Molecular Bonding

Unit 5

Covalent Bonds Sharing pairs of electrons Covalent bonds are the inter-atomic

attraction resulting from the sharing of electrons between atoms.

They result in ‘localized overlaps’ of orbitals of different atoms.

They also are the result of the attraction of electrons for the nucleus of other atoms.

Typical of molecular substances.

Example of a Covalent Bond

Covalent Bonds Cont. Atoms bond together to form

molecules– molecules are electrically neutral groups

of atoms joined together by covalent bonds

– strong attraction Molecules attracted to each other

weakly form molecular compounds

Properties of Molecular Compounds

Strong covalent bonds hold the atoms together within a molecule.

The intermolecular forces that hold one molecule to another are much weaker.

Properties vary depending on the strength of the intermolecular forces.

Lewis structures

A Lewis structure is a representation of the valence e- in an atom, ion or molecule.

• Element symbol represents nucleus and core e-.

• Dots represent valence e-.

Electron pairs

• In covalent compounds electrons are shared between atoms creating electron pairs.

• Bonding pairs: e- that are shared between 2 atoms.

•Lone or unshared pairs: e- that are NOT involved in bonding.

Covalent bonds

Hydrogen follows the duet rule: sharing 2 electrons.

Non-metals Carbon through Fluorine follow the octet rule: sharing 8 electrons.

Writing Lewis Structures of Molecules

1. Determine the central atom (atom in the middle)- usually is the “single” atom- least electronegative element- H never in the middle; C always in the middle

2. Count the total number of valence e- (group #)- add ion charge for “-”- subtract ion charge for “+”

3. Divide the total number of electrons by 2- sharing involves 2 electrons

4. Attach the atoms together with one pair of electrons

5. All remaining e- = LONE PAIRS! - lone pairs are NOT involved in bonding

Writing Lewis Structures Cont.

6. Place lone pairs around non-central atoms to fulfill the “octet rule” - some elements may violate this octet rule – (H=2, Be=4, B=6)

7. If more e- are still needed, create double or triple bonds around the central atom.

- single = 1 pair of shared electrons (2 e-) - double = 2 pair of shared electrons (4 e-) - triple = 3 pair of shared electrons (6 e-)

Lewis structure water

Practice

Give the Lewis structure for: HCl NH3

C2H6

CO2

NAS When drawing Lewis structures, remember: terminal atoms, atoms

that can only make one bond, must be on the outside. N – A = S #e- needed for – #e- available = #e- shared octet or duet (valence e-) (bonds)

S = the number of pairs of shared electrons S = the number of bonds (a dash may be used to 2 represent a pair of shared electrons)

Resonance When there is more than one Lewis

structure for a molecule that differ only in the position of the electrons they are called resonance structures– Lone Pairs and Multiple Bonds in different

positions Resonance only occurs when there are

double bonds and when the same atoms are attached to the central atom

The actual molecule is a combination of all the resonance forms.

•••• •• ••••••••

•• ••O S O O S O•••••• ••••

••••

••••

Exceptions to the Octet Rule

• H, Be, B (stable with 2, 4, and 6 e-, respectively)

Some molecules cannot be drawn with the Lewis structure rules, due to odd # of e-. i.e. NO & NO2 (There is no way for N to get an octet)

• Some molecules are stable when the center atom has more than an octet. i.e. SF6, PCl5

sulfur has 12 electrons P has 10 electrons

Coordinate Covalent Bond A covalent bond in which one atom

contributes both bonding electrons.

Polar Bonds: Electronegativity

Measure of the ability of an atom to attract shared electrons– Larger electronegativity means atom attracts more

strongly– Values 0.7 to 4.0

Increases across period (left to right) on Periodic Table

Decreases down group (top to bottom) on Periodic Table

Larger difference in electronegativities means more polar bond– negative end toward more electronegative

atom Covalent bonding between unlike atoms

results in unequal sharing of the electrons– One end of the bond has larger electron

density (more electronegative) than the other– Polar covalent – unequal sharing– Nonpolar covalent – equal sharing

Bond Polarity The result is bond polarity

– The end with the larger electron density gets a partial negative charge

– The end that is electron deficient gets a partial positive charge

H F••+d -d

Predicting Molecular Geometry

VSEPR Theory– Valence Shell Electron Pair Repulsion

The shape around the central atom(s) can be predicted by assuming that the areas of electrons on the central atom will repel each other

Each Bond counts as 1 area of electrons– single, double or triple all count as 1 area

Each Lone Pair counts a 1 area of electrons– Even though lone pairs are not attached to

other atoms, they do “occupy space” around the central atom

– Lone pairs generally “push harder” than bonding electrons, affecting the bond angle

Shapes

Straight Line– molecule made up of only 2 atoms

Shapes- Linear

– 2 atoms on opposite sides of central atom, no lone pairs around CA

– 180° bond angles Trigonal Planar

– 3 atoms form a triangle around the central atom, no lone pairs around CA

– Planar– 120° bond angles

180°

120°

Tetrahedral– 4 surrounding atoms form a tetrahedron

around the central atom, no lone pairs around the CA

– 109.5° bond angles

109.5°

Shapes

Trigonal Pyramidal– 3 bonding areas and 1 lone pair around

the CA– Bond angle = 1070

V-shaped or Bent– 2 bonding areas and 2 lone pairs around

the CA– bond angle = 104.50

Dipole Moment Bond polarity results in an unequal electron

distribution, resulting in areas of partial positive and partial negative charge

Any molecule that has a center of positive charge and a center of negative charge in different points is said to have a dipole moment

If a molecule has more than one polar covalent bond, the areas of partial negative and positive charge for each bond will partially add to or cancel out each other

The end result will be a molecule with one center of positive charge and one center of negative charge

The dipole moment affects the attractive forces between molecules and therefore the physical properties of the substance

Charge distribution in the water molecule

Polarity of Molecules

Molecule will be NONPOLAR if:– the bonds are nonpolar (Br-Br, F-F)– there are no lone pairs around the

central atom and all the atoms attached to the central atom are the same

Molecule will be POLAR if:– the central atom has lone pairs– there are no lone pairs around the

central atom and all the atoms attached to the central atom are NOT the same

Intermolecular forces From weak to strong:

- Dispersion- Dipole-dipole- Hydrogen bonding- Ion-dipole (attraction between ions and dipole molecules)- IonicDispersion, dipole-dipole, and hydrogen bonding are Van der Waals forces

Intermolecular Forces Hydrogen Bonding – extreme dipole

bonding involving hydrogen and a very electronegative element (FON)

Examples:– H2O

– NH3

Properties – universal solvent (H2O), unique properties (H2O)

Hydrogen Bonding in H2O

H-bonding is especially strong in water becausethe O—H bond is very polar

There are 2 lone pairs on the O atom

Accounts for many of water’s unique properties.

Intermolecular Forces

Dipole-Dipole – interactions between 2 polar bonds or molecules

Examples:– sugar and H2O– acids

Properties – produce acids, dissolve molecular (organic) solids in H2O

Intermolecular Forces Dispersion (aka London Dispersion,

Induced Dipole) – interaction that is proportional to the number of e- and proportional to the size of the e- cloud– Results from motion of electrons

Examples: non-polar molecules

Properties – help explain the states of matter and the states of the halogens

Hybridization

refers to a mixture or a blending Biology – refers to genetic material Chemistry – refers to blending of

orbitals Remember, orbitals can only predict

an area in space where an e- may be located.

Sometimes blending orbitals can produce a lower, more stable bonding opportunity.

Orbital hybridization occurs through e- promotion in orbitals that have similar energies (i.e. same energy level).

Hybridization Cont. Hybridization occurs WITHIN the

atom to enhance bonding possibilities.

Do not confuse this concept with orbital overlap (bonding).

Hybridization is a concept used to explain observed phenomenon about bonding that can’t be explained by dot structures.

EXAMPLES – draw box diagrams for Be, B, and C (use noble gas core).

How do I know if my central atom is hybridized?

If your central atom is B, Be, C, Si, or Al then it is hybridized.

If your molecule has multiple bonds in it then it is hybridized.– Double bonds – sp2 hybridized– Triple bonds – sp hybridized

Sigma & Pi Bonds Sigma Bond (σ) – combination of orbitals that is

symmetrical around the axis connecting the two nuclei. Pi Bond – parallel overlap of p orbitals causing bonding

electrons to be found above & below the bond axis.Sigma bonds: 2 “s” orbitals overlapping

1 “s” and 1 “p” orbital overlapping

2 “p” orbitals overlapping (same axes; end-to-end)

Pi bonds: 2 “p” orbitals overlapping (parallel axes; side-by-side)