1. Chapter 9
Chemical Bonds

2. Core Concept
Electron structure will explain how and why atoms join together in certain numbers.

3. Valence Electrons and Ions
Outer electrons determine the chemical properties of an atom
Octet rule
Atoms attempt to acquire an outer shell of eight electrons
Electrons can be gained/lost/shared in the process
Example: sodium (Na)

4. Chemical Bonds Covalent Three types: Ionic Metallic bonds
Attractive forces holding atoms together in compounds
Three types:
Ionic
Electrons transferred between atoms
Electrostatic force = binding force
Covalent
Octets achieved through sharing electrons
Typically between nonmetallic elements
Metallic bonds
Outer electrons move freely throughout metal
“Electron sea” within rigid lattice of metal atoms
Conduct heat and electricity well

5. Ionic Bonds Chemical bond of electrostatic attraction
Forms between metals and nonmetals
Form crystalline solids with orderly geometric structure
Example: NaCl
Na loses; Cl gains
No single NaCl molecule, per se

6. Ionic Compound Formulas
Ionic compounds have an overall neutral charge
If charges aren’t complementary on ions, multiple ions are needed
Multiple ions are represented by subscripts

7. Ionic Compound Formulas
Two rules
Write symbol for positive ion first followed by negative ion symbol
Assign subscripts to assure compound is electrically neutral
Example: Calcium chloride

8. Ionic Compound Names Name of metal (positive) ion first
Name is unchanged
Stem name of second element next ending in “-ide” (for two elements)
Many elements have variable charges
Historical suffix usage
“-ic” for higher of two;
“-ous” for lower
Modern approach
English name of metal followed by Roman numeral indicating charge

9. Lewis Dot Structures for Ionic Compounds
For anions fill the octet by adding the required number of electrons
Example O atom becoming O2-.
For cations empty the octet by losing all electrons from the valence shell.
Example Mg atom becoming Mg2+.

10. Lewis Dot Structures for Ionic Compounds
Place squared brackets around each ion and place the charge as a superscript following the brackets.
To represent the number of each ion that is needed place a coefficient in front of the squared brackets for each ion to represent the required quantity.
Ex: Calcium chloride
[Ca]+2 + 2[Cl]-1

11. Covalent Bonds Chemical bonds formed by sharing pairs of electrons
Formed between 2 nonmetals
Electrons shared to form octets, ideally
Overlap of shared electron clouds between nuclei yields net attraction
Atoms within covalent compounds are electrically neutral, or nearly so

12. Covalent Compounds and Formulas
Covalent compound - held together by covalent bonds
Electrons shared in covalent bonds
Electron dot representation
Bonding pairs shared
Lone (non-bonding) pairs not shared

13. Multiple Bonds Sharing of more than one electron pair Examples
Ethylene - double bond
Acetylene - triple bond

14. Composition of Compounds
Millions of different combinations of over 90 elements
Common names
Often related to historical usage (baking soda, washing soda,…)
Difficult to relate to actual molecular composition
Modern approach - systematic sets of rules
Different for ionic and covalent compounds
One common rule - “-ide” means compound contains only two different elements

16. Covalent Compound Names
Molecular - composed of two or more nonmetals
Same elements can combine to form a number of different compounds
Two rules
First element in formula named first with number indicated by Greek prefix
Stem name of second element next; Greek prefix for number; ending in “-ide” (for two elements)

17. Lewis Structures
Lewis structures are representations of molecules showing all valence electrons, bonding and nonbonding.

18. Electronegativities

19. Writing Lewis Structures
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
If it is an anion, add one electron for each negative charge.
If it is a cation, subtract one electron for each positive charge.
PCl3
(7) = 26

20. Writing Lewis Structures
The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.
Keep track of the electrons:
26  6 = 20

21. Writing Lewis Structures
Fill the octets of the outer atoms.
Keep track of the electrons:
26  6 = 20  18 = 2

22. Writing Lewis Structures
Fill the octet of the central atom.
Keep track of the electrons:
26  6 = 20  18 = 2  2 = 0

23. Writing Lewis Structures
If you run out of electrons before the central atom has an octet…
…form multiple bonds until it does.

24. Bond Polarity Result of unequal sharing of electrons Electronegativity
Measure of an atom’s ability to attract electrons
Differences:
1.7 or greater - ionic
polar covalent
Less than covalent