Chem/Phys Flashcards
Atomic Weight
The weighted average of the masses of the naturally occurring isotopes of an element, in amu per molecule or grams per mole
Mole
A unit used to count particles; represented by Avogadro’s number 6.022 x 10^23 particles
mass of sample/molar mass
Also the amount of substance that contains the same number of particles as a 12.00g sample of C12.
Isotopes
For a given element, multiple species of atoms with the same number of protons but different numbers of neutrons
Planck’s quantum theory
Energy emitted as electro-magnetic radiation from matter exists in discrete bundles called quanta
Bohr’s Model of H Atom
Energy of electron = E = (2.18 * 10^-18 J/e-)/n^2
EM energy of photons = E = hc/wavelength
(h = 6.626 x 10^-34 J/s)
Balmer series vs Lyman series
Group of hydrogen emission lines corresponding to transitions from upper levels n > 2 VS Group of hydrogen emission lines corresponding to upper levels n > 1 to n = 1
Absorption spectra
Characteristic energy bands where electrons absorb energy
Heisenberg Uncertainty Principle
Impossible to know the location and momentum of an electron at the same time
Quantum numbers
Principle (n) - the larger the integer value, the higher the energy level and radius of the electrons orbit; max # e- in energy level n = 2n^2
Azimuthal (l) - Subshells; four, corresponding to I = 0,1,2,3 are s, p, d, f; max # e- within subshell = 4l + 2
Magnetic (ml) - orbital within a subshell where highly likely to find electron; between 1 and -1
Spin (ms) - spin of a particle or intrinsic angular momentum; +1/2 or -1/2
Hund’s Rule
Within a given subshell, orbitals are filled such that there are a maximum number of half filled orbitals with parallel spins
Valence electrons
Electrons of an atom that are in its outer energy shell and available for bonding
Coordination Compounds
Lewis acid-base adduct with a cation bonded to at least one electron pair donor (including H2O); Donor molecules are called ligands and use coordinate covalent bonds.
Chelation
When the central cation in a coordination compound is bound to the same ligand multiple times
Hydrogen bonding
The partial positive charge of the hydrogen atom interacts with the partial negative charge located on the electronegative atoms (F, O, N) of nearby molecules
Dipole-dipole Interactions
Polar molecules orient themselves such that the positive region of one molecule is close to the negative region of another molecule.
Dispersion forces
The bonding electrons in covalent bonds may appear to be equally shared, but at any particular point in time they will be located randomly throughout the orbital, permitting the unequal sharing of electrons and leading to transient polarization and counterpolarization of the electron clouds of neighboring molecules
Units for rate constant?
zero-order - M/s
first-order - s^-1
second-order - M^-1s^-1
third-order - M^-2s^-1
Combustion reaction
A fuel, such as a hydrocarbon, is reacted with an oxidant, such as oxygen, to produce an oxide and water
Combination reaction
Two or more reactants form one product
Decomposition reaction
A compound breaks down into two or more substances, usually as a result of heating or electrolysis
Single-displacement reaction
An atom (or ion) of one compound is replaced by an atom of another element
Double displacement reaction
Also called metathesis reactions; elements from two different compounds displace each other to form two new compounds
Net ionic equation
Written showing only species that actually participate in the reaction
Neutralization reaction
A specific type of double-displacement reaction that occurs when an acid reacts with a base to produce a solution of a salt and usually water
Factors that affect reaction rates?
Reactant concentrations, temperature, medium, catalysts
Law of Mass Action
aA + bB –> cC + dD
Keq = [C]^c[D]^d/[A]^a[B]^b
Properties of equilibrium constant
Pure solids and liquids don’t appear in the expression
If Keq»_space; 1, little of the reactants compared to the products.
If Keq «_space;1, little of the products
compared to the reactants.
If Keq is close to 1, approximately equal amounts of the two
Le Chatelier’s Principle
Used to determine the direction of the reaction at equilibrium when subjected to stress (ie. change in concentrations, pressure, volume, or temp)
Isolated vs closed vs open systems
Isolated - no exchange of energy or matter
Closed - no exchange of matter but exchange of energy
Open - exchange of matter and energy
Isothermal Process
Constant temperature; ∆U = 0; first law is Q = W
Adiabatic process
Constant heat; Q = 0, first law is ∆U = -W
Isobaric process
Constant pressure
Isovolumetric (Isochoric) process
Constant volume; W = Q, first law is ∆U = Q
Endothermic vs Exothermic
Absorb energy (positive ∆H) vs Release energy (negative ∆H)
Heat absorbed or released in a given process
q = mc∆T (c - specific heat)
State functions
Describe the macroscopic properties of the system; pressure, density, temperature, volume, enthalpy, internal energy, free energy, and entropy
Enthalpy
Used to express heat changes at constant pressure
Standard heat of formation
The enthalpy change that would occur if one mole of a compound was formed directly from its elements in their standard states
Standard heat of reaction
the hypothetical enthalpy change that would occur if the reaction were carried out under standard conditions
Hess’s Law
the enthalpies of reactions are additive; the reverse of any reaction has the same magnitude with opposite sign
Bond dissociation energy
The average energy required to break a particular type of bond in one mole of gaseous molecules
Bond enthalpy
The standard heat of reaction can be calculated using the values of bond dissociation energies of particular bonds
Entropy
The measure of the distribution of energy (“randomness”) throughout a system
Reaction spontaneity by ∆H and ∆S signs
-/+ : spont at all temps
+/- : nonspont at all temps
+/+ : spont at high temps
-/- : spont at low temps
Reaction quotient
Once a reaction commences, the standard state conditions no longer hold; Q is the same equation as K
Pressure Equivalents
1 atm = 760 torr = 760 mmHg = 101,325 Pa
STP vs Standard Conditions
STP - 0˚C, 1 atm; used for gas law calculations
Standard conditions - 25˚C, 1 atm, 1 M concentrations; used for standard enthalpy, entropy, free energy, or emf
Boyle’s Law
P1V1 = P2V2
Charle’s Law
V1/T1 = V2/T2
Gay Lussac’s Law
P1/T1 = P2/T2
When do ideal gases behave unideally?
Small volume; high pressure or low temperature
Van der Waals equation of state
accounts for deviations from ideality that occur
(P + n^2a/V^2)(V-nb) = nRT
1 mol of Gas at STP is how many liters?
22.4
Dalton’s law of partial pressures
The total pressure of a gaseous mixture is equal to the sum of the partial pressures of the individual components
Calculate partial pressure of a gas
Pgas = Xgas*Ptotal
Kinetic molecular theory of gases
An explanation of gaseous molecular behavior based on the motion of individual molecules
Average molecular speeds
K = 1/2mv^2 = 3/2kBT
Root-mean-square speed
urms = √3RT/M
Colligative properties
Properties derived solely from the number of particles present, not the nature of the particles
Freezing point depression
∆Tf = iKfm (m is molality)
Boiling point elevation
∆Tb = iKbm (m is molality)
Osmotic pressure
π = MRT
Raoult’s Law
vapor pressure lowering; the partial vapor pressure of each component of an ideal mixture of liquids is equal to the vapor pressure of the pure component multiplied by its mole fraction in the mixture
Pa = XaP˚A; Pb = XbP˚b
Graham’s law of diffusion and effusion
Diffusion - occurs when gas molecules distribute through a volume by random motion
Effusion - occurs when gas flows under pressure from one compartment to another through a small opening
r1/r2 = 1√M1/M2
Solubility rules
1) All group 1 salts or ammonium cations are water sol.
2) All salts with nitrate (NO3-) or acetate anions are water sol.
3) All chlorides, bromides, and iodides are water soluble, except Ag+, Pb+ and Hg2+
4) All sulfate ion salts (SO42-) are water soluble, except Ca2+, Sr2+, Ba2+, and Pb2+
5) All metal oxides are insoluble with the exception of alkali metals, CaO, SrO and BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides.
6) All hydroxides are insoluble with the exception of alkali metals and Ca2+, Sr2+ and Ba2+
7) All carbonaes (CO3 2-), phosphates (PO4 3-), sulfides (S2-) and sulfites (SO3 2-) are insoluble, with the exception of alkali metals and ammonium.
Units of concentration
percent comp. by mass: mass solute/mass solution *100 mole fraction: moles solute/total moles molarity: moles/liter molality: moles/kg normality: g/L
Arheinius definition of acids and bases
Acids as source of H+ and bases as source of OH-
Brønsted-Lowry vs Lewis
Proton donors/acceptors vs Electron donors/acceptors
pH and pOH
-log[H+] or -log[OH-]
Kw
10^-14; pH + pOH = 14
Properties of Weak acids and Bases
HA (aq) + H2O (l) –> H3O+ (aq) + A- (aq)
Ka = [H3O+][A-]/[HA]
Kb = [B+][OH-]/[BOH]
Salt formation
Acids and bases may react with each other, forming a salt and (often) water in a neutralization reaction
Hydrolysis
The reverse of a neutralization, where salt reacts with water to produce the acid and base
Amphoteric species
Can act as either an acid or a base depending on chemical environment
Titration
A procedure used to determine the molarity of an acid or base by reacting a known volume of a solution of unknown concentration with a known volume of a solution with a known concentration. The half-equivalence point defines pH = pKa
Henderson-Hasselbach Equation
Used to estimate the pH of a solution in the buffer region where the concentrations of the species and its conjugate are present in approximately equal concentrations
pH = pka + log [A-]/[HA] or pOH = pKb + log [BH+/B]
Oxidation vs reduction
loss of electrons vs gain of electrons
Oxidizing agent vs reducing agent
is reduced vs is oxidized
Galvanic cells
∆G is negative; supply energy used to do work; place oxidation-reduction half reactions in separate containers (half cells) and connect allowing the flow of electrons
Electrolytic cells
Nonspontaneous –> electrical energy is required to induce a reaction; everything in one container.
Reduction potential
The tendency of a species to acquire electrons and be reduced; standard reduction potentials calculated under standard conditions (25˚C, 1 atm, 1 M)
E˚cell or emf
emf = E˚red,cathode - E˚red, anode
Spontaneity of electrochemistry
∆G = -nFEcell
Sn1
2 steps; likes polar protic solvents; 3˚ > 2˚ > 1˚ > methyl; rate = k[RL]; racemic products; strong nucleophile not required
Sn2
1 step; likes polar aprotic solvents; methyl > 1˚ > 2˚ > 3˚; rate = k[Nu][RL]; optically active and inverted products; favored with strong nucleophile
Nucleophilicity in protic solvents:
F- > Cl- > Br - > I- bc protic solvents can inhibit by protonation or H bonding
Nucleophilicity in aprotic solvents:
I- > Br- > Cl- > F-
Physical properties
Characteristics of processes that don’t change the composition of matter (ie. mp, bp, solubility, odor, color, and density)
Chemical properties
Have to do with the reactivity of the molecule with other molecules.
Conformational isomers
Different Newman projections; strain maxed in eclipsed conformations
Cyclic strain types
Angle strain - stretch or compress angles from normal size
Torsional strain - from eclipsing conformations
Nonbonded strain - from interactions with substituents on nonadjacent carbons; in cyclohexane, the largest substituent usually takes equatorial position to reduce
Configurational isomers
Can only be interchanged by breaking and reforming bonds; ie. enantiomers and diastereomers
Properties of alcohols
Higher BP than alkanes; weakly acidic hydroxyl hydrogen
Synthesis of alcohols
- Addition of water to double bonds
- SN1 and SN2
- reduction of carboxylic acids, aldehydes, ketones, and esters (*aldehydes and ketones with NaBH4, esters and acids with LiAlH4)
Reactions of alcohols
Substitution reactions after protonation and leaving group (tosylate) conversion
Levels of organic oxidation-reduction
0) alkanes
1) alcohols, alkyl halides, amines
2) aldehydes, ketones, imines
3) carboxylic acids, anhydrides, esters, amides
4) carbon dioxide
Organic redox definitions of oxidation and reduction
oxidation - loss of e-, fewer bonds to hydrogens, more bonds to heteroatoms (O, N, halogens)
reduction - gain of e-, more bonds to hydrogens, fewer bonds to heteroatoms
Organic oxidizing agents
Have high affinity for electrons (O2, O3, Cl2) or unusually high oxidation staes (Mn7+ in MnO4- or Cr6+ in CrO4 2-)
Organic reducing agents
Sodium, magnesium, aluminum, and zinc, which have low ENs and ionization energies.
Also metal hydrides such as NaH, CaH2, LiAlH4, NaBH4 (because of H- ion)
Strong and weak oxidizers
PCC - alcohol to aldehyde
KMnO4 (jones’s reagent) or alkali chromate salts - 2˚ alcohols to ketones and 1˚ alcohols to carboxylic acids
*tertiary alcohols cannot be oxidized
Alcohols as protecting groups
Can be used as protecting groups for carbonyls –> rxn with a dialcohol forms an acetal which can be removed with aqueous acid
Phenol acidity
Hydrogen of alcohol is particularly acidic because the anion is stabilized by resonance with the ring
Quinone synth? Hydroxyquinone importance/synth?
Oxidation of phenol (cyclohexane w vertical DBs and carbonyls at bottom and top corner)
further oxidation –> hydroxyquinone which have bio activity
Ubiquinone
Also called coenzyme Q, a vital electron carrier associated with complexes I, II, and III of the ETC
Q cycle
Reduction of ubiquinone to ubiquinol, rinse and repeat. In the ETC
Aldehyde properties
Higher bp than alkanes due to polarity, but not as high as alcohols because no H-bonding
Synthesis of aldehydes
- Oxidation of primary alcohols
- Ozonolysis of alkenes
Reactions of Aldehydes
-reactions of enols (michael’s addition) ?
-nucleophillic addition to a carbonyl
-aldol condensation
(aldehyde acts as nucleophile in enol form and electrophile in keto form; after aldol is formed, dehydration forms an a,B-unsaturated carbonyl)
-decarboxylation
Carboxylic acid properties
pKa around 4.5 bc of resonance stabilization of conj. base; EN atoms increase acidity via inductive effects; higher BP than alcohols bc of 2 hydrogen bonds
Synthesis of Carboxylic acids
- oxidation of primary alcohols with KMnO4
- hydrolysis of nitriles
Reactions of Carboxylic acids
- formation of soap by reacting carb. acids with NaOH, arrange in micelles
- nucleophilic acyl substitution (incl. ester formation and reduction to alcohols)
- decarboxylation
Lactam
Cyclic amides; named according to the carbon bound to the nitrogen (B lactams contain a bond between the B carbon and the N, etc. B = four membered ring, G = 5 etc)
Lactone
Cyclic esters; named by carbon bound to oxygen and by length of chain; ie. alpha-acetolactone, beta-propiolactone
Reactivity of Carb. Acid Derivatives
1) Acyl halides (most)
2) Anhydrides
3) Carboxylic acids and esters
4) Amides
*reactions down this chain are spontaneous via acyl substitution, up requires special catalysts and specific catalysts
Synthesis of Anhydrides (normal and cyclic)
- Dehydration of two carboxylic acids
- Intramolecular reaction of two carboxylic acids (cyclic)
Synthesis of Amides
- ammonia and anhydride
- ammonia and ester
Reactions of Amides
- acid catalyzed hydrolysis
- reduction to an amine with LiAlH4
Reactions of Esters
- transesterification (alcohol and ester)
- acid catalyzed hydrolysis to carb. acid
- reduction to alcohols with LAH
- saponification with strong base
Nitrogen containing compounds
- amide (carbonyl with NH2)
- imine (carbon db to nitrogen)
- enamine (alkene with NRR’ as one group)
- azide (RN-N+ triple bond N)
- nitrile (RC triple bond N)
- isocyanate (RN=C=O)
Strecker synthesis
Synth of amino acids with aldehyde, ammonium chloride, and potassium cyanide
Gabriel (Malonic Ester) Synthesis
Synth of amino acids with potassium phthalamide and diethyl bromomalonate
Phosphoric acid (bio app?)
A phosphate group aka Pi; at phys pH includes both HPO4 2- and H2PO4 -
Pyrophosphate (PPi)
P2O7 4-; released during formation of phosphodiester bonds in DNA; unstable in aqueous solution and is hydrolyzed to form two molecules of Pi
Organic Phosphates
Nucleotides w phosphate groups ie. ATP, GTP, those in DNA
Extraction
Separates dissolved substances based on differential solubility in aqueous vs. organic solvents
Filtration
Separates liquids from solids
Distillation
Separated liquids based on boiling point, which depends on IMF. Types are simple, fractional, and vacuum.
Simple distillation
Can be used to separate two liquids with boiling points below 150˚C and at least 25˚ apart.
Vacuum distillation
Should be used when a liquid to be distilled has a boiling point above 150˚C; to precent degradation of the product, the incident pressure is lowered, thereby lowering the boiling point.
Fractional distillation
Should be used when two liquids have boiling point less than 25˚C apart; by introducing a fractionation column, the sample boils and refluxes back down over a larger surface area, improving the purity of the distillate.
Recrystallization
Separates solids based on differential solubility in varying temperatures
Electrophoresis
Used to separate biological macromolecules based on size and/or charge
IR Spectroscopy
Measures molecular vibrations of characteristic functional groups.
UV Spectroscopy
Passing UV light through a chemical and plotting absorbance vs. wavelength; useful for studying compounds containing double bonds and heteroatoms with lone pairs.
TLC vs reverse phase
in TLC - polar plate
in reverse phase - nonpolar plate
Average velocity
v = ∆x/∆t (m/s)
Acceleration
The rate of change of an object’s velocity
a = ∆v/∆t (m/s^2)
Linear motion equations
v = vo + at
x = vot + 1/2at^2
v^2 = v^2o + 2ax
avg v = (vo + v)/2
x = avg v*t = (vo + v)/2 *t
Projectile motion (vertical and horizontal components of velocity)
vertical = vsinø horizontal = vcosø
Static friction
the force that must be overcome to set an object in motion
0 ≤ fs ≤ usN
Kinetic friction
opposes the forces of objects moving relative to each other
fk = ukN
Work
W = Fdcosø (if a force is perpendicular to the displacement, there is no work)
Joules (N/m)
System work
When the piston expands, work is done by the system (W >0)
When the piston compresses, work is done on the system (W<0)
The area under the Pvs.V curve is the amount of work done in a system.
Power
The rate at which work is performed; W/∆t (Watts, J/s)
Kinetic energy
1/2mv^2
Potential energy
mgh
Total mechanical energy
E = U + K
Work-energy theorem
When there are no nonconservative forces acting on a system, the total mechanical energy remains constant.
Newton’s first law
Objects will be stationary unless acted on by a force
Newton’s second law
When a net force is applied to a body of mass, m, the body will be accelerated in the same direction as the force.
F = ma (N = kg*m/s^2)
Newton’s third law
Every action has an equal and opposite reaction
Newton’s law of gravitation
all objects experience attraction to each other
Fg = Gm1m2/r^2
Uniform circular motion
Ac = v^2/r Fc = mv^2/r
Linear expansion
the increase in length by most solids when heated; “when temperature increased, the length of a solid increase a LoT”
∆L = aL∆T
Volume expansion
The increase in volume of fluids when heated
∆V = BV∆T
Conduction
heat transfer involving direct molecular collisions
Convection
heat transfer by physical motion of a liquid
Radiation
heat transfer by EM waves
Specific heat
Q = mc∆T; use to calculate Q when the object does not change phase
Heat of transformation
The quantity of hear required to change the phase of 1g of substance
Q = mL
First law of thermodynamics
∆U = Q - W
Second law of thermodynamics
Entropy of the system and the surroundings increases or remains unchanged
Specific gravity
psubstance/pwater
Density of water
10^3 kg/m^2 or 1 g/cm^2 or 1g/ml
Pressure
F/A (pascals = N/m^2)
Absolute pressure
P = Po + pgz
Gauge pressure
Pg = P-Patm
Continuity equation
A1V1 = A2V2
Bernoulli’s Equation
P + 1/2pv^2 + pgh = constant
Archimedes’ Principle
Fbuoy = pfluidgvsubmerged
Pascal’s Principle
A change in the pressure applied to an enclosed fluid is transmitted undiminished to every portion f the fluid and to the walls of the containing vessel
P = F1/A1 =A2/A2 and A1d1 = A2d2
so
W = F1d1 = F2d2
Coulomb’s law
kq1q2/r^2 = F
Electric field
Lines point toward negative
E = F/q = Q/r^2 (N/C or V/M)
Electric potential energy
U = q∆V = qEd = kQq/r (J)
Dipole
p = qd
The dipole feels no net translational force, but experiences a torque about the center causing it to rotate so that the dipole moment aligns with the electric field.
Electrical potential
The amount of work required to move a + test charge q from infinity to a particular point divided by the test charge
V = U/q (J/C)
Potential difference (voltage)
∆V= W/q = kQ/r
Current
the flow of electric charge
I = Q/∆t (A or C/s)
Ohm’s law
V = IR
Resistance
opposition to the flow of charge
R = resistivity*L/A (ohms Ω)
*resistance increases with increasing temperatures
Kirchoff’s laws
1) At any junction within a circuit, the sum of current flowing into that point must equal the sum of current leaving
2) the sum of voltage sources equals the sum of voltage drops around a closed-loop circuit
Resistors in series
Add
Resistors in parallel
1/Rt = 1/R1 + 1/R2… add as reciprocals and and take reciprocal of sum
Power dissipated by resistors
P = IV = V^2/R = I^2R
Capacitance
The ability to store charge per unit voltage
C = Q/V
Capacitors in parallel and series
Add in parallel, reciprocals in series
Energy stored by Capacitors
U = 1/2QV = 1/2CV^2 = 1/2Q^2/C
Wave formulas
f = 1/T; v = wavelength * f
Standing waves in strings
wavelength = 2L/n; f = nv/2L
Both ends are always nodes
Standing waves in open pipes
wavelength = 2L/n; f = nv/2L
Both ends are are antinodes
Standing waves in closed pipes
wavelength = 4L/n; f = nv/4L
The closed end of the pipe is always a node and open end is always an antinode
Sound propogation
Moves through deformable medium by the oscillation of particles parallel to the direction of the wave’s propogation
Intensity
I = P/A (W/m^2)
Sound level
ß = 10log(I/I0) (dB)
- increase of 10 dB = increase of intensity by factor of 10
- *increase of 20 dB = increase of intensity by factor of 100
Doppler effect
f’ = f(v± vD)/(v±vS)
Observer and detector moving closer: + in numerator, - sign in denominator
Observer and detector moving apart: - in numerator, + in denominator
Refraction
n = c/v
Speed of light
c = 3.00 * 10^8 m/s
Snell’s law
n1sinø1 = n2sinø2
when n2>n1, light bends toward the normal; when n1 > n2, light bends away from the normal
Diffraction
the bending of light around the corners of an obstacle or aperture into the region of geometrical shadow of the obstacle
dark fringes are located via
asinø = n*wabelength (n = 1, 2, 3..)
Optics eqution
1/o + 1/i = 1/f = 2/r
Concave mirrors
If an object is placed inside the focal length of a concave mirror instead, the image formed is behind the mirror, enlarged, and virtual
Convex mirrors
Regardless of the position of the object, a onvex mirror only forms virtual upright image
Converging lenses
Convex; for an object beyond the focal length the image is real and inverted; for an object inside the focal focal length, the image formed is virtual, upright and enlarged; no image if object is at the focal point
Diverging lenses
Concave; image is always virtual and between the object and lens
Magnification
-1/o
Photoelectric effect
E = hf = hc/wavelength
K = hf - W
K is the max kinetic energy of an ejected electron, W is the minimum energy required to eject an electron.
Mass defect
The difference between the sum of protons and neutrons and the atomic mass
Results from the conversion of matter to energy (E = mc^2) where energy is the binding energy that holds nucleons within the nucleus
Half-life
n = noe^-wavelength*t
Alpha decay
Release of an alpha particle He4/2
for new element - mass is -4, protons is -2
Beta-minus decay
Release of an electron
for new element - protons is +1
Beta -plus decay
Release of a positron
for new element - protons is -1
Log identites
Log A x Log B = Log A + Log B
Log A/B = Log A - Log B
Log A^B = B log A
Log 1/A = -log A
Hill’s criteria
Helps determine causal relationships; only temporality is necessary
