bonds, lower melting points
• Nonconductors
• Soft

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Molecular Shape

• Try CH4
• Build with toothpicks and styrofoam balls
• Not 2-D!
• Electrons will repel each other

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VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory
• In a small molecule, the electrons are arranged
  as far apart as possible
• Explains most molecules
• 5 basic shapes

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Types

• CH4 Tetrahedral
• NH3 Trigonal Pyramidal
• H2O Bent
• HCl Linear
• Mixed

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Polarity

• Electrons are often not equally shared
• Polar bond Covalent bond between atoms that
  pull unequally on the electrons
• How do you know?
• Electronegativity difference (handout)
• Example H-O Polar
• Try a few

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Polar Molecule

• Molecules contains polar bonds
• Arrangement is not symmetric
• Try H2O
• Try CCl4

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Some Implications

• Shape gives molecules its properties (smell)
• Polarity determines solubility (Milk kaliedoscope)

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Isomers

• Molecules with the same formula but different
  arrangements
• Three types
• Structural (C and H)
• Functional (alcohol and ether)
• Stereoisomer (mirror image)

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Stereoisomers

• Drug interactions
• Genetic mistakes
• Smells (Caraway and Spearmint)

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Naming

• Compunds are binary so they end in ide
• Element name first, ide form second
• Prefixes tell how many of each
• N2O3 - Dinitrogen trioxide
• Dont use mono if one comes first
• Try CO2 and CO
• Write these on your periodic tables
• Mono,di,tri,tetra,penta,hexa,hepta,octa,
• nona,deca

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Try these!

• H2O
• Na2O
• CuCl
• N2O3
• AlN
• Check for a metal first!!