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
Theoretical, actual, percent yield
Theoretical yield: mass product based on stoichiometry (max)
Actual yield: mass obtained in the lab
Percent yield = (actual yield / theoretical yield) * 100
Solutions, solute, solvent
Solutions: homogeneous mixture of 2+ substances
Solute: smaller amount
Solvent: larger amount (ie. water is universal solvent)
Why is water an excellent solvent?
Polar nature & ability to form hydrogen bonds
Water can dissolve polar & ionic substances
Ionic compounds in water:
Forces of attraction between solute & solvent particles dissolve
Not all ionic substances are soluble in water
Solubility depends on attraction between ions & water
Polar covalent compounds:
strong, weak, non ELECTROLYTES
Intermolecular hydrogen bonding
Strong electrolytes: dissociate completely in aqueous solutions (conduct electricity)
Weak electrolytes: produce fewer number of ions (don’t conduct well)
Non electrolytes: dissolve in water but don’t produce ions (don’t conduct)
Concentration solution:
amount of solute in a given quantity of
Molarity:
moles of solute in 1L of solution
n = mass / molar mass
Dilution: adding additional solvent to a solution
Dilution: adding additional solvent to a solution
Dilute: small amount of solute dissolved
Concentrated: large amount of solute dissolved
Dillution equation
M: molarity
V: volume (initial & final)
Aqueous chemical reaction types:
Aqueous chemical reaction types: Precipitation reaction
2 soluble ionic compounds react to give an insoluble product (precipitate)
Soluble salts (aq) form clear solutions, insoluble salts form precipitates (s)
Aqueous ionic reaction can be expressed in 3 different ways:
- Molecular equations
- Ionic equation
- Net ionic equation
Molecular equations:
all reactants/products as if they were intact, undissociated compounds
Complete ionic equations:
all soluble ionic substances dissociated into ions
Net ionic equation:
eliminates spectator ions, shows only actual chemical change
What are spectator ions?
Spectator ions: not involved in actual chemical change (appear unchanged on both sides of ionic equation)
Solubility rules for Ionic Compounds in Water
Don’t need to memorize rules, but must know how to interpret
Arrhenius definition of acid & base
limited to aqueous solutions (water as solvent)
Arrhenius acid: acid is a substance that produces H+ ions (protons) in water
- hydronium ion = proton = H+ ion
Arrhenius base: substance that produces OH- ions in water
Bronsted-lowry definition of acid & base
much broader definition of acid & base
Acid: proton donor
Base: proton acceptor → have a lone pair
Can acids and bases act as electrolytes (conduct electricity)?
Strong acids & bases dissociate completely in aq solution → form strong electrolytes
Weak acids & bases: dissociate very weakly into ions → form weak electrolytes
Reversible reaction (⇌)
Neutralization reaction:
acid reacts with a base to form salt & water
Complete ionic equation:
- strong acids (SA) & strong bases (SB) dissociated
- weak acids (WA) & weak bases (WB) (undissociated)
Titration:
titrant
analyte
equivalence point
analytical technique to determine concentration of unknown acid/base solution
Titrant: solution known concentration
Analyte: solution unknown concentration
Equivalence point: amount titrant added is enough to completely neutralize analyte solution
How to determine the equivalence point?
- Add indicator: change color at/near equivalence point (ie. phenolphthalein)
- Color changes at the end point. Ideally, we want the end point to be very close to the equivalent point. - Titration curve
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)
Delocalized pi bond
Pi orbital extend over 2+ atoms
In lewis structure, this occurs when with resonance structures involving double & triple bond
Molecular orbital diagram (MO)
What does a +, 0, negative bond order mean?
+ BO: stable enough to exits
≤ 0: doesn’t exit
Precipitation reaction
2 soluble ionic compounds (aq) react to give an insoluble product (precipitate (s))
Redox reaction
transfer of electron(s) from one species to another
even though a solid could be produced, it is not precipitation because the reactants are not soluble salts
Oxidation number
assume all bonds are ionic
Break bond any two atoms, give both electrons to most electronegative atom (electron transferred from least → most electronegative)
Calculate similar to FC (periodic table - atom electrons)
Oxidation number rules
Identify a redox reaction
calculate oxidation # (state) of each atom before & after reaction
OIL RIG
Oxidation: increase in ON (loss electrons) → OIL RIG
Oxidizing agent/oxidant
Reduction: decrease in ON
(gain electrons)
Reducing agent/reductant
Cation displacement reaction:
metal with high reactivity in the activity series is added to a solution containing a cation with lower reactivity
- Identify the Reacting Metals:
- Locate the Metals on the Activity Chart
- Compare Reactivities: The metal that’s higher on the list is more reactive.
- Predict the Reaction: If the free metal (the one not in a compound) is higher on the activity series than the metal in the compound, the reaction is likely to occur. The free metal will replace the metal in the compound, forming a new compound and releasing the less reactive metal.
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
Ideal gas law
PV = nRT
Temperature must be in K
Standard temp & pressure (STP):
0ºC, 1atm
Standard molar volume:
1 mole of any ideal gas at standard temp & pressure occupies 22.4L
Determine molar mass of gas using ideal gas law:
Real gasses behave as ideal gasses
high temp and low pressure
less influenced by IMF, negligible volume
Standard solution:
Analyte:
Standard solution: solution known concentration
Analyte: solution unkown concentration
Equivalence point (indicator) in acid, base titration
moles H+ from acid = moles OH- from base
amount H+ ions in flask = amount OH- ion added
Endpoint point (indicator) in acid, base titration
slight XS of base (OH-), indicator changes color
Diprotic acid
donate two hydrogen ions (H⁺) per molecule in an aqueous solution
