Exam 2 Flashcards
Lewis structure rules
- Atom lowest electronegativity in center (except for H cannot be placed as center)
- Find total number of v.e
(Add +1 v.e for negative charge, subtract -1 v.e for positive charge) - Connect all atoms using single bonds (lines), add remaining electrons to terminal atoms (pairs)
- If central atom has < 8 electrons, move lone pair from outside atom to make multiple bonds
lewis exceptions
Resonance structures:
same arrangement of atoms but diff placement of electrons
distribution of electrons is an avg of all lewis structures
Resonance hybrid: superposition/avg of two resonance structures
How do lewis structures depict electrons incorrectly?
depict electrons as localized
In nature, electrons are delocalized: density spread over entire molecule
Bond order
(# bonding e - # antibonding e) / 2
Higher bond order = stronger bond = shorter bond length = higher bond energy
Equivalent vs. non-equivalent resonance structure:
Equivalent resonance structure: same distribution formal charges
Non-equivalent resonance structure: different distribution FC
(Don’t equally contribute to resonance hybrid)
Formal charge (FC):
FC = #v.e periodic table – #v.e in bonded atom
hypothetical charge an atom would have if we could redistribute the electrons in the bonds evenly between the atoms
Sum formal charges on all atoms = overall; charge molecule
Criteria for choosing greatest contribution to resonance hybrid:
What are some drawbacks of lewis structures?
doesn’t explain shape/geometry of molecule (effect properties)
Valence Shell Electron Pair Repulsion Theory (VSEPR)
predict shape/geometry molecule from lewis dot structure
Molecular shape notation
Electron pair geometry:
all regions where electrons, bonds, lone pairs are located
Same as molecular structure when there are no lone pairs around central atom
Molecular structure:
location of the atoms, not electrons
Valence bond theory (VB):
When atoms approach each other for bonding, atomic orbitals of one atom overlap with the atomic orbitals of the other atom
Each of the atomic orbitals that overlap should have 1 electron with opposite spin
After orbitals overlap, pair of electrons occupy the overlapped region
Strength bond: greater orbital overlap (closer nuclei is to bonded electrons) = stronger bond
Hybridization:
Hybridization: orbital mixing to form hybrid orbitals (linear combination of atomic orbitals)
- Number hybrid orbitals formed = number of atomic orbitals combined
- Hybrid orbitals formed are equivalent in shape and energy
- Hybrid orbitals are more effective in forming bonds than unhybridized orbitals
- Hybrid orbitals orient themselves in 3D to max distance between then and min repulsions between electrons
Hybridized vs non-hybridized orbitals
Non hybridized orbitals are non equivalent in energy and shape, whereas hybridized orbitals are equivalent in shape and energy
rule for number of hybrid orbitals?
Number of electron groups around central atom
Sigma vs. pi bond
Sigma bonds (σ): electron pair is in the region centered on the internuclear axis
- Lobes point toward each other (end to end overlap)
- Free rotation around sigma bond
Pi bonds (π): electron pair in the region above and below the internuclear axis
- Orbitals that are parallel to each other (sideways overlap)
- Restricted rotation around pi bond (transiomer)
π bond weaker than σ bond bc sideways overlap is less effective
Polar vs. nonpolar
POLAR:
- asymmetric electron distribution around molecule
- Dipole moment: measure polarity
- Arrow pointing toward more electronegative
NONPOLAR:
- symmetric electron distribution, electronegativity difference = 0
In a molecule with 2+ atoms, what does the overall polarity depend on?
Individual bond polarity (polar bonds: asymmetry, individual bond dipoles dont cancel: asymmetry in molecular shape)
Shape of molecule
Effect of molecular polarity on behavior:
Boiling point (intermolecular forces)
Intermolecular forces (IMF):
Intermolecular forces: attraction between molecules with partial charges, or between ions/molecules
Weaker than bonding forces (F = q1q2 / r^2)
IMF: Dipole dipole forces
ALL molecule with net dipole moment have dipole dipole forces
- Positive pole of one polar molecule attracts the negative pole of another
- Same molar mass: ↑ dipole moment = ↑dipole dipole forces = ↑higher boiling point
IMF: hydrogen bond
extreme form of dipole dipole forces
Hydrogen atom covalently bonded to a small, highly electronegativity atoms with lone electron pairs (N, O, F)
hydrogen bond, stronger than dipole dipole (higher boiling point)
Boiling point increases with size apart from hydrogen bond
IMF: London dispersion forces
Between nonpolar molecules/atoms
Instantaneous dipole in one particle induces a dipole in another, resulting in an attraction
Larger particles, greater molar mass, more easily polarizable, stronger dispersion forces
Molecular orbital theory (MO):
When individual atoms combine, atomic orbitals in bonding atoms combine to form molecular orbitals (spread out over entire molecule)
Valence bond theory vs. molecular orbital theory
Electrons are localized in VB but in MO they are delocalized over entire atom
bond order
+ bond order: # bonding > # antibonding → molecule stable (exists)
– or 0 bond order: # bonding # antibonding → molecule unstable (doesn’t exist)
Higher bond order = stronger bond
Relationship between # electrons in bonding, antibonding molecular orbital & stability
Higher number of electrons in bonding MO = greater stability
Higher number of electrons in antibonding orbitals = lower stability
Percent composition:
% by mass of each element in compound
- Elemental makeup of a compound defines its chemical identity
- Useful to determine relative abundance of a given element in different compounds of known formula
Chemical formula:
relative numbers, not masses, of atoms in substance
Empirical formula
derived from experimentally measured element masses by:
Empirical formula mass:
avg atomic masses of all atoms in empirical formula
Stoichiometry:
quantitative relationships between amount of reactants & products in chemical equation
- atoms are neither created nor destroyed
- Identity of reactants & products are experimentally determined → cannot change identity (chemical formula) in balancing
- Start with more complicated molecules first (greatest #atoms)
Limiting reactant:
completely consumed first, limits amount of product formed
Molarity:
moles of solute in 1L of solution
n = mass / molar mass
Dillution equation
M: molarity
V: volume (initial & final)
Aqueous chemical reaction types:
Solubility rules for Ionic Compounds in Water
Don’t need to memorize rules, but must know how to interpret
Relative energies of p-molecular orbitals in homonuclear diatomic molecules (2nd period)
Paramagnetic molecules
have at least one unpaired electron
diamagnetic molecule
All electrons are paired
Molecular orbital diagram vs lewis structure in explaining paramagnetic nature of O2 molecule
Atomic orbital vs hybrid orbital vs molecular orbital
large size and planarity of the naphthalene molecule allow its London dispersion forces to be substantial enough to make it a solid at room temperature, despite these forces being weaker than the hydrogen bonds in water.
Bonding orbital
increased electron density directly between nuclei
Antibonding orbital
electrons are located away from region between two nuclei (nodal plane)
Molecular orbital diagram (MO)
Oxidation number rules
Atmospheric pressure:
pressure exerted by column if air from top of atmosphere to surface of earth (higher altitude = smaller atmospheric pressure)
Barometer device to measure atmospheric pressure → pressure indicated by height (mm) of mercury column [1atm=760mmHg]
Gas pressure
Boyle’s Law
Charles law
Avogadro’s law
Determine molar mass of gas using ideal gas law:
