Bonding, IMF, Thermodynamics Flashcards
Lewis Dot Structure
Each dot represents a valence electron, single bonds (two electrons) are shared between two atoms in the bond, double bonds (four electrons), triple bonds (six electrons)
Formal Charges
Tell us whether or not the atoms are sharing their electrons in the best way possible (not representative of actual charge of the element)
Formal Charge = Valence - (1/2 of Bonded Electrons) - # lone paired (nonbonded electrons)
Resonance Structures
Accurate depiction of a molecule’s bonding that cannot be one in one single structure (often needed when there are double/triple bonds or formal charges and lone pairs)
Resonance Hybrid
Depiction of resonance structures (ex. dotted line drawing of SO2)
Major vs. Minor Resonance
Major resonance forms are those structurally favorable (all atoms have octet and no formal charges)
Bond Dissociation Energy
Energy required to break a bond homolytically
Homolytic Bond Cleavage
One electron of the bond being broken goes to each fragment of the molecule (formation of radicals)
Heterolytic Bond Cleavage
AKA Dissociation, both electrons of the electron pair end up on the same atom (forming cation and anion)
Bond length
The distance between two nuclei that are bonded to one another
Greater the s character, shorter the bond, stronger the bond
Similar Bonds
Varying of atomic radii makes it hard to compare bond lengths/BDE of unlike bonds
Bond Order
Directly related to bond length
Covalent Bonds
Can be nonpolar or polar, with dipole moment
Nonpolar - electron density between two nuclei density is even
Polar - electron density is uneven bc electronegativity attracts electrons and creates a more partially negative side on one side of the molecule
Coordinate Covalent Bond
One atom donates both of the shared electrons in a bond
Ionic Bond
The gaining/losing of electrons that results in a cation/anion held together by electrostatic attraction
Valence Shell Electron Pair Repulsion (VSEPR) Theory
Prediction of shapes of different molecules
Guided by the fact that since electrons repel one another, electron pairs (whether bonding or non bonding) attempt to move as far as possible from one another
Hybridization
Method to rationalize observed chemical/structural trends. Character is based on how many things are attached to the atom in question (p number = number of things its attached to minus 1) *inclusive of lone pairs
sp - connected to two things
sp2 - 3 things
sp3 - 4 things
Sigma bonds
2 electrons localized between two nuclei - formed by end-to-end overlap
Pi bonds
2 electrons localized to region that lies on opposite sides of the plane (due to already existing sigma bond) side-to-side overlap
Therefore less stable than sigma bonds
Molecular Polarity
Polarity can apply to a bond as a whole -> if a molecule contains no polar bonds, cannot be polar. If a molecule has 2+ polar bonds that have even electron density, makes them polar
IMF
Weak interactions that take place between neutral molecules (NOT BONDS)
Ion-dipole forces
Polar molecules attracted to ions
Dipole-Dipole Forces
Attractions between positive end of one polar molecule and negative end of another molecule
Ex. Hydrogen bonding (strongest dipole-dipole force)
Dipole-induced Dipole Force
Permanent dipole in one molecule may induce dipole in another neighboring nonpolar molecule (momentary)
London Dispersion (instantaneous Dipole-induced-dipole)
Very weak transient interactions between instantaneous dipoles in nonpolar molecules.
Strength of dispersion forces increase as a result of size/molecular weight of molecule increasing
Van Der Waals
Dipole/H bonding/London collective name, sometimes referring to only London Dispersion
Hydrogen Bonding (NOF)
1) must have a covalent bond between H and either NOF
2) must have another molecule with a lone pair on either NOF atom
Ex. Water
More hydrogen bonding increases boiling point
Vapor Pressure
Pressure exerted by gaseous phase of a liquid that evaporated from exposed surface of the liquid
Weaker a substance’s IMF, higher the vapor pressure, more easily it is to evaporate
Also temperature dependent (increases in kinetic energy increase vapor pressure)
Volatile
Liquids (with high vapor pressure) that are easily vaporized
Ionic Solids
Held together by electrostatic attraction between cations and anions in a lattice structure. This lattice structure held together by ionic bonds as well.
Most ionic substances are solid at room temperature, with bonds being strong
Bond strength increases as atomic radii/space between the two elements decrease; varying electronegativities (greater charge differences, stronger attraction)
Network Solids
Connected lattice of covalent bonds. Network of intermolecular forces are same as intramolecular forces (covalent bonding) which makes them very strong. Hard solids at room temperature
Diamonds, quartz are examples
Metallic Solids
Covalently bound lattice of nuclei and their inner shell electrons surrounded by sea/cloud of electrons
At least one valence electron (the conduction electrons) per atom not bound to anything in particular and freely moving
Creates excellent conductivity of heat and electricity (malleable)
All metals (except Hg solid at room temp)
Molecular Solids
Held together by three Intermolecular interactions (H-bonding, Dipole-dipole, and London)
Significantly weaker than ionic, network or metallic bonds. Much lower melting/boiling points
Liquids/gas at room temperature
Dipole Moment
Two atoms involved in the bond differ in electronegativity
