CHAPTER 8

COVALENT BONDS

• OCTET RULE = atoms lose, gain or share electrons in order to acquire a full set of 8 valence electrons– (form stable configuration)

• COVALENT BOND- 2 or more elements combine by sharing electrons; generally occurs with elements close together on the table

• MOLECULE- formed when 2 or more atoms bond covalently

– Examples: DNA; fats; cotton; polymers

– video

• ELECTRONEGATIVITY = tendency for an atom to attract electrons to itself when bonded– Covalent bonded atoms have

electronegativities that are close• The difference between the two values

is less than 1.70

Dog Bonds• http://ithacasciencezone.com/chemzone/lessons/

03bonding/dogbonds.htm

CHARACTERISTICS OF IONIC VS COVALENT BONDS

• CHARACTERISTIC IONIC COVALENT• Type of Particle formula unit molecule

• Bond formed by transferring e- by sharing e-

• Types of elements metal to nonmetal nonmetal to nonmetal

• Physical state solid solid, liq., gas

• Melting point high low

• Solubility in water high low (some will dissolve)

• Conductivity- good poor to non- Electrical conductive

FORMATION OF COVALENT BOND

• Covalent bond forms when the two atoms come close together and the distance is just right for the attraction between one atom’s protons and the other atom’s electrons

Each covalent bondrepresents one pairof shared electrons

•Molecules and Covalent Bonds

DIATOMIC MOLECULE• Occur in nature as a molecule of two

atoms because they are more stable than the individual atoms alone

• Examples– H2

– O2

– F2

– Cl2

– Br2

– I2

– N2

Except for Hydrogen, the diatomic molecules take the shape of the number 7 on the periodic table

NAMING MOLECULES

• NON-ORGANIC OR INORGANIC– Uses prefixes that stand for the

number of atoms present– Prefixes =

• 1 = mono 6= hexa• 2= di 7=

hepta• 3=tri 8=

octa• 4=tetra 9= nona• 5= penta 10= deca

Rules for inorganic mol.• 1st element– name as written on

periodic table, use a prefix if the # of atoms is greater than one

• 2nd element- always add a prefix indicating # of atoms present– Change ending to – ide– Example:

• SO2 – sulfur dioxide

• N2O3 – dinitrogen trioxide

INORGANIC PRAC. PROB.

• Write the name for the following compounds:– SiF4

– S5O6

– PCl3

• Write the formula for the following compounds:– Oxygen difluoride– Disulfur trioxide

ACID NAMING• ACID = a compound that

produces H+ ions in solution• 2 types of acids:

– 1. Binary acid-contains hydrogen and one other element (or pair without Oxygen)

– 2. Oxyacid- has a polyatomic ion containing oxygen with hydrogen attached

Naming Binary Acids• Use the prefix (hydro-)• Add the root of the second

element and attach a (-ic) ending

• Examples:– Hydrochloric acid HCl– Hydrofluoric acid HF

Naming oxyacids• If the ion ends in (-ate) change

to (-ic) and add the word “acid”

Example: HNO3

NO3 is nitrate so change the ending and

the compound is then called nitric acid

• If the ion ends in (-ite) change to (-ous) and add the word “acid”– Example: H2SO2 (SO2 is

sulfite) change to sulfurous acid

Naming organic molecules• HYDROCARBON = a compound

composed of hydrogen and carbon

• Organic names- consist of a prefix that indicates the # of carbon atoms followed by an ending that indicates the # of bonds between carbons

Organic Prefixes• 1 carbon atom = meth-• 2 carbon atoms = eth-• 3 carbon atoms = prop-• 4 carbon atoms = but-• 5 carbon atoms = pent-• 6 carbon atoms = hex-• 7 carbon atoms = hept-• 8 carbon atoms = oct-• 9 carbon atoms = non-• 10 carbon atoms = dec-

Types of organic bonds• ALKANE- all carbon-carbon

bonds are single• Have “-ane” ending• General formula= CnH2n+2 where n= number of carbons

• C2H6 = ethaneH- C - C -H

H

H

H

H

• ALKENE= have at least one carbon-to-carbon bond that is double

• Have “-ene” ending

• General formula is CnH2n

• Example:C3H6 propene

C = C

• ALKYNE – at least one carbon-to-carbon bond is a triple bond

• Have “-yne” ending• General formula CnH2n-2

• Example:C4H6 butyne

C C

Drawing hydrocarbons

• Use sticks to represent bonds– Single bond - share 1 pair of

e---Double bond = share 2 pairs of e-– Triple bond share 3 pairs of e-

• Use the symbol to represent the element

• C= carbon H = hydrogen F= Fluorine • Examples: H H H H

H-C-C-C-C-H C4 H10

H H H H butane

Remember:Carbon Needs 4

electrons to be shared

4 lines of attachment

HydrogenFluorine, chlorine, bromine, iodine

Needs 1 electron

1 line of attachment

Oxygen,sulfur

Nitrogen

Needs 2 electrons

3 electrons

2 line of attachment

3 lines

Practice Problems:

• Draw the following structures and write their formula:– Methane octene heptyne

– Hexyne propane butene

Carbon

LEWIS STRUCTURE

• Model that uses electron dot structures to show how electrons are arranged in molecules

• Pairs of dots or lines represent bonding pairs

• Hydrogen needs 2 dots, all other elements need 8 dots around them

• Example:H : H or H - HH:O: or H-O H H

::

SHAPES- MOLECULAR GEOMETRY

• For covalent bonds shape is very important, it determines some of its properties ( ex. Smells)

• Ex: NH3 has shape

N

HH

H

:

• VSEPR THEORY- (valence shell electron pair repulsion theory)– In a small molecule, the electrons

are arranged as far apart as possible

SHAPES• 1. LINEAR- straight line

– Bond angle = 180 degrees– Includes molecule with just 2

atoms (O = O, H-Cl) or more O=C=O

• 2. TRIGONAL PLANAR– Flat triangle– Central atom with 3 atoms

attached

B

Cl

Cl Cl

120o

• 3. TETRAHEDRAL– Four-surfaced shape– Triangular base

C

H

HH

H

109.5o

• 4. PYRAMIDAL– Has central atom bonded to three

others with an unshared pair of electrons

NH

HH

:

107o

• 5. BENT– Has one angle

O

H H

105o

: :

Shape simulationAnimated molecules

WEAK FORCES

• Intramolecular forces- forces within a molecule that holds atoms together in a covalent or ionic bond

– Van der Waals- weak forces involving the attraction of the electrons of one atom for the protons of another (covalent)

• Intermolecular forces- forces of attraction between molecules (like in solid or liquid state)– Dipole-dipole forces- forces of

attraction between two polar molecules

– Simulation

• Bond length- distance between two bonded atoms (also called bond axis)

• Bond angle- distance between axis

• Bond energy- energy required to break a chemical bond

OHH

Bond axis (length)

Bond Angle

• COVALENT RADIUS-radius of an atom when bonded to another

Cl Cl

Covalent radius

Polar and Nonpolar

MOLECULAR POLARITY• Not all covalent bonds are the

same

• The bonding shared pairs of electrons are pulled between the nuclei of the atoms sharing them

• Look at electronegativity differences to see what type of covalent bond

• Types– Nonpolar covalent – Polar covalent

NONPOLAR COVALENT• When the atoms are chemically

similar• The bonding e- are shared

EQUALLY (ex. H2, O2)

• OR bonds are symmetrically arranged (everything cancels)– Ex. O = C= O or CH4

• Electronegativity difference is < .4 generally

Non- Polar: Symmetrical(POLAR bonds that aren’t POLAR molecules)

• A polar molecule always contains polar bonds but some molecules with polar bonds are non-polar molecules (symmetrical)– CF4 = non polar

POLAR COVALENT• The 2 atoms are joined by

sharing electrons UNEQUALLY (E.N. difference is usually .4-1.7)

• The atom with stronger e- attraction (higher E.N. #) acquires a slightly negative charge

• The atoms with the lower E.N. # gets a slightly positive charge

• Set up poles

Points to side w/higher EN #

Electronegativity Type of bondDifference

0.0 - 0.4 covalent- nonpolar0.4 – 1.0 covalent-

moderately polar

1.0 – 1.7 covalent- very polar

> than 1.7 ionic

• Polar molecules develop partially charged ends, and ARE NOT symmetrical, like water

DIPOLE• Created by equal but opposite

charges that are separated by a short distance (use to show direction of (-) charge)

DISSOLVING RULE• RULE = “LIKE DISSOLVES LIKE”• Polar molecules dissolve in

polar water or polar solvent– Washable markers– Ionic dissolves in polar

• Nonpolar molecules dissolve in nonpolar solvents– Perm. Markers- alcohol– Dry cleaning

CONDUCTIVITY• Transmission of electric current

by ions• Ionic bonds- have electric

charge (e-) so they conduct• Covalent bonds- no charges, so

no charged particles- no conduction

• ELECTROLYTE- substance that dissolves in water to give a conducting solution (salts, acids, bases)

• In order to have conductivity:– 1. you need charged particles– 2. the charged particles (ions)

must be free to move– 3.the higher the concentration,

the more ions, the greater the conductivity

CHROMATOGRAPHY• Used to separate the

components of a solution to identify them (method of fractionation- separating parts from a whole)

• Separates by polarity

Parts of chromatography:• 1. Solvent- used to move the

mixture (needs to be same polarity as mobile phase)

• 2. mobile phase- consists of the mixture to be separated

• 3. stationary phase- thing the mobile phase is going to travel through- can be a solid or a liquid adhered to the surface of the solid– Has an attraction for polar

molecules

Chromatography process:

• The substances in the mobile phase will travel at different rates thru the stationary phase depending on their polarity– Slowest moving substance- has

greatest attraction for the stationary phase (most polar) found at the bottom of the paper

– Fastest moving substance- least attraction (least polar); top of paper; closer to the solvent front

TYPES OF CHROMATOGRAPHY• Paper

• Column • TLC (thin layer)• Gas chromatography• Ion exchange