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Chapter 9
Molecular Compounds
Covalent Bonding
If only the nonmetals in groups 5A, 6A, & 7A existed, ionic bonds couldn't form. The atoms in those groups need electrons so they are not willing to lose any. If two Hydrogen atoms are locked in a room together, how would they become stable?
Keep in mind, there are attractive and repulsive forces in a covalent bond. The nuclei are attracted to the electrons, but are still repelling one another.
The sharing of valence electrons between two nonmetal atoms
There is no transfer of electrons because there are no metals present to lose electrons.
Formation of Covalent Bonds
Atoms overcome the nucleus' repulsive forces due to their mutual attraction to the same electrons.
Neutral group of atoms joined together by covalent bonds. Several atoms and several covalent bonds can be present in a molecule, unlike the formula unit of an ionic compound. Atoms are attached by more than just electrical attraction.
a molecule consisting of two identical atoms There are 7: must memorize (Hint: Starts at 7, makes a 7)
H 2 , N 2 , O 2 , F 2 , Cl 2 , Br 2 , I 2 These are elemental compounds that are more stable as a compound than by themselves.
Chemical formula for molecular compounds.
Shows the type of atoms and the the number of atoms.
Subscripts are not always in lowest whole number ratio.
Ionic compounds are in lowest whole number ratio called formula units.
Does not give the structure or shape of the molecule
Lower melting and boiling points than ionic compounds. Most are gases or liquids at room temperature. Poor conductors of electricity
Why do you think they possess these properties?
Gilbert Lewis Stated
Sharing of electrons occurs if the atoms involved acquire the electron configurations of noble gases. Stable by sharing instead of losing or gaining electrons. Unpaired vs. Paired Electrons are significant in bonding
(Show Boxes) Oxygen It is slightly different when 2 carbon atoms form covalent bonds. A covalent bond is formed by the unpaired electrons of two atoms, but Carbon only has 2 unpaired electrons Carbon needs 4 more electrons, so it needs 4 unpaired electrons. 1s^2 2s^1 2p^3 , not 1s^2 2s^2 2p^2
Carbon's Dot Structure Unique
Sigma Bonds
Single Covalent Bonds (Greek letter: ) When electrons are shared, the valence orbitals of one atom overlap the valence orbital of the other atom. (Strong bonds) The bonding orbital is the region between the two atoms in which the shared electrons are likely to be found. Can be s/s, s/p, or p/p orbitals overlapping. The bond is being pulled straight toward the central atom.
The sigma bond is the area centered between the two atoms that are sharing electrons.
Images of Sigma Bonds
Multiple Covalent Bonds
More than two valence electrons are being shared be two atoms.
O 2 N 2
Multiple Covalent Bonds
More Multiple Covalent Bonds
H 2 O CH 4 Parellel Orbitals overlap (Greek letter:^ ) Sideways overlapping, which means the shared electrons are not being pulled toward the central atom. A Pi bond always accompanies a sigma bond when forming double and triple covalent bonds. Formed only by p orbitals (Weak bonds) Remember, p orbitals are dumbell shaped, which makes them different
Pi Bonds
Sigma vs. Pi Bonds Images of Pi Bonds
The shared electron pair of the Pi bond occupies the space above and below the line where the atoms are joined. Image looks like there is 2 pi bonds, but it is one large pi bond.
Strength of Covalent Bonds
Several factors control the strength of covalent bonds.
Remember: the nuclei of bonded atoms are attracted to each other's electrons but repelling their nuclei. Bond Length is the most important factor: the distance that separates the bonded nuclei. (Shorter is stronger) Bond length decreases with more shared electrons. Triple bonds are strongest, single bonds are weakest.
Bond Dissociation Energy
The amount of energy required to break a covalent bond, It is different for all compounds. In a chemical reaction, bonds are broken and formed. Endothermic: reactions in which more energy is required to break the bonds of the reactants than is released when bonds form. Exothermic: reactions in which more energy is released forming new bonds than is required to break the bonds of the reactants.
9.3 Molecular Structures
Lewis Dot Formulas: show which atoms are pairing up to form bonds
Structural formulas: show the arrangement of atoms in molecules and polyatomic ions.
Dashes are used in structural formulas 1 dash: 2 shared electrons 2 dashes: 4 shared electrons 3 dashes: 6 shared electrons
How many electrons are donated by each chlorine? How many unshared pairs are in the molecule? How many electrons are being shared? How many shared pairs are in the molecule?
C 2 H 6 O C 4 H 6 O 2
Mg has 2 valence electrons.
- How many covalent bonds must it form to be stable?
- How many electrons does it have to donate? How about Aluminum?
Mg
Emergency Bonds
A coordinate covalent bond is formed when one atom contributes both bonding electrons in a covalent bond. Once formed, they act as normal covalent bonds. Electrons are mobile, which allows them to be “fluid” Used in Polyatomic ion formation Arrows are used to indicate a coordinate covalent bond Ex.) CO, NH 4 +, H 3 O+, SO 3 , SO 4 2-
C=O
C=O
C=O
Oxygen is stable! 8 valence electrons &
Carbon needs 2 more electrons, but Oxygen is stable.
S S
S S
S
NH H^ O
Negative Polyatomic Ions
O
O
Electrons donated from metals
Resonance
Resonance structures occur when two or more valid electron dot formulas can be written for a molecule. Differ in the position of the electron pairs, not the position of the atoms. Ex. O 3 , CO 3 2- Same formula but different structure due to the locations of the electron pairs.
Exceptions to the Octet Rule
Sometimes it is impossible to write electron dot structures that fulfill the octet rule. Occurs whenever the total number of valence electrons in the species is an odd number. Only certain substances can form covalent bonds and have less that 8 valence electrons. Boron can only form 3 bonds.
H
Occurs in atomic orbitals during bonding. Hybrid = two different types of orbitals are involved. Shows which orbitals of a bonded atom have paired electrons. Hybrid orbitals can be occupied by a shared pair of electrons or an unshared pair of electrons. (aka: Lone pairs)
B
H
H H
C
H
H
H H
: HCN, CO
: BH
Bent triatomic
: NH
: CH ,CCl 4 , NH 4 +
105
Attractions Between Molecules Gases
No attraction Nonpolar Molecules
Liquids Dipole Attraction Polar Molecules
Solids Ionic Attraction Ions Form Crystals
Intermolecular Attractions (Between)
These are what hold molecules together.
Weaker than either an ionic or covalent bond. They are responsible for whether a molecular compound is a gas, liquid, or solid.
Attractions between polarized molecules
The weakest attractions between molecules. Not Bonds!!!!!! Three types are Dispersion forces, Dipole interations, and Hydrogen bonds Hydrogen > Dipole > Dispersion
Van der Waals forces
The weakest of all intermolecular interactions. Thought to be caused by the motion of electrons, and remember, electrons are always moving. Strength of dispersion forces increases as the number of electrons in a molecule increases Electrons are not lost or gained
Due to movement, the electrons move to one side and create a separation of charge.
Dispersion forces Dipole Interactions
Electrostatic attractions occur between the
Similar to ionic bonding, but much weaker attraction.
Ion - Dipole Interactions Dipole - Dipole Interactions
Hydrogen Bonds
Strongest of all intermolecular attractions. Must involve hydrogen! Dipole interactions with hydrogen. An atom or molecule is attracted to a Hydrogen atom that is already bonded to an atom with high electronegativity. The covalently bonded hydrogen becomes slightly positive. Unshared electron pairs and atoms with high electronegativity become attracted to the slightly(+)^ Hydrogen.
Why is there so much water?
Water molecules are polar. The oxygen atom becomes slightly negative and each hydrogen becomes slightly positive. This causes an intermolecular attraction between water molecules. The attraction water molecules have for one another is called Hydrogen bonding.
