Chapter 22: Chemical Bonding

1
Chapter 22 Chemical Bonding
2
Section 1 Electrons and Chemical Bonding

• Chemical bonding joining of atoms to form new
  substances
• Chemical bond interaction that holds 2 atoms
  together

3
Valence Electrons

• e- in outermost shell
• Determines an atoms chemical properties
• Used to form bonds
• Within a group, or family, atoms have the same
  of valence e-
• Atoms with fewer than 8 valence e- are more
  likely to form bonds than an atom with 8 e-

4
(No Transcript)
5
Types of Bonds

• Ionic, Covalent, Metallic
• Atoms bond by sharing, gaining or losing e- to
  have a filled outermost energy level.
• A full set 2 e- for a few of the elements
• A full set 8 e- for most elements
• Which ones? Why?

6
Section 2 Ionic Bonds

• Bonds form by gaining or losing e-, resulting in
  charged atoms called ions.
• Oppositely charged ions are attracted to one
  another
• Metal Nonmetal
• Positive and negative charges cancel each other
  out to form an overall neutral compound

7
Metal Atoms

• Have few valence electrons
• Usually lose these valence e- and form positive
  ions (cations)
• Some transition metal ions can have multiple
  charges. For example, iron can have a 2 or a 3
  charge.
• The charge is written as a superscript of the
  symbol
• Ex. K, Ca2, Al3

8
Nonmetal atoms

• Have almost full valence shells
• Tend to gain e- from other atoms and form
  negative ions (anions)
• The charge is written as a superscript of the
  formula
• Ex. P3-, S2-, Cl-

9
Polyatomic Ions

• Poly many
• Polyatomic many atoms
• A group of atoms that behave as a single ion with
  an overall positive or negative charge
• Treat the polyatomic ion as a single unit with a
  single charge.

10
(No Transcript)
11
Writing formulas for ionic compounds

• The number of positive charges and negative
  charges must balance in an ionic compound
• The formula represents this balance
• Subscripts are used to indicate the ratio of
  elements in the compound (no 1)

12
Writing formulas for ionic compounds

• Find oxidation number (charge) for both parts
• For elements in groups 1 and 2, use group .
  Boron family is 3, Carbon family 4, Nitrogen
  family is -3, Oxygen family is -2 and halogens
  are -1.
• For cations followed by a roman numeral, the
  roman numeral is the oxidation
• For polyatomic ions check the list. Do not change
  the subscripts within the polyatomic ion formula.

13

• 2. Write symbols. The positive ion first and
  negative ion second.
• 3. Put polyatomic ions in parentheses if more
  than one is needed.
• 4. Use subscripts to designate the number of each
  part for the total and charges to be
• a. Find the least common multiple of both
  charges
• b. determine the factor needed to get that
  charge and use that as the subscript.

14

• Sodium sulfide
• Potassium iodide
• Lithium oxide
• Barium fluoride
• Iron(III) oxide
• Copper(II) chloride

• g. Sodium acetate
• h. Zinc(II) carbonate
• i. Chromium(II) sulfate
• j. Cobalt(III) iodide

15
Answers

1. Na2S
2. KI
3. Li2O
4. BaF2
5. Fe2O3

• f. CuCl2
• g. NaC2H3O2
• h. ZnCO3
• i. CrSO4
• j. CoI3

16
Writing names for ionic compounds

• Write the names of the and - part of the
  formula
• The part is the name of the element or
  polyatomic ion
• Check the list of elements to see if it needs a
  roman numeral. If so, use the negative part of
  the formula to figure out the positive charge on
  the metal.
• To name the second/- part, if it is an element
  change the ending to -ide
• If it is a polyatomic ion, keep the name as is

17
Practice

• LiCl
• MgCl2
• BeO
• CaCl2
• HgS
• SnF2

g. (NH4)3PO4 h. ZnCO3 i. Sn(OH)2 j. Li2SO4 k.
KC2H3O2
18
Answers

• Lithium chloride
• Magnesium Chloride
• Beryllium oxide
• Calcium Chloride
• Mercury(II) sulfide
• Tin(II) fluoride

• g. Ammonium phosphate
• h. Zinc(II) carbonate
• i. Tin(II) hydroxide
• j. Lithium sulfate
• k. Potassium acetate

19

20
Section 3 Covalent and Metallic Bonds

• Covalent bonds are formed when atoms share one or
  more pairs of valence e-.
• Forms between 2 nonmetals.
• Covalent bonds result in the formation of
  molecules.
• Define molecule

21
Octet Rule

• Atoms combine in such a way so as to fill the
  valence shell (usually that means 8 e- but could
  be just 2 e-)

22
How many bonds?

• The number of e- that an atom needs to fulfill
  its valence is equal to the number of covalent
  bonds it can form.
• Ex. N can make 3 covalent bonds because it has 5
  e- in its valence shell
• Ex. H can make 1 covalent bond because it has 1
  e- its valence shell.

23
Diatomic Elements

• Certain elements exist as pairs in nature because
  that is how they are most stable.
• Di 2
• Just remember Professor BrINClHOF (Bromine,
  Iodine, Nitrogen, Chlorine, Hydrogen, Oxygen and
  Fluorine)
• You need to memorize the 7 diatomic elements!!

24
Practice! Draw a Bohr diagram and Lewis structure
for the following

• Water (H2O)
• Diatomic Fluorine (F2)
• Silicon tetrafluoride (SiF4)

25
Double Bonds

• When atoms share 2 pairs of e-, it is a double
  bond

26
Triple Bonds

• When atoms share 3 pairs of e-, it is a triple
  bond

27
Naming Covalent Compounds

• Many compounds have common names such as
  "methane", "ammonia" and "water
• Simple covalent compounds can be named using
  prefixes to indicate how many atoms of each
  element are in the formula.
• The ending of the last (most negative) element is
  changed to -ide.

28
Naming Covalent Compounds

• Prefixes
• Mono 1
• Di 2
• Tri 3
• Tetra 4
• Penta 5
• Hexa 6
• If there is just 1 of the first element no
  prefix is used

29
Practice

• CO2
• SiF4
• CO
• NBr3
• P2O5
• BCl3

30
Answers

• CO2 carbon dioxide
• SiF4 silicon tetrafluoride
• CO carbon monoxide
• NBr3 nitrogen tribromide
• P2O5 diphosphorus pentoxide
• BCl3 - boron trichloride

31
Practice

• Carbon tetrabromide
• Phosphorus triiodide
• Bromine
• silicon monoxide
• Silicon disulfide

32
Answers

• CBr4
• PI3
• Br2 (Bromine is a diatomic molecule!)
• SiO
• SiS2

33
Metallic Bonds

• A bond formed by the attraction between
  positively charged metal ions and surrounding e-
• Think of a metal as being made up of positive
  ions with electrons swimming around keeping the
  ions together

34
Properties of Metals

• Because e- can move freely about, metals have
  particular propeties.

35

• Conduct electricity well

36

• Metals can be reshaped (ductile, malleable)

37

• Metals can bend without breaking