Bonding and Intermolecular Forces - Thermodynamics Flashcards
How to calculate formal charge
formal charge = valence electrons - sticks (bonds) - bonds (lone pairs)
where do you place positive charges
on the less electronegative atoms
where do you place negative charges
on the more electronegative atoms
2 groups
- sp
3 groups
- sp2
4 groups
- sp3
sp2 hybridization with no lone pairs
trigonal planar
sp2 hybridization with one lone pair
bent
sp3 hybridization with no lone pairs
tetrahedral
sp3 hybridization with one lone pair
trigonal pyramidal
sp3 hybridization with 2 lone pairs
bent
sp hybridization with no lone pairs
linear
the strength of a chemical bond is dependent upon
- more electrons shared = stronger bond
- shorter distance between atoms = stronger bond
stronger bonds
- higher bond dissociation energies
- more energy to break the bond
- always an endothermic process
types of bonds
- intramolecular (nuclear)
- intermolecular
intramolecular (nuclear bonds)
- covalent
- coordinate covalent
- metallic
covalent
- sharing of electrons
- polar
- nonpolar
polar covalent
- 2 atoms share valence electrons
- different electronegativity - unequal sharing
- dipole moment (partial charges)
nonpolar covalent
- 2 atoms share valence electrons
- similar electronegativity = equal sharing
coordinate covalent bond
- metal/lewis acid - acceptor
- ligand/lewis base - donor
- ligand brings both electrons for bond.
metallic bond
- sea of free floating electrons
- metals like to lose electrons
ionic bond
- one gives other takes electrons
- very different electronegativity
- atoms with formal charges
- cation and anion
- Na+ Cl-
Intermolecular forces
- Hydrogen bond
- Dipoles
- London Dispersion (Van Der Waals)
Dipoles
- ion-dipole
- dipole-dipole
- dipole-induced dipole
- easily cleaved
hydrogen bond
- strong intermolecular forces
- produced between very polar molecules
- 2 criteria
- FON w/ lone pairs
- FON with hydrogen
ion-dipole
- polar molecule + ion
- higher charged ion and more polar molecule make for stronger force
dipole-dipole
- polar covalent + polar covalent
- partial charges attract
dipole-induced dipole
- polar covalent + nonpolar covalent
- partial charge on one causes temporary partial charge on the other
- very short lived so very easily cleaved
London dispersion
- nonpolar molecules
- all molecules
- instantaneous induced dipoles
- electron cloud is deformed by collisions that produce temporary but small dipoles
enthalpy
- energy stored within chemical bonds or any attractive force
high energy reactants and low energy products
- exothermic
- ΔH < 0
low energy reactants and high energy products
- endothermic
- ΔH > 0
breaking bonds
- exothermic
forming bonds
- endothermic
enthalpy of formation
- the amount of energy associated with forming one mole of a compound from its constitutive elements in their standard states
Δ H of a standard state
0
Hess’s law
- combine two or more reactions and to find their total
reversing the direction of a reaction
changes the sign of ΔH and ΔS
changing the stoichiometric coefficient of a reaction
you must scale the value of the ΔH or ΔS for the reaction
Entropy
- the measure of disorder/potential randomness
- increasing number of particles
- changing from solid to liquid to gas
- increasing the temperature
- increasing volume
Gibb’s free energy
- energy available to do work
- ΔG = ΔH-TΔS
spontaneous process
- exergonic
- high energy reactants, low energy products
nonspontaneous process
- endergonic
- low energy reactants, high energy products
core electrons
- do nothing!!
