MCAT General Chemistry Chapter 3: Bonding and Chemical Interactions

Octet Rule

• Many atoms form bonds according to this rule 

  • An atom tends to bond with other atoms so that it has eight electrons in its outermost shell, forming a stable octet 

Exceptions to Octet Rule

• Incomplete octet

• Expanded octet

• odd numbers of electrons

Exceptions to Octet Rule: Expanded Octet

• any element in period 3 and greater can hold more than 8 electrons (ex. Phosphorous, sulfur, chlorine, etc.)

Exceptions to octet rule: incomplete octet

• describes elements that are stable with fewer than 8 electrons in their valence shell

  • hydrogen

  • helium

  • lithium

  • beryllium

  • boron

Exceptions to Octet Rule: Odd Numbers of Electrons

• any molecule with an odd number of valence electrons cannot distribute those electrons to give eight to each atom 

Common Elemetns that Abide by Octet Rule

• Carbon

• Nitrogen

• Oxygen

• Fluorine

• Sodium

Ionic Bonding

• one or more electrons from an atom with a low ionization energy (typically a metal) are transferred to an atom with a high electron affinity (typicallu a nonmetal) 

  • Resulting electrostatic attraction between opposite charges holds bond together 

Covalent Bonding

• an electron pair is shared between two atoms, typically non-metals, that have relatively similar values of electronegativity 

  • Degree to which the pair of electrons is equally or unequally shared is dependent on the degree of polarity in the covalent bond; this results

    • Nonpolar = shared equally

    • Polar = shared unequally; greater electronegativity difference (0.5-1.7)

    
    

Principles of Ionics Compounds

• Cation (typically a metal) loses electrons and anion (typically a nonmetal) gains electrons 

• Difference in electronegativity must be greater than 1.7

• Generally between metal and nonmetal  

• Electrostatic attractions permits them very high melting and boiling points 

• Readily dissolve in water and other polar solvents 

• In solid states, form crystalline lattices 

Crystalline lattices

consist of repeating positive and negative ions that maximise attractive forces between opposing charges and minimize repulsive forces between like charges 

NaCl crystal lattice structure 

Principels of Covalent Compounds

• Atoms have similar electronegativities and attract 

• When atoms of similar tendency to attract electrons form a compound, it is energetically unfavorable to create ions 

• Contain discrete molecular units relatively weak intermolecular interaction

• Lower melting and boiling points 

• Joined by single, double or triple bonds 

• Exhibit variable bond order, length, energies, and polarities

Bond Order

number of shared electron pairs between two atoms

Bond Length

average distance between the two nuclei of atoms in a bond 

• As the number of shared electron pairs increases, two atoms are pulled closer together 

Bond Energy

energy required to break a bond by separating its components into their isolated, gaseous, atomic states 

• The greater the number of electron pairs between atoms, the more energy required to break the bond 

Polartiy

occurs when two atoms have a relative difference in electronegativities 

• Atom with higher electronegativity gets larger share of electron density

• Polar bond creates dipoles

  • More electronegative element takes on a partial negative dipole charge 

  • Less electronegative element takes on a partial positive dipole charge 

  • Polar molecules can be nonpolar if dipole moments cancel out

Dipole Moment

Equation 3.1:

• Describes the point at which their is separation of positive and negative charges

p = qd

• P is dipole moment 

• Q is magnitude of charge

• D is displacement vector

Coordinate Covalent Bonds

• both of shared electrons are contributed by only one of the two atoms by one atom attacking the unhybridized p-orbital of the other 

Covalent Bond Notation

• Bonding electrons: electrons in valence shell involved in covalent bond

• Nonbonding electrons: electrons in the valence shell that are not involved in covalent bonds 

Lewis Structures

• keep track of bonded and non bonded electron pairs 

  • When lewis structures for an atom differ in bond connectivity or arrangement, then they represent different compounds 

  • If they say the same connectivity of arrangement, they represent resonance structures 

  • Arrangements that minimize formal charge are most stable

Drawing Lewis Structures

1. Central atom is the most electronegative 

• Exception is in HCN, where H occupies an end position 

2. Count # of valence electrons for each atom of a molecule 

3. Draw single bonds between central atom and the atoms surrounding 

4. Complete the octets of atoms bonded to the central atom 

   1. Hydrogen is an exception as it only holds 2 valence electrons

5. Place any extra electrons on central atom

   1. If central atom has less than octet, try to write double or triple bonds between central and surrounding atoms 

Formatl Charge

• When calculating, assume a perfectly equal sharing of all bonded electron pairs, regardless of difference in electronegativity 

Formal Charge vs. Oxidation Number

• Differs from oxidation number which assumes that the more electronegative atom has 100% share of the bonding electron pair

Resonance Hybrid

•  the nature of the bonds of a compound with multiple structures is actually a hybrid of all these structures 

Principles of Resonance Hybdrids

• More stable resonance structures contribute more to the resonance hybrid than others 

• Lewis structure with small or no formal charge is preferred

• Resonance structures with negative formal charges placed on more electronegative atoms is more stable 

Valence Shell Electron Pair Repulsion Theory (VSEPR)

• Uses Lewis structures to predict molecular geometry of covalently bonded molecules 

• States that three dimensional arrangement of atoms surrounding a central atom is determined by the repulsions between bonding and nonbonding electron pairs in the valence shell of the central atom 

• Electrons arrange themselves as far apart as possible, minimising repulsive force

Count the regions of density with bonding electrons, and consider how they are arranged as a result of the presence of lone pairs

Electron Geometry vs. Molecular Geometry

• Electronic geometry describes spatial arrangement of all pairs of electrons, nonbonding and bonding, implicates ideal bonding angle 

• Molecular geometry describes spatial arrangement of only bonding pairs of electrons 

Focus on coordination # , number of atoms that surround/bonded to central atom

Quantum Numbers Specify an Atomic Orbital

• When l = 0, s

• When l = 1, p 

  • Three orbitals around x,y, and z axises

• When l = 2, d

• When l = 3, f

Molecular Orbitals

• When atomic orbitals interact, this describes the probability of finding the bonding electrons in a given space

• Bonding orbital: Formed when the signs of the atomic orbitals are the same

• Antidbonding orbital: formed when the signs of the atomic orbitals are different 

Sigma Bonds

• formation of a molecular bond between orbitals head to head, on the same plane

  • Allows free rotation 

Pi Bonds

• formed when there are two parallel electron cloud densities 

  • Do not allow free rotation 

Intermolecular Forces

From lowest to greatest strength 

1. Van der waals (london dispersion forces) 

2. Dipole-dipole

3. Hydrogen bonds

4. Ion-dipole

5. Ionic

London Disperison Forces (van der waals forces)

• Randomized short lived dipole moments between bonding electrons in nonpolar covalent bonds

  • Interact with dipoles of neighboring electron clouds to form more dipoles

  • Weakest of all intermolecular forces, dipoles shift moment-to-moment

• Strength is dependent on degree to which by which molecules can be polarized, (larger molecules are more easily polarized)

Dipole-dipole Interactions

• The interactions of oppositely charged ends of molecular dipoles that are organized closest to each other in a polar molecule

• Present in solid and liquid phases, but negligible in gas phase 

• Differ from london dispersion as they exude a longer duration and higher level of permanence

Hydrogen Bonds

• A specific strong form of dipole-dipole interaction that may be intra- or intermolecular 

• Are not actually bonds (no sharing or transferring)

• Exist when hydrogen is bonded to one of 3 highly electronegative atoms (fluorine, oxygen, or nitrogen)

• Positively charged hydrogen atom interacts with partial negative charge of these atoms on nearby atoms

• Have higher melting and boiling points than substances of similar weights that lack hydrogen bonds