Unit 2: Polar Covalent Bonds; Acids and Bases Flashcards
Describe how differences in electronegativity give rise to bond polarity
A polar bond is a covalent bond in which there is a separation of charge between one end and the other
- one end is slightly positive, the other slightly negative
Electronegativity inductive effect
Polar covalent bond
A polar bond is a covalent bond in which there is a separation of charge between one end and the other
- one end is slightly positive, the other slightly negative
What if two atoms of equal electronegativity bond together?
If the atoms are equally electronegative (aka, they’re the same atom) both have the same tendency to attract the bonding pair of electrons so it will be found on average halfway btw the 2 atoms (still in molecular orbital)
- This bond can be thought of as a “pure” covalent bond- where electrons are shared evenly btw the 2 atoms
When is an ionic bond, opposed to a covalent bond, formed?
If B is a lot more electronegative than A, then the electron pair is dragged right over to B’s end of the bond; A has lost control of its electron, and B has complete control over both = ions have been formed
Dipole moment
Mathematically, dipole moments are vectors; they possess both a magnitude and a direction. The dipole moment of a molecule is therefore the vector sum of the dipole moments of the individual bonds in the molecule. If the individual bond dipole moments cancel one another, there is no net dipole moment
Explain how dipole moments depend on both molecular shape and bond polarity
Formal charge
Compares the number of electrons around a “neutral atom” (an atom not in a molecule) versus the number of electrons around an atom in a molecule. Assigned to an atom in a molecule by assuming that electrons in all bonds are shared EQUALLY, regardless of electronegativity. The sum of the formal charges of each atom must be equal to the overall charge of the molecule or ion.
We assign electrons in the molecule to individual atoms according to these rules:
1. Non-bonding electrons are assigned to the atom on which they are located
2. Bonding electrons are divided equally btw the 2 bonded atoms
FC = (# of valence electrons in free atom) - (# of LP electrons) - (1/2 # of bond pair electrons)
Valence electrons
Bonding electrons
Non-bonding electrons
Carbocations, Carbanions, Carbon Radical
Carbocations: occur when a C has only 3 bonds, and no LPs. Have only 6 valence electrons and a FC of +1
Carbanions: occur when a C has 3 bonds plus 1 LP. Have 8 valence electrons and a FC of -1
Carbon Radical: has 3 bonds and a single, unpaired electron. They have 7 valence electrons and a FC of 0
Tetravalent
Carbon is tetravalent, meaning that it commonly forms four bonds
Resonance form
Sets of Lewis structures that describe the delocalization of electrons in a polyatomic ion or a molecule
- based on molecular orbitals NOT averages of different bonds btw atoms
- we describe the electrons in such molecular orbitals as being delocalized, that is they cannot be assigned to a bond btw 2 atoms
6 rules for estimating stability of resonance structures
- The greater the number of covalent bonds, the greater the stability since more atoms will have complete octets
- The structure with the least number of formal charges is more stable
- The structure with the least separation of formal charge is more stable
- A structure with a negative charge on the more electronegative atom will be more stable
- Positive charges on the least electronegative atom (most electropositive) is more stable
- Resonance forms that are equivalent have no difference in stability and contribute equally (eg. benzene
When are resonance structures used?
When one Lewis structure for a single molecule cannot fully describe the bonding that takes place between neighboring atoms relative to the empirical data for the actual bond lengths between those atoms
Resonance hybrid
The net sum of valid resonance structures is defined as a resonance hybrid, which represents the overall delocalization of electrons within the molecule
- A molecule that has several resonance structures is more stable than one with fewer
Delocalization
In resonance structures, the electrons are able to move to help stabilize the molecule, this movement of the electrons is called delocalization
Resonance Hybrid
The combination of all resonance structures
- Though the FC closest to zero is the most accepted structure, in reality the correct Lewis structure is a combination of all the resonance structures
Bronsted-Lowry Acid
Acids are defined as being able to donate protons in the form of hydrogen ions
Bronsted-Lowry Base
Bases are defined as being able to accept protons
Conjugate Acid
Substance on the right side of the equation; base turns into conjugate acid
Conjugate Base
Substance on the right side of the equation; acid turns into a conjugate base
Acidity constant, Ka
*Also known as the acid dissociation constant
The relative acidity of different compounds or functional groups, aka their relative capacity to donate a proton to a common base under identical conditions, is quantified by a number called the dissociation constant
Any particular acid will always have the same pKa (assuming we are talking about an aqueous solution at room temp) but different aq solutions of the acid could have different pH values, depending on how much acid is added to how much water.
Equilibrium constant, Keq
Lewis acid
An electron-pair acceptor
Lewis base
An electron-pair donor
Lewis theory of acids and bases
A Lewis acid is an electron-pair acceptor and a Lewis base is an electron-pair donor. This covers Bronsted-Lowry proton transfer rxns, but also includes rxns in which no proton transfer is involved
Coordinate or Dative Bonds
Any covalent bond that arose bc one atom brought a pair of its electrons and donated them with another
Nucleophile
Electron donor (aka a Lewis base)
Electrophile
Electron acceptor (aka Lewis acid)
Dipole-dipole forces
London dispersion forces
Hydrogen bond
Intermolecular forces
Hold molecules together in a liquid or solid, 3 types of forces
Noncovalent interaction
Van der Waals forces
How do changes between the solid, liquid, and gaseous states almost invariably occur for molecular substances without breaking covalent bonds?
Due to the large difference in the strengths of intra- and intermolecular forces
Intramolecular vs Intermolecular forces
Intramolecular forces
- such as covalent bonds that hold atoms together in molecules and polyatomic ions
Intermolecular forces
- hold molecules together in a liquid or solid
- generally much weather than covalent bonds
- determine bulk properties such as melting points and boiling points
- electrostatic = they arise from the interaction btw positively and negatively charged species; bc electrostatic interactions fall off rapidly with increasing distance btw molecules, intermolecular interactions are most important for solids and liquids where molecules are close
3 kinds of intermolecular interactions
- dipole-dipole - van der Waals force
- London dispersion - van der Waals force
- hydrogen bond
