chapter 8: covalent bonding

Chapter 8: Covalent Bonding Matter takes many forms in nature: In this

chapter, we are going to learn to distinguish the type of compound that we have already studied, the “ionic compound” (which contains oppositely-charged particles: metal cations and non-metal anions), from a different type of compound – a “molecular compound”. Additionally, we are going to focus on a type of molecular compound known as a binary molecular compound

II. Binary compounds: A “binary” compound contains atoms from

two different elements. A. NaCl”, “CaF2”, and “Al2O3” (3 ionic

compounds) are binary ionic compounds. “NH4Cl” is an ionic compound, but because it contains more than 2 elements, it is not a binary ionic compound.

B. “N2O5”, “SF6”, and H2O” (3 molecular compounds) are binary molecular compounds.

“C6H12O6” is a molecular compound, but because it contains more than 2 elements it is not a binary molecular compound.

Comparison/Contrast between an ionic compound and a molecular substance.

A. Molecular substances are made of molecules.

1. There is no “molecule” in an ionic compound.

B. A “molecule” contains a specific number of atoms, connected in a specific manner, to give a specific shape.

If even one atom is “missing” or “different”, the molecule would be an entirely different substance.

1. Not so with an ionic compound: In an ionic compound there is a specific ratio of atoms.. In salt (NaCl) for example, there is a ratio of 1 Na for every 1 Cl.

If a clump of salt lost 1 Na and 1 Cl, it would still be the same original substance: NaCl

C. The formula of a “molecule” should never be simplified. C2H8 is not the same substance as CH4.

1. The formula of an ionic compound should always be simplified. Ba2O2 is the same substance as BaO.

D. A molecule will not crack apart.1. Ionic compounds can crack apart if

hammered….. If the cations come close too close together and the anions come too close together, the structure cracks apart.

E. Molecular substances may have low melting points.

1. All ionic compounds have a very high melting point.

F. Molecular substances may, at room temperature, be found as solids, liquids, or gases.1. All ionic compounds are solids (at room

temperature).

G. Molecular substances contain atoms which are held together by covalent bonds.

1. Ionic substances are held together by ionic bonds.

IV. Covalent bonding: The type of bonding that occurs within a

molecular substance, in which atoms share their valence electrons in order to become more stable.

A. Occurs between atoms of nonmetallic elements.

B. Not all “molecules” or “molecular substances” are compounds!

In addition to the binary molecular compounds that we will study, there are 7 nonmetallic elements found in nature (in their elemental form) as pairs of atoms. These are the 7 “diatomic” elements:

N2, O2, F2, Cl2, Br2, I2, H2.

V. An important review:A. Metallic elements: Found to the left side

of the staircase boundary on the periodic table.

Non-metallic elements: Elements found to the right side of the

staircase boundary on the periodic table.2. Hydrogen is a nonmetallic element also.

VI. The octet rule: When a molecule is formed: “Nonmetal

atoms share electrons in covalent bonds in order to obtain a full octet of electrons.” An octet = 8 valence electrons.

A. Exception: A hydrogen atom will end up with a total of two electrons by sharing with 1 other atom.

B. There are a few other notable exceptions to the octet rule:

1. A few molecular compounds which contain an odd number of valence electrons are known to exist.

2. A few molecular compounds have either a boron or an aluminum atom with 6 valence electrons.

2. A few molecular compounds have a central atom with 10 or 12 valence electrons.

(1) One common example is “sulfur hexafluoride”.

In this compound, the central sulfur atom contains 6 x 2 = 12 valence electrons. Be sure to remember that this compound is an example in which the central atom does not follow the octet rule.

VII. Types of covalent bonds. A. Single covalent bond – 1 shared pair of

valence electrons: 2 dots, or a single dash, represent 2 electrons that are simultaneously being attracted by, or “shared” by, the nuclei of two neighboring atoms.

1. The formula in the center is a type of structural formula called a “Lewis dot structural formula”.

The formula on the right is the molecular formula.

H – H H H H2

B. Double covalent bond – two pairs of shared valence electrons: 4 dots

or 2 parallel dashes.

C C C C C2H4

H H

H H

H H

H H

C. Triple covalent bond– three pairs of shared valence electrons: 6 dots or 3 parallel dashes.

N N or N N N2

Notice the two “unshared pairs” of electrons (one pair is to the far left and one pair to the far righ)t of the nitrogen structure. You may never use a long dash to represent an unshared pair of electrons.

Unshared pairs of electrons don’t bond the atoms together….but, the repulsive forces of unshared pairs of electrons do dramatically influence the shape of a molecule!

D. Notice how an ion can react with a molecule to generate a polyatomic ion. In the example below, a hydrogen ion bonds to a molecule of ammonia(NH3) to make the ammonium ion (NH4)+:

H + + N H H N H

H

H H

H +][

VIII. Drawing a Lewis Dot Structure:

A. Certain elements are known as “central” atoms…. They will be found in the center of a structure. The first element given in a formula is usually the central atom (exception: hydrogen and the halogens).

1. Position the central atom in the center of your work space.

B. Hydrogen and the halogens are known as “peripheral” atoms. They will be found only connected to one other atom.

Position hydrogen and halogen atoms so that they “touch”, or “go around” only 1 other atom.

C. Add up all the valence electrons. Position the valence electrons as dots around the atom they belong to - the valence electrons may never leave the original atom.

Position the dots to form a “doorway” with 4 sides, in which the symbol of the element appears centered in the doorway.

Start with no more than 2 dots on each side of the 4 sided doorway.

D. If you can’t easily achieve a Lewis dot structure which has each atom (other than hydrogen) surrounded by 8 dots by doing what is described above, then you either need a double bond (2 pairs of shared electrons) or a triple bond (3 pairs of shared electrons).

For CO2, you will need two double bonds.

1. To make a double bond, move one “un-shared electron” simultaneously from each of two neighboring atoms, and place those 2 electrons in between the two neighboring atoms.

2. To make a triple bond, start with a double bonded pair of atoms, and simultaneously move one more unshared electron from each of the two atoms. Reposition those two electrons in between the atoms.

Important points regarding nonmetal atoms and their bonding charcteristics:

Atoms of the Following Nonmetallic Elements:

Have This Number of Electrons (dots) when in a Stable Structure:

Type of Bonds Permitted by these atoms:

Things to Remember about atoms of these elements; or to remember about a specified molecule.

Hydrogen 2 Single covalent Diatomic element; H2 is a “linear” diatomic molecule

Boron, Aluminum

8 or 6 Single covalent When only 6 electrons surround a boron or an aluminum atom, the molecule’s shape will be “trigonal planar” (a flat pancake).





Sulfur 8 EXCEPT

with “SF6” when there are 12

Single, double, and/or triple covalent

In sulfur hexafluoride sulfur does NOT follow the octet rule. This is one “exception” to the octet rule.

The Halogen family: F, Cl, Br, I, At

8 Single covalent All are diatomic elements; and,F2, Cl2, Br2, I2, At2 are all linear molecules, with only single bonds.





The Noble Gas family:

8….Except for helium (2). Noble gas atoms don’t form compounds.

Do NOT form compounds easily (no bonds).

Always found as single atoms in the gaseous state.

Nitrogen and oxygen

8 Single, double, and/or triple covalent

Diatomic elements; linear molecules. N2 has one triple bond, while O2 has 1 double bond.

All other nonmetal atoms

8 Single, double, and/or triple covalent

IX. Lewis Dot structural formulas for polyatomic ions:

A. Covalent bonds occur within a polyatomic ion (not between polyatomic ions).

B. When drawing polyatomic ions, place the first element in the center of the structure, and place the second element around the first element (placing 1 atom of the second element along each different side of the first element).

C. When the charge of a polyatomic ion is +, you need to subtract the indicated number of electrons from the total of the valence electrons in the molecule. So, for +1 ions: take away 1 electron from the molecular ion’s number of valence electrons.

D. When the charge of a polyatomic ion is –, you need to add the indicated number of electrons to the molecular ion’s number of valence electrons.

1. If the charge is 1-, then add 1 more electron to the molecule’s total number of valence electrons.

2. If the charge is 2-, then add 2 more electrons; if the charge is 3-, then add 3 more electrons.

E. Last, for a polyatomic ion: Draw a large bracket around the ion; and, place its charge at upper right.

# of N valence electrons: 5 5# of H valence electrons: 4 x 1 = 4 4Charge of ion = +1, therefore less 1 -1Therefore, total = 8Ammonium ion (NH4)+

H N H

H

H +][

Steps for Dot Structures:

Step 1: total # valence electrons.

Step 2. Position central atoms: carbon atoms form a straight line; assume only single bonding.

Step 3. Position other atoms; remember “special” molecules.

A. Peripheral atoms: Hydrogen and the halogens- connect to only 1 other atom, use only 1 single bond.

B. binary polyatomic ions: first element is central, second element is peripheral. assume all single bonds.

C. hydrocarbons” – molecular formula gives list of atoms

(from left to right) connecting to each central atom (usually carbon)

CH3CH2OH means “first carbon touches 3 H atoms, second carbon touches 2 H atoms, then there is an O touching an H. Assume all single bonds.

D. Memorize: CO2 C in the middle; use two double bonds

• Step 4: Make each atom stable.

Work from left to right: Assume all single bonds. Position unshared pairs to provide octets.

Exception: Hydrogen atoms = only 2 dots.

Step 5: Count the dots you’ve used.

Make sure the # you used = the # you were supposed to. Erase extras.

Step 6: Make corrections -

If your structure “needs” 2 extra dots, it really needs a double bond….

Erase 2 unshared dots, and share them (as part of a double bond).

If your structure “needs” 4 extra dots, it really needs a triple bond.....

Erase 4 unshared dots, and share them (as part of a triple bond).

Every atom should now be stable.

X. VSEPR Theory – Valence Shell Electron Pair Repulsion theory.

[Remember: Like charges repel!]A. A theory to predict the 3-dimensional

geometry, ie. the“shape” of a molecule

1. The theory is based on “electrostatic repulsion”: Molecules will adjust their shape to keep the negatively-charged pairs of valence electrons as far apart as possible from each other.

B. When NOT to use VSEPR theory: When there are only 2 atoms in a molecule. These molecule’s shapes are called linear – it doesn’t matter if there are single bonds, double bonds, triple bonds, or unshared electron pairs.

C. Using VSEPR theory: 1. Draw the Lewis dot structure for the

molecule. 2. Identify its central atom. 3. Identify the sets of valence electrons as one

of two possibilities:A. Those connecting two atoms.B. Those that do not connect two atoms.

These are called “unshared pairs”.

4. The unshared pairs found on a central atom strongly repel each other; and molecules that would otherwise be linear, will be forced into a bent (or angular) shape.

5. Unshared pairs also cause a molecule that would be shaped like a flat triangle (trigonal planar), to be forced into a not flat ( trigonal pyramidal) shape.

6.Count the number of connections separately from the number of unshared pairs.

1 single bond counts as 1 connection. 1 double bond counts as 1 connection. 1 triple bond counts as 1 connection. Each unshared set of 2 dots counts as 1 un-

shared pair.

D Predicting Shapes Using VSEPR Table

Read horizontally across the table.Connections To

the Central AtomUnshared Pairs of

Electrons Around Central Atom

Molecular ShapeAround Central

Atom

2 0 Linear

3 0 Trigonal Planar

4 0 Tetrahedral, 109.5o

2 1 or 2 Bent

3 1 Pyramidal

Shapes:

,E. Shapes:

Linear diatomic

Linear triatomic

Trigonal Planar

Bent

Pyramidal

Linear diatomic

Tetrahedral

Linear triatomic

Trigonal planar

bent

Trigonal pyramidaltetrahedral

“Molecular Polarity” – A term that is used to distinguish two types

of molecules…. Based on the presence or absence of a separation of charge.

Some molecules show characteristics indicating that they have oppositely-charged ends (a positive end and a negative end). This is called a separation of the charges (or “separation of charge”).

Other molecules show characteristics indicating that their structure doesn’t have a separation of charge, or their structure hides the presence of their oppositely-charged ends.

How to determine a molecule’s polarity.

The first part of determining a molecule’s polarity is to calculate each individual bond’s polarity.

Be careful with the vocabulary being used – An individual bond’s polarity is called the

“bond polarity”The polarity of the entire molecule is called

the “molecular polarity”

To calculate a bond polarity, first identify the “electronegativity value” of each of the 2 atoms in the bond you are working on.

The electronegativity value is number (from 0 to 4) which informs us of an atom’s ability to attract electrons when in a compound.

The electronegativity value is given on your periodic table, side 2, within each element’s square…..upper right corner of the square, in black print.

The closer an element’s electronegativity is to “4, the better that an atom of that element will attract electrons when that atom is found in a compound.

The closer an element’s electronegativity is to “0”, the less likely it is for that atom to be able to attract electrons when in a compound.

After identifying the two electronegativity values, subtract the two electronegativity values. Take the absolute value of the answer (make the answer positive).

After subtracting and taking the absolute value of the two electronegativity values, then think about your answer, and determine whether your answer indicates that there is, or that there is not, a situation in which one of the two atoms in the bond “overpowers” the other atom in terms of electron attracting ability. If an atom is able to overpower the other atom, it will “hog” the electrons, as opposed to sharing the electrons equally with the other atom. You will see (on the next page) that I’ve placed a “δ–”sign next to an oxygen atom, and a “δ+” sign next to 2 hydrogen atoms, in a sketch of a water molecule. By doing this, I am indicating that the oxygen atom is hogging the negatively-charged valence electrons belonging to the 2 hydrogen atoms; and, the 2 hydrogen atoms are both overpowered and “partially lose” their own valence electrons to the oxygen atom.

δ- indicates the “partial negative” atom; and, δ+ indicates the “partial positive” atom.

Finally, you are now able to conclude that a bond is either “nonpolar covalent” or “polar covalent”.

A bond is to be called “non-polar covalent” when the two atoms share the electrons more-or-less equally.

Labeling a bond as nonpolar covalent means that the difference in electronegativity values fell between 0 and 0.4

Labeling a bond as nonpolar covalent means that the electrons are distributed practically equally along the bond.

Labeling a bond as nonpolar covalent means that there is no separation of charge in that particular bond.

A bond is to be called “polar covalent” when the electrons are NOT distributed equally along the bond, rather, the electrons are found much or most of the time toward the atom that has the higher electronegsativity value.

This will occur when difference in electronegativity values is greater than 0.4, but less than 2.

We place a δ – next to an atom to identify it as the “partial negative” atom; and, we place a δ + next to an atom to identify it as the “partial positive” atom in a polar covalent bond.

A bond is designated as ionic when one atom has stripped the other of some of its valence electrons, yielding a cation and an anion. Then, the cation and anion are attracted, and more cations and anions are attracted and an ionic compound forms.

This should occur when the difference in electronegativity values is greater than 2.

We studied this type of compound previously. Ionic compounds have very different characteristics than compounds in which polar covalent and nonpolar covalent bonds exist.

The second part of determining the “molecular polarity” (a molecule’s polarity) is visualizing the effect of the individual bond polarities in conjunction with the shape of the molecule. Some common occurrences are listed below:

If the molecule’s shape is either tetrahedral, linear triatomic or trigonal planar and if all the bonds in the structure are identical, then the overall molecule’s polarity (the molecular polarity) is nonpolar

This is because, if every bond is nonpolar, there isn’t a substantial separation of charge to begin with.

And, if every bond is polar (as long as each bond has the same 2 atoms), the polarity of each bond will be cancelled due to the symmetry of the polar bonds around a central atom. When this symmetry exists, the δ – charge(s) cancel the δ + charge(s) in each bond.

If the molecule’s shape is either tetrahedral , trigonal planar, or linear, and the structure contains non-identical bonds in which 1 is polar, then the overall (molecular) polarity is polar (because 1 polar bond won’t have another bond to be cancelled with).

If the molecules’ shapes are pyramidal or bent, and one or more of the bonds is polar, the overall (molecular) polarity is “polar”. This is because in molecules having these shapes, the bond polarity is not situated symmetrically relative to the molecule’s center, so the polarity doesn’t cancel.

Polar bonds NonPolar molecule Polar bonds Polar molecule.

Polar bonds combine to cancel out; Polar bonds do NOT cancel out;

Producing Non-polar molecules: Producing Polar Molecules:

“Like dissolves like”: An expression stating that ionic compounds

and molecules whose molecular polarity is polar will be able to dissolve only in solvents containing polar molecules; and, conversely molecules whose molecular polarity is nonpolar will be able to dissolve only in solvents containing nonpolar molecules.

Some characteristics of a water molecule:

A H2O molecule has 2 identical polar bonds, a bent structure, and its molecular polarity is “polar”.

The partial negative (or slightly negative) region of a water molecule is the area closest to the oxygen atom.

The partial positive (or slightly positive) region of a water molecule lies within the area closest to the two hydrogens atoms.

Because of having the above characteristics, water behaves in an unusual manner:

When placed between a positively charged metal plate and a negatively charged metal plate, the water molecules all line up, with their positively-charged region attracted toward the negative plate and their negatively- charged region attracted toward the positive plate.

Inter- molecular Bonds (or forces of attraction)

are DIFFERENT from the covalent bonds you have been studying so far!

The covalent bonds you have been studying so far (single, double, triple bonds) are bonds within a molecule; these would be called______ intra-molecular bonds.

Inter-molecular bonds (forces) are the attractions between 2 molecules. They hold separate molecules together.

They are weaker than polar and non-polar covalent bonds (single/double/triple).

They are weaker than the ionic bonds which connect cations to anions.

Breaking an inter-molecular bond is a physical change; whereas, breaking an intra-molecular bond is a chemical change.

When you boil water (to generate water vapor), or melt ice (to generate liquid water), you are breaking inter-molecular bonds. When you treat a water molecule with electricity, you destroy the water molecule and generate from it oxygen gas and hydrogen gas….this is the breaking of intra-molecular bonds….this is a chemical change.

Types of intermolecular bonds:Hydrogen bond: Is the strongest

intermolecular force; and, it is perhaps the most important intermolecular bond, as it is necessary for life as we know it.

Hydrogen bonds take effect when you have a combination of a few certain atoms in a polar molecule.

The polar molecule must contain a hydrogen atom, and that hydrogen atom must be connected to one of the highly electronegative atoms listed below:

Fluorine, oxygen, OR nitrogen. Ex: H2O, HF, NH3

Hydrogen bonds determine the properties of water and biological molecules (such as proteins).

Cause water to predominate as a liquid (rather than as a gas) on earth.

Cause ice to expand upon freezing (rather than contract as the kinetic molecular theory would predict).

Holds the DNA double helix structure together.

Time-permitting, we will also learn about one of the weaker types of intermolecular forces of attraction by doing a lab in which we use evaporation rates to identify between hydrogen bonding and a weaker force, present between nonpolar molecules.

chapter 8: covalent bonding

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