Ch 9 Notes Covalent Bonding

9.1 p. 240 -241

Octet Rule: “atoms gain, lose, or share e-‘s to achieve a stable configuration of 8 valence e-‘s,  or an “octet”.”

There are exceptions to the octet rule- See p.256-257 for examples.

Covalent bond- occurs between two non-metals; means that atoms share e-‘s; the e’s are considered part of the complete outer energy level of both atoms involved.

Most compounds are covalent.

The word “molecule” indicates a covalent compound.

 

“bonding pair” – a shared pair of electrons;

“ lone pair”- unshared pair of electrons;

 

Why are some bonds covalent rather than Ionic?

What is electronegativity?

Attraction for electrons

Using electronegativity to predict type of bond:

·         Subtract electronegativities

·         If difference is <1.7, it is covalent

·         If difference is >1.7, it is ionic

 

Example:  SiO2

Electronegativity values: Si =1.90 and O = 3.44

Subtract: 3.44-1.90=1.54  , therefore it is covalent

Because atoms that have similar electronegativities (similar attractions to the electrons), they will share rather than  steal away the electrons of another.

 

Some elements only occur as diatomic molecules. Diatomic means double atom. Ex: O2;

These elements are diatomic:            

H2      N2      O2      F2      Cl2      Br2     I2       At2

 

Bond Length and Bond Strength (p. 246-247)

Look at fig. 9-7; Bond length is determined by the size of the atoms & how many electron pairs are shared.            Single>double>triple       (single is longest, triple is shortest)

Bond Strength-

Bond dissociation energy- energy required to break a bond.  (dissociation means to “come apart”);

Shorter bonds are stronger, thus their bond energy is greater. Example:  a C=C bond has a higher bond energy than a C-H bond.

Single<double<triple         (single is weakest, triple is strongest)

Lewis Structures:

Simple examples:  p. 243 Figures 9-2 & 9-3; p.244 Example problem 9-1 & Practice Problems # 1-5;

Things to remember:

·         H will always be on the end b/c it will only have 1 pr. of  shared electrons.  H never has lone pairs.

·         All other atoms will have 8 electrons (or four bonding pairs) to complete the octet.

·         You must account for the total number of valence electrons by showing bonding pairs as a line between 2 atoms, and lone pairs as a pair of dots.

·         The central atom is determined by: the atom with the lowest electron affinity (whichever element is to the left, or lowest on the periodic table.)

9.1 Assessment: p. 247 Q #6-12;

For a list of steps (procedure) for more complicated problems, see p. 252-255 (section 9.3);

assignment: p. 244 # 1-5 & p. 247 # 6-12;

RULES FOR DRAWING LEWIS STRUCTURES

1.    Count the number of valence electrons.

2.    Arrange atoms (often around a unique central atom).

3.    Draw single bonds between connected atoms.

4.    Fill octets around outer atoms with lone pairs (excluding H atoms).

5.    Put remaining electrons around central atom (even if more than octet).

6.    If less than octet around central atom, make multiple bonds.

Study these Example Problems:

                   9-3 on p. 253 & 9-4 on p. 254 & 9-5 on p. 255.

 

ASSIGNMENT: P. 255 # 30-34 & P. 258 #48

 

Section 9.3 Molecular Structures

Structural Formulas use symbols & bonds to show position of atoms.

There are exceptions to the octet rule- See p.256-257 for examples.

3 reasons:

1.      molecules have an odd # of valence electrons (NO2 has 17 valence electrons)

2.    Not enough valence electrons to form octet ( BH3 has only 6 valence electrons)

3.    Central atoms with more than 8 valence e’s (PCl5 has 10, an expanded octet)

 

Section 9.2 Naming Covalent Compounds

Some compounds have common names (“water”) but we need to know scientific names.

See the flow chart on p.251

·       Binary molecular compounds (CO2) end in “-ide” (as long as they are NOT an acid)

·       Prefixes are used to indicate the # of atoms of each type of element.

·       (Copy table 9-1 on p.248 for prefixes/meanings) *exception to the prefix rule-do not use the prefix “mono-“ on the first element.

Naming practice: Binary only-example problem on p.248-249 & practice problems # 13-17 on p. 249.

See table 9-2 p.249-common names/sci. names-good examples

 

Naming Acids:

How can I recognize an acid by its chemical formula? (p.250 & 597)

Most acids will have one or two hydrogen atoms at the beginning of their chemical formula. For Example, HNO3 is Nitric Acid, HCl is Hydrochloric Acid, and H2SO4 is Sulfuric Acid. There are some exceptions to this rule, such as Acetic Acid, the stuff in vinegar, whose chemical formula is CH3COOH.

1)          BINARY ACIDS- HCl          HBr

Use the prefix “hydro” + root of the 2nd element + “ic” ending.

Hydrochloric acid         Hydrobromic Acid

2)        OXYACIDS-  contains Hydrogen and an Oxyanion.  What’s an oxyanion? A polyatomic ion that contains oxygen. (p.224 has polyatomic ions table)

H2SO4        Sulfuric Acid

If the polyatomic ion ends in “ate”, then the acid ends is “ic” (SO42- is sulfate ion)

H2SO3        Sulfurous Acid

If the polyatomic ion ends in “ite”, then the acid ends is “ous” (SO32- is sulfite ion)

Bases:

Definition: a substance that gives off  hydroxide ions (or that accepts protons) in solution. 

Example: NaOH “sodium hydroxide”

(Review Video: Std. Deviants School of Chemistry: Chemical Bonds)

Lewis Structures

1.           Count valence electrons

2.         Decide what will be the central atom (pick the one with the lowest electronegativity)

3.         Arrange atoms with single bonds only

4.         Add lone pairs (electrons) to the OUTER atoms

5.         Count the # of e’s you used

All have been used?

·       Check to see if you have octet on all atoms

·       If you have octet, structure is complete

·       if molecule has a chg. Place brackets and charge outside

 

6.         If left overs?

 

7.          If you have left over electrons

   Place them on the central atom.

 

8.          Ck to see if you have octet

 

9.         If you don’t have octet…..

Create multiple bonds by shifting electrons to a shared position

 

·       Double Check: # of electrons, octets for all atoms, put brackets & charge around a charged molecule

 

Molecular Geometry

1.           Linear- 2 shared pairs, no lone prs; (CO2 & all diatomic molecules)

2.         Trigonal Planar- 3 shared pairs, no lone prs; (BH3)

3.         Tetrahedral-4 single cov. Bonds/ no lone prs. (CH4)

4.         Trigonal Pyramidal-3 single cov. Bonds/1 lone pr. (PH3)

5.         Bent-2 lone prs/2 single cov. Bonds (H20)

 

Practice: Draw lewis structure and predict the mol. Shape for:

 

NCl3                

OCl2

 

 

What’s on the Test?

·         Memorize list of elements that form diatomic molecules;

·          Naming covalent compounds;

1.     (flow chart p.251)

2.     Including binary(using prefixes)

3.     Including acids (binary acids, & oxyacids)

·         Drawing lewis structures of molecules and polyatomic ions

1.     Be able to determine which atom is the central atom, and which are terminal

2.     Be able to show multiple bonds when necessary

3.     make sure that you know where to place lone prs of e-

·         Relate length of bonds to strength of bonds (single, dbl, triple)

Draw resonance structures (example is:

http://www.mikeblaber.org/oldwine/chm1045/notes/Bonding/Resonan/IMG00007.GIF

·         Using electronegativity to predict type of bond

·         Be able to write the formula if given the name of the compound.

·         Tell the shape of the molecule (5 shapes) see p. 260

·         Study materials:  notes, book assignment, & worksheets;  see vocabulary list- p. 271