Unit 7 – Introduction to Bonding and Chemical

							Unit 7 – Introduction to
Bonding and Chemical
       Formulas
Addison Wesley – Ch 6, 15,
         and 16
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
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?
Ionic Compound

Made up of ions
Electrically neutral
Charges must equal each other
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 doesn’t conduct,
  molten does
Crystal is Brittle

Ionic solids shatter along a plane
This is because like charges align as soon
 as they are hit
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
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 Cu+1
Monatomic Anions

Can be negative 1,2, or 3
Never vary
Change element name to “ide” ending
Example: chloride
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
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
Shortcut

Place charge above each ion
Crisscross
Reduce if necessary (empirical formula is
 smallest ratio of ions)
Try some
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
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
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
Molecule
 Atoms joined together
  by covalent bonds
 Substance is
  molecular if it is made
  up of molecules
 Examples: CO2, H2O
  C6H12O6
Examples
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)
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
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
Try single bonds first!

Try O2
Double up if you come up with a shortage
 of one electron.
Triple bonds

Try N2
Triple up if you end up with a shortage of 2
 electrons.
Practice:

CO2     H2O    H2CO     HCN
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
Properties of Molecular Substances
                  •Let’s look at these
                  substances
                  •Can be solids, liquids,
                  or gases
                  •Much weaker bonds,
                  lower melting points
                  •Nonconductors
                  •Soft
Molecular Shape

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

CH4 – Tetrahedral
NH3 – Trigonal Pyramidal
H2O – Bent
HCl – Linear
Mixed
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
Polar Molecule

Molecules contains polar bonds
Arrangement is not symmetric
Try H2O
Try CCl4
Some Implications

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

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

Drug interactions
Genetic mistakes
Smells (Caraway and Spearmint)
Naming
Compunds are binary so they end in “ide”
Element name first, “ide” form second
Prefixes tell how many of each
N2O3 - Dinitrogen trioxide
Don’t 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
Try these!

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