COVALENT COMPOUNDS Back to Lewis Dot Structures and Valence Electrons!

Review of Lewis Dot Structures

¨  Electron Dot Structures contain: ¤ Element’s Symbol: representing the atom’s nucleus and

inner electrons ¤ Dots: representing all of the valence electrons

n 8 maximum

¨  Examples:

N O Cl

Covalent Bonding

¨  A joining of two atoms through the sharing of valence electrons. ¤ Most common with non-metals.

¨  Octet Rule will apply.

¨  They share electrons to become more stable.

Sharing Electrons

¨  Atoms want to become stable and look like the noble gas (in terms of their electron configuration)

¨  Because non-metals are close to having 8 electrons,

they will not give up electrons! ¨  In order to look like the noble gas, they must share

their valence electrons with each other. ¤ Both atoms get to count the shared electron(s) toward

the noble gas configuration

Example: Fluorine

¨  One atom of Fluorine has 7 valence electrons. ¤  It needs 1 more electron to become stable.

¨  A second fluorine atom also has 7 valence electrons. ¤ Again, it needs only 1 more electron to become stable.

¨  By sharing electrons both atoms will end with full orbitals.

Types of Covalent Bonds

¨  Single Covalent Bond: ¤ Two atoms are held together by sharing one pair of

electrons ¤ Usually represented with a single line

¨  These are the longest bonds and also the weakest ¨  Example:

¤ Cl-Cl

Types of Covalent Bonds

¨  Double Covalent Bond: ¤ Two atoms held together by sharing two pairs of

electrons. ¤ Usually represented with a double line

¨  Example: ¤ O=O

Types of Covalent Bonds

¨  Triple Covalent Bond: ¤ Two atoms held together by sharing three pairs of

electrons ¤ Usually represented with a triple line

¨  These are the shortest and strongest bond

Electrons Involved in Bonding

¨  Depending on the number of valence electrons available: ¤ Most atoms can have up to 4 covalent bonds.

¨  This is based on unshared electron pairs or “Lone Pairs” of electrons ¤ These are valence electrons that are NOT involved in a

covalent bond.

Example: Oxygen

Lone Pairs of Electrons

Rules for Drawing Lewis Dot Structures

¨  Rule 1: ¤ Add together the number of valence electrons for EACH

atom in the molecule

¨  Example: CF4

¤ Carbon has 4 valence electrons ¤ Each Fluorine has 7 valence electrons ¤ Therefore, the total number of valence electrons =

4 + 4(7) = 32

Rules for Drawing Lewis Dot Structures

¨  Rule 2: ¤ Write out the elements of the molecule so that the least

electronegative element is in the center surrounded by the other elements n TIP: Fluorine is the most electronegative. The closer the

elements are to Fluorine, the higher the electronegativity!

¨  Example: CH4

C F F

F F

Rules for Drawing Lewis Dot Structures

¨  Rule 3: ¤ Place a covalent bond between the central atom and

the outside atoms. n Covalent bonds are represented by a line. n Remember: each covalent bond contains two electrons.

¨  Example: CF4

C F

F

F

F

Rules for Drawing Lewis Dot Structures

¨  Rule 4: ¤ Add electrons to the outer atoms as “lone pairs” to

satisfy the Octet Rule.

¨  Example: CF4

¤ We started out with 32 electrons ¤ Each line you drew connecting Carbon to Fluorine was 2

electrons. So in total you used 8 electrons. ¤ There are now 24 valence electrons remaining.

C F

F

F

F

Practice Problem #1

¨  Write the Lewis Dot Structure for TeH2 ¤ What is the total number of valence electrons in this

molecule? n Te: 6 valence electrons n H: 1 valence electron n Total: 6 + 2(1) = 8

¤ Which atom is at the center of the molecule? n The least electronegative is Te

Practice Problem #1

¨  Draw Te in the middle and H on opposite sides of Te.

¨  How many valence electrons do you have left?

¤  You have 4 valence electrons left. ¤  You started with 8 electrons and use four for the bonds

n  Two bonds each with two electrons

Te HH

Practice Problem #1

¨  Will the remaining 4 electrons go on Te or H? ¤ Te!

¨  You have now written the Lewis Dot structure for TeH2

Te HH

Practice Problem #2

¨  Write the Lewis Dot Structure for CO2 ¤ What is the total number of valence electrons in this

molecule? n Carbon: 4 valence electrons n Oxygen: 6 valence electrons n Total: 4 + 2(6) = 16 valence electrons

¤ Which atom is at the center of the molecule? n The least electronegative element is Carbon

Practice Problem #2

¨  Draw Carbon in the middle and O on opposite sides

¨  How many valence electrons do you have left? ¤ You have 12 valence electrons left. ¤ You started with 16 electrons and use four for the

bonds n Two bonds each with two electrons

C OO

Practice Problem #2

¨  Will the remaining 16 electrons go on C or O? ¤  They will be split amongst both, C and O, in order to give

each atom a complete octet!

¨  I have now used up my remaining 12 electrons. ¨  Do they form a complete octet for each atom?

¤ NO.

C OO

Practice Problem #2

¨  Each atom does not have a complete octet and we need to fix this problem.

¨  To fix the problem, you need pair up any electrons

that are not paired

C OO

Practice Problem #2

¨  To pair them together, just draw another line between the C and O. ¤ This will form a double bond!

¨  Now, does each atom contain a full octet? ¤ Yes!

C OO

COVALENT BONDS Naming Compounds, Writing Formulas

Properties of Covalent Bonds

Bond Formation Electrons are shared between atoms

Type of Structure Molecules

Physical State Can be liquids, gases, or brittle solids

Melting Point Low

Soluble in Water? Most are not, BUT there are a few that are

Electrical Conductivity No

Other Some have odors

Strength of Covalent Bonds

¨  Strength depends on the distance between 2 nuclei (or bond length)

¨  As length increases, strength decreases!

Molecular Formulas

¨  Two or more atoms joined together by covalent bonds are called a Molecule

¨  Molecular Formula: ¤ A system for telling you the type and number of each

element in a molecule ¤ H2O; CO2; C6H12O6

¨  Remember: A formula unit for an ionic compound has the same style as a molecular formula. However, the unit tells you the ratio of cations to anions in an ionic compound

Diatomic Molecules

¨  Molecules are individual particles ¤  In comparison, Ionic Compounds exist as large clusters

of ions arranged in a crystal lattice

¨  Diatomic Elements: ¤ Seven elements that, when in their pure form, will bond

in pairs. n They are more stable in pairs than alone.

Naming Covalent Compounds

¨  The following are the naming rules for a compound that is composed of all non-metals

¨  We know the number of atoms from each element by the prefixes used in the name of the compound

¨  If there is more than one atom of an element used, then a prefix is needed.

Covalent Pre-fixes

Number of Atoms Prefix

1 Mono (use only w/ Oxygen)

2 di-

3 tri-

4 tetra-

5 penta-

6 hexa-

7 hepta-

8 octa-

9 nona-

10 deca-

Naming Covalent Compounds

¨  The second element ends in “ide” as it did for naming ionic compounds ¤  If the vowel combo is o-o or a-o, omit the first vowel (ex:

monoxide)

¨  IF there is only one atom of the first element, then it DOES NOT get a prefix

¨  Example: ¤ NCl3

n  nitrogen trichloride

Practice

¨  P2O5 ¤ Diphosphorous pentoxide

¨  CO

¤ Carbon monoxide

¨  CF4 ¤ Carbon tetrafluoride

¨  CCl4

¤ Carbon tetrachloride

¨  N2O ¤ Dinitrogen monoxide

¨  SF6

¤ Sulfur hexafluoride

Writing Formula Names

¨  Use the prefixes in the names of each element to determine how many of element

¨  Write the element’s symbol and use subscripts to indicate

the number of each element

¨  Remember: ¤  If the first element has only one, it WILL NOT use mono-

¨  Example: ¤ Carbon Tetrachloride

n  CCl4

General Names

¨  Some compounds get more common names because they are used often

¨  Example: ¤ H2O (water) ¤ CH4 (Methane) ¤ NH3 (Ammonia) ¤ C6H6 (Benzene)

Practice

¨  Arsenic trichloride ¤ AsCl3

¨  Dinitrogen pentoxide ¤ N2O5

¨  Tetraphosphorus decoxide ¤  P4O10