Covalent Bonding and Nomenclature Notes I. Writing Formulas for Binary Molecular Compounds-those...

Covalent Bonding and Nomenclature Notes

I. Writing Formulas for Binary Molecular Compounds-those containing 2 nonmetals. Prefix naming system - know theses prefixes:

mono – one di – two tri – threetetra – four penta - five hexa –

sixhepta – seven octa - eight nona –

ninedeca – ten

Simply write what it says.

Ex: phosphorus pentachloride PCl5

dihydrogen monoxide H2O Practice: nitrogen tetrasulfide ______________ carbon dioxide ________________oxygen monofluoride _____________ sulfur hexachloride __________________trioxygen decanitride ______________ tetrafluorine monophosphide ___________hexafluorine nonasulfide ___________ heptabromine octanitride ____________

II. Writing Names for Binary Molecular Compounds1. The less electronegative element is given first. It is given a prefix only if it contributes more than one atom to a molecule of the compound. (All this means is that you will never start with mono-)2. The second element is named by combining a prefix indicating the number of atoms contributed by the element to the root of the name of the second element and then adding –ide to the end.

The o or a at the end of a prefix is usually dropped when the word following the prefix begins with another vowel. (monoxide or pentoxide)

Common RootsH: hydr C: carb N: nitr O: oxF: flor Si: silic P: phosph S: sulCl: chlorBr: brom I: iod

Practice: CCl4 _________________________

NF3 _______________________

PBr5_________________________

SF6_____________________________

SO3 _________________________

PCl5 _______________________

N2O_________________________

PF6_____________________________

Remember what happens when an ionic bond forms?

• One or more electrons from 1 atom are removed and attached to another atom, resulting in a cation and an ion which attract each other

Write this down!!Ionic Compounds - never exist as individual formula units, are solids

Molecular Compounds-can exist as an individual molecule, are usually liquids or gaseshttp://web.jjay.cuny.edu/~acarpi/NSC/6-react.htm

III. Formation of Covalent Bonds and Molecular Compounds

A. Covalent Bonds – a bond in which electrons are shared. Which compounds have covalent bonding?

1. Molecular (or covalent) compounds - these are two NON-METALS. These compounds always have covalent bonding

2. Polyatomic ions (PO-3, NO-1, CN-1). These ions are held together with covalent bonds.

B. Type of Covalent Bonds

l. Nonpolar covalent bond-a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in an evenly balanced charge. If the difference in electronegativity between two bonded atoms is less than 0.3 a nonpolar bond will exist.

Nonpolar Covalent Bond (equal sharing)

Oxygen AtomOxygen Atom Oxygen AtomOxygen Atom

Oxygen Molecule (O2)

2. Polar covalent bond-a bond is which the bonded atoms do NOT share the bonding electrons equally. A polar covalent bond is a bond in which the atoms have an unevenly balanced charge. If the difference in electronegativity between two bonded atoms is from 0.3 to 1.7, a polar bond will exist If the difference in EN is less than 0.3 then the bond is nonpolar covalent. The atom with the greater electronegativity will pull the electrons toward it, giving that atom a slightly negative charge. A partial negative charge is shown by - and the less electronegative atom will have a partial positive charge, designated +.

Polar Covalent Bonds: Unevenly matched, but willing to share.

The H-O bonds in water are polar covalent because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

Practice: Find the differences in electronegativity (EN page 151 or on calculator) in the following pairs of atoms. Designate which, if any, atom is partially negative and partially positive.

a. H and Cl

b. F and Br

c. S and I

d. O and H

Another Example: cesium-fluorine bonding

Cs EN = 0.7F EN = 4

4 - 0.7 = 3.3

Ionic Bond

D. The Octet Rule and Dot Structures -chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of valence electrons. Electron dot structures (also known as Lewis dot diagrams) show valence electrons as dots around the element’s symbol. Dot structures for molecules show atoms sharing dots (covalent bonds).

Ar

Covalent bonds are single, double, or triple

single bond-two atoms share one pair of electrons (1 sigma bond)

double bond-two atoms share two pair of electrons (1 sigma and 1 pi bond)

triple bond-two atoms share three pair of electrons (1 sigma and 2 pi bonds)

Rules for correctly illustrating the dot structure of a molecule:1. Add up the TOTAL number of valence electrons in the substancea) be sure to subtract 1 electron if it is a positively-charged ion (NH4

+1)b) be sure to add electrons for each negative charge on an ion (SO4

-2)2. Decide what is the central atom. The central atom is the one that is

least represented. (or the least electronegative)3. Hook the particles together using a short straight line (or 2 dots) to

indicate a covalent bond between atoms. Each of these "bonds" represents 2 shared electrons.

4. Subtract the number of electrons used in "hooking" the atoms together from the total valence electrons.

5. Use the "leftover" electrons, if any, to fill the octets of the peripheral atoms.

6. Place anymore "leftover" electrons on the central atom (in pairs).

http://chemsite.lsrhs.net/d_bonding/flashLewis.html

Practice: Draw the electron dot structures for the following:

Practice: Draw the electron dot structures for the following:

1. carbon tetrachloride (CCl4)

2. F2O

3. NF4+1

4. PCl3

5. CO2

6. N2

E. Exceptions to the Octet Rule

a. Some atoms have less then an octet.

Example: Hydrogen only needs 2 electrons surrounding it and boron only needs 6.

H2

BF3

b. Some atoms have more than an octet (One reason is because of bonding d orbitals as well as s and p orbitals.) Example: Sulfur can have up to 12 electrons surrounding it.

SF6

F. Resonance – a concept in which two or more Lewis structures for the same arrangement of atoms (resonance structures) are used to describe the bonding in a molecule or ion. To show resonance, a double-headed arrow is placed between a molecule’s resonance structures.

Example: Ozone

G. Coordinate Covalent Bond is formed when one atom contributes BOTH bonding electrons in a covalent bond.

Examples:

carbon monoxide

SO42-

HCN

Covalent Bonding and Nomenclature Notes Part 2

I. Metallic Bonds-a third type of bond. This is what holds pure metal atoms together.

What happens to form a metallic bond?

1. each metal donates its valence electron(s) to form an electron cloud

2. this leaves positive particles which are "cemented" together with the negative electron cloud, often called a “sea of electrons.”

Metallic Bonds: Mellow dogs with plenty of bones to go around.

A Sea of Electrons

Metals Form Alloys

Metals do not combine with metals. They form Alloys which is a solution of a metal in a metal.Examples are steel, brass, bronze and pewter.

II. Polarity - Polar and nonpolar molecules - if a molecule contains a

polar bond, is the molecule itself polar also? It depends!!

1. Polar molecules - a polar molecule is positive at one point and negative at another point. For example, HBr contains a polar bond. As a result the hydrogen side of the molecule is partially positive and the bromine side of the molecule is partially negative. It "acts like a magnet".

Water is a molecule with two polar bonds. A molecule of water is also polar because there is an area of positive charge on the hydrogen atoms and an area of partially negative charge on the oxygen. Its bent shape allows it to act somewhat like a magnet. The presence of 2 unshared pair of electrons is a main factor for it being a polar molecule.

2. Nonpolar molecules - Carbon tetrachloride has four C-Cl bonds. Each bond is a polar covalent bond.

The molecule itself is nonpolar because of its

1.) shape. It is perfectly symmetrical, and

2.) the partially-positive carbon in the center which is covered by the 4 partially negative chlorine atoms. It cannot “act like a magnet”.

Cl

Cl

Cl

3. Helpful hints and practice: A. Hints to help you decide if a molecule is POLAR: 1. Does it have at least one polar bond? If so, it's probably polar. 2. Does it have any unshared pairs of electrons around the central atom? If so,

it is probably polar. 3. Can the molecule act like a magnet? If so, it is probably polar. B. Practice: Which of the following molecules are polar and which ones are

nonpolar molecules? If the molecule is polar, tell why it is polar.

1.) SO2

2.) H2S 3.) CO2

4.) BF3

5.) CH4

6.) ClO2-1

7.) CH3Cl8.) PO4

-

9.) MgCl2

III. Hydrates: Some compounds trap water inside their crystal structure and are known as hydrates. You will not be able to predict which compounds will form hydrates. CuSO4 5H20 is an example of a hydrate. This says that one formula unit of cupric sulfate will trap 5 molecules of water inside its crystal.

Hydrates are named by naming the ionic compound by the regular rules and then adding (as a second word) a prefix indicating the number of water molecules. You will use the word “hydrate” to indicate water. The above compound would be called cupric sulfate pentahydrate.

To find the formula mass of a hydrate, simply find the mass of the ionic compound by itself and then ADD the mass of water molecule(s) to that mass.

Practice: What is the formula mass of barium chloride dihydrate?

1. What is the formula mass of aluminum sulfate octahydrate?