ChemPhys, Chapter 20

1
Chapter 20 Chemical Bonds
2
Chemical Compounds

• The one cent piece is made of copper and is
  copper colored.
• The Statue of Liberty is covered with copper but
  is not copper colored. Instead it is green.
  Why?
• The copper on the outside of the statue of
  liberty has reacted with other chemical in the
  atmosphere to make the green compound we see.

3
Compounds Have New Properties

• When elements combine, the new substances has new
  and different properties.
• Sodium (Na) is a silvery metal that reacts
  violently with water.
• Chlorine (Cl) is a poisonous greenish-yellow gas.
• They combine to make NaCl, table salt.

4
Chemical Formulas

• A chemical formula tells what elements a compound
  contains and the exact number of atoms of each
  element in a unit of that compound.
• Water, H2O, has two atoms of hydrogen and one
  atom of oxygen.
• The subscript (2 in H2O) tells the number of that
  atom in the molecule.
• No subscript means 1 of that element.

5
Familiar Compounds
6
Atoms Form Compounds

• The electric forces between oppositely charged
  electrons and protons hold atoms and molecules
  together. (Positives and negatives attract.)
• Group 18, the noble gases, generally do not form
  bonds. Why?

7
Atomic Stability

• The outer energy level of each atom has a place
  for 8 electrons.
• An atom is chemically stable when its outer
  energy level is complete with 8 electrons.
• Noble gases have completed outer energy levels
  making them chemically stable and unreactive.

8
Electron Dot Diagrams

• We can used electron dot diagrams to show the
  outer electrons of an atom.
• For example

9
Energy Levels

• In the first row, hydrogen and helium have a
  place for only two electrons. So, helium is
  complete with two electrons.
• From the second row on down, each atom has room
  for 8 electrons. So, the other noble gases are
  complete with 8 electrons.
• When an atom fills the outer level it is most
  stable.

10
When Na and Cl React

• When sodium and chlorine react, the outer energy
  levels of each are complete
• Opposite charges hold the atoms together.

Electron Transfer
Na
Cl
Na
Cl-
Table Salt
11
When H and O React

• When hydrogen and oxygen react, the outer energy
  levels of each are complete
• The atoms share electrons.

H O H
H O H
12
Chemical Bond

• A chemical bond is the force that holds atoms
  together in a compound.
• This force is the result of the attraction
  between positive and negative charges in atoms
  and molecules.

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Chapter 20, Section 1 Review

• Describe how a chemical compound differs from its
  component elements.
• Explain what a chemical formula represents.
• Explain the electrical forces that exist between
  oppositely charged electrons and protons and how
  that is involved in forming compounds.
• Why does chemical bonding occur?

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Gain or Loss of Electrons

• When sodium and chlorine react an electron is
  transferred from Na to Cl
• An Ion is an atom which has become charged by
  gaining or losing electrons.

Electron Transfer
Na
Cl
Na
Cl-
Table Salt
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Gain or Loss of Electrons

• When potassium and iodine react an electron is
  transferred from K to I

Electron Transfer
K
I
K
l-
Potassium iodide
16
Ionic Bond

• An ionic bond is the force of attraction between
  the opposite charges of the ions in an ionic
  bond. Now, MgCl2

Mg
Cl
Cl
Becomes
Mg2
Cl-
Cl-
Magnesium Chloride
17
Zero Net Charge

• When an ionic compounds form, the number of
  positive charges equals the negative charges.
• In NaCl, one positive charge (Na) and one
  negative charge (Cl-)
• In MgCl2, two positive charges (Mg2) and two
  negative charges (2 x Cl-)

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Sharing Electrons

• The attraction that forms between atoms when they
  share electrons is known as a covalent bond.
• A molecule is the neutral particle that forms as
  a result of electron sharing.
• A single bond is made up of two shared electrons.

19
Water is Covalently Bonded

• The water molecule is held together by two single
  covalent bonds

H O H
Two Single Bonds
20
Multiple Bonds

• A double bond is formed when two atoms share two
  pairs of electrons.
• A triple bond is formed when two atoms share
  three pairs of electrons

N N ? N N
21
Unequal Sharing

• Sometimes one atom tugs harder on the pair of
  electrons in a covalent bond than the other.
• In HCl (stomach acid) the chlorine tugs harder on
  the shared pair of electrons than the hydrogen.
• Since the electrons spend more time on the
  chlorine, it has a partial negative charge.

22
Unequal Sharing

• In HCl the shared pair of electrons spend more
  time on the chlorine than the hydrogen
• The lowercase Greek letter d is read partial

d-
d
H Cl
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Polar or Nonpolar

• When two atoms equally share a pair of electrons,
  the covalent bond is nonpolar and there is no
  partial charge on either atom.
• When two atoms share a pair of electrons
  unequally, then the covalent bond is polar and
  the unequal sharing results in a partial charge
  on each atom.

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Chapter 20, Section 2 Review

• Describe ionic bonds and covalent bonds.
• Identify the particles produced by ionic bonding
  and by covalent bonding.
• What is the difference between a nonpolar
  covalent bond and a polar convalent bond?

25
Binary Ionic Compounds

• A binary compound is one that is composed of two
  elements
• Examples
• Sodium chloride NaCl
• Potassium iodide KI
• Magnesium chloride MgCl2

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Oxidation Number

• The number of electrons that an element gains or
  loses is called the oxidation number.

27
Oxidation Numbers of Elements
28
Special Ions
29
Compounds are Neutral

• When an ionic compounds form, the number of
  positive charges equals the negative charges.
• In CaF2, two positive charges (Ca2) and two
  negative charges (2 x F-)
• In Al2O3, six positive charges (2 x Al3) and six
  negative charges (3 x O2-). Six is the lowest
  common multiple of 2 and 3.

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Rules for Writing Formulas

• Write the symbol of the element or the polyatomic
  ion that has a positive oxidation number.
• Write the symbol of the element or the polyatomic
  ion that has a negative oxidation number.
• The charge (without the sign) becomes the
  subscript of the other ion. Reduce the
  subscripts to the ratio.

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Examples of Writing Formulas

• Lithium Nitrogen ? Li1 N3- ? Li3N
• Ca O ? Ca2 O2- ? Ca2O2 ? CaO
• Na Cl ? Na1 Cl1- ? NaCl
• Iron (III) O ? Fe3 O2- ? Fe2O3

32
Writing Names

• Write the name of the positive ion.
• Check to see if the positive ion is one of the
  special ones requiring a Roman numeral after it.
• Write the root name of the negative ion.
• Drop the ending of the negative element and add
  -ide.

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Names of Ionic Compounds
34
Polyatomic Ions

• Some positive and negative ions have more than
  one atom in them. These are called polyatomic
  ions.
• Some of the common polyatomic ions are in the
  next table and on charts on the wall of the
  classroom.

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Polyatomic Ions
36
Writing Names

• Write the name of the positive ion first followed
  by the name of the negative ion.
• K2SO4 becomes Potassium Sulfate
• Sr(OH)2 becomes Strontium Hydroxide
• FeCl3 becomes Iron (III) Chloride
• CuClO3 becomes Copper (I) Chlorate

37
Writing Formulas

• Follow the same rules as with binary compounds.
  Be sure to put polyatomic ions in parentheses
  because they go as a unit.
• Barium Chlorate Ba2 ClO31-
• becomes Ba (ClO3)2
• Ammonium Phosphate NH41 PO43-
• becomes (NH4)3PO4

38
Compounds With Added Water

• A hydrate is a compound that has water chemically
  attached to each molecule.
• The word hydrate comes from the word that means
  water.
• The formula is written followed by a dot and the
  water
• CoCl2.6H2O (pink) CoCl2 (blue)
• CaSO4.2H2O

39
Naming Binary Covalent Compounds

• Using what we have learned so far, N2O, NO, NO2,
  and N2O5 would all be nitrogen oxide.
• To solve this problem, prefixes are added
• N2O is named dinitrogen oxide
• NO is named nitrogen oxide
• NO2 is named nitrogen dioxide
• N2O5 is named dinitrogen pentoxide (a dropped
  to avoid the ao combination)

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Numerical Prefixes
Number Prefix Number Prefix
1 mono- 6 hexa- 2 di- 7 hepta-
3 tri- 8 octa- 4 tetra- 9 nona-
5 penta- 10 deca-
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Chapter 20, Section 3 Review

• How do you determine oxidation numbers?
• How do you write formulas and names for ionic
  compounds?
• How do you write formulas and names for covalent
  compounds?