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Chem/Phys Flashcards

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1
Q

Atomic Weight

A

The weighted average of the masses of the naturally occurring isotopes of an element, in amu per molecule or grams per mole

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2
Q

Mole

A

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.

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3
Q

Isotopes

A

For a given element, multiple species of atoms with the same number of protons but different numbers of neutrons

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4
Q

Planck’s quantum theory

A

Energy emitted as electro-magnetic radiation from matter exists in discrete bundles called quanta

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5
Q

Bohr’s Model of H Atom

A

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)

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6
Q

Balmer series vs Lyman series

A

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

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7
Q

Absorption spectra

A

Characteristic energy bands where electrons absorb energy

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8
Q

Heisenberg Uncertainty Principle

A

Impossible to know the location and momentum of an electron at the same time

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9
Q

Quantum numbers

A

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

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10
Q

Hund’s Rule

A

Within a given subshell, orbitals are filled such that there are a maximum number of half filled orbitals with parallel spins

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11
Q

Valence electrons

A

Electrons of an atom that are in its outer energy shell and available for bonding

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12
Q

Coordination Compounds

A

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.

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13
Q

Chelation

A

When the central cation in a coordination compound is bound to the same ligand multiple times

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14
Q

Hydrogen bonding

A

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

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15
Q

Dipole-dipole Interactions

A

Polar molecules orient themselves such that the positive region of one molecule is close to the negative region of another molecule.

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16
Q

Dispersion forces

A

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

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17
Q

Units for rate constant?

A

zero-order - M/s
first-order - s^-1
second-order - M^-1s^-1
third-order - M^-2s^-1

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18
Q

Combustion reaction

A

A fuel, such as a hydrocarbon, is reacted with an oxidant, such as oxygen, to produce an oxide and water

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19
Q

Combination reaction

A

Two or more reactants form one product

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20
Q

Decomposition reaction

A

A compound breaks down into two or more substances, usually as a result of heating or electrolysis

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21
Q

Single-displacement reaction

A

An atom (or ion) of one compound is replaced by an atom of another element

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22
Q

Double displacement reaction

A

Also called metathesis reactions; elements from two different compounds displace each other to form two new compounds

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23
Q

Net ionic equation

A

Written showing only species that actually participate in the reaction

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24
Q

Neutralization reaction

A

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

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25
Q

Factors that affect reaction rates?

A

Reactant concentrations, temperature, medium, catalysts

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26
Q

Law of Mass Action

A

aA + bB –> cC + dD

Keq = [C]^c[D]^d/[A]^a[B]^b

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27
Q

Properties of equilibrium constant

A

Pure solids and liquids don’t appear in the expression

If Keq&raquo_space; 1, little of the reactants compared to the products.
If Keq &laquo_space;1, little of the products
compared to the reactants.
If Keq is close to 1, approximately equal amounts of the two

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28
Q

Le Chatelier’s Principle

A

Used to determine the direction of the reaction at equilibrium when subjected to stress (ie. change in concentrations, pressure, volume, or temp)

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29
Q

Isolated vs closed vs open systems

A

Isolated - no exchange of energy or matter
Closed - no exchange of matter but exchange of energy
Open - exchange of matter and energy

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30
Q

Isothermal Process

A

Constant temperature; ∆U = 0; first law is Q = W

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31
Q

Adiabatic process

A

Constant heat; Q = 0, first law is ∆U = -W

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32
Q

Isobaric process

A

Constant pressure

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33
Q

Isovolumetric (Isochoric) process

A

Constant volume; W = Q, first law is ∆U = Q

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34
Q

Endothermic vs Exothermic

A

Absorb energy (positive ∆H) vs Release energy (negative ∆H)

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35
Q

Heat absorbed or released in a given process

A

q = mc∆T (c - specific heat)

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36
Q

State functions

A

Describe the macroscopic properties of the system; pressure, density, temperature, volume, enthalpy, internal energy, free energy, and entropy

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37
Q

Enthalpy

A

Used to express heat changes at constant pressure

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38
Q

Standard heat of formation

A

The enthalpy change that would occur if one mole of a compound was formed directly from its elements in their standard states

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39
Q

Standard heat of reaction

A

the hypothetical enthalpy change that would occur if the reaction were carried out under standard conditions

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40
Q

Hess’s Law

A

the enthalpies of reactions are additive; the reverse of any reaction has the same magnitude with opposite sign

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41
Q

Bond dissociation energy

A

The average energy required to break a particular type of bond in one mole of gaseous molecules

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42
Q

Bond enthalpy

A

The standard heat of reaction can be calculated using the values of bond dissociation energies of particular bonds

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43
Q

Entropy

A

The measure of the distribution of energy (“randomness”) throughout a system

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44
Q

Reaction spontaneity by ∆H and ∆S signs

A

-/+ : spont at all temps
+/- : nonspont at all temps
+/+ : spont at high temps
-/- : spont at low temps

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45
Q

Reaction quotient

A

Once a reaction commences, the standard state conditions no longer hold; Q is the same equation as K

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46
Q

Pressure Equivalents

A

1 atm = 760 torr = 760 mmHg = 101,325 Pa

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47
Q

STP vs Standard Conditions

A

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

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48
Q

Boyle’s Law

A

P1V1 = P2V2

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49
Q

Charle’s Law

A

V1/T1 = V2/T2

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50
Q

Gay Lussac’s Law

A

P1/T1 = P2/T2

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51
Q

When do ideal gases behave unideally?

A

Small volume; high pressure or low temperature

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52
Q

Van der Waals equation of state

A

accounts for deviations from ideality that occur

(P + n^2a/V^2)(V-nb) = nRT

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53
Q

1 mol of Gas at STP is how many liters?

A

22.4

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54
Q

Dalton’s law of partial pressures

A

The total pressure of a gaseous mixture is equal to the sum of the partial pressures of the individual components

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55
Q

Calculate partial pressure of a gas

A

Pgas = Xgas*Ptotal

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56
Q

Kinetic molecular theory of gases

A

An explanation of gaseous molecular behavior based on the motion of individual molecules

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57
Q

Average molecular speeds

A

K = 1/2mv^2 = 3/2kBT

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58
Q

Root-mean-square speed

A

urms = √3RT/M

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59
Q

Colligative properties

A

Properties derived solely from the number of particles present, not the nature of the particles

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60
Q

Freezing point depression

A

∆Tf = iKfm (m is molality)

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61
Q

Boiling point elevation

A

∆Tb = iKbm (m is molality)

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62
Q

Osmotic pressure

A

π = MRT

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63
Q

Raoult’s Law

A

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

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64
Q

Graham’s law of diffusion and effusion

A

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

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65
Q

Solubility rules

A

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.

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66
Q

Units of concentration

A 
percent comp. by mass: mass solute/mass solution *100
mole fraction: moles solute/total moles
molarity: moles/liter
molality: moles/kg
normality: g/L
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67
Q

Arheinius definition of acids and bases

A

Acids as source of H+ and bases as source of OH-

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68
Q

Brønsted-Lowry vs Lewis

A

Proton donors/acceptors vs Electron donors/acceptors

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69
Q

pH and pOH

A

-log[H+] or -log[OH-]

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70
Q

Kw

A

10^-14; pH + pOH = 14

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71
Q

Properties of Weak acids and Bases

A

HA (aq) + H2O (l) –> H3O+ (aq) + A- (aq)

Ka = [H3O+][A-]/[HA]

Kb = [B+][OH-]/[BOH]

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72
Q

Salt formation

A

Acids and bases may react with each other, forming a salt and (often) water in a neutralization reaction

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73
Q

Hydrolysis

A

The reverse of a neutralization, where salt reacts with water to produce the acid and base

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74
Q

Amphoteric species

A

Can act as either an acid or a base depending on chemical environment

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75
Q

Titration

A

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

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76
Q

Henderson-Hasselbach Equation

A

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]

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77
Q

Oxidation vs reduction

A

loss of electrons vs gain of electrons

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78
Q

Oxidizing agent vs reducing agent

A

is reduced vs is oxidized

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79
Q

Galvanic cells

A

∆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

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80
Q

Electrolytic cells

A

Nonspontaneous –> electrical energy is required to induce a reaction; everything in one container.

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81
Q

Reduction potential

A

The tendency of a species to acquire electrons and be reduced; standard reduction potentials calculated under standard conditions (25˚C, 1 atm, 1 M)

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82
Q

E˚cell or emf

A

emf = E˚red,cathode - E˚red, anode

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83
Q

Spontaneity of electrochemistry

A

∆G = -nFEcell

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84
Q

Sn1

A

2 steps; likes polar protic solvents; 3˚ > 2˚ > 1˚ > methyl; rate = k[RL]; racemic products; strong nucleophile not required

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85
Q

Sn2

A

1 step; likes polar aprotic solvents; methyl > 1˚ > 2˚ > 3˚; rate = k[Nu][RL]; optically active and inverted products; favored with strong nucleophile

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86
Q

Nucleophilicity in protic solvents:

A

F- > Cl- > Br - > I- bc protic solvents can inhibit by protonation or H bonding

87
Q

Nucleophilicity in aprotic solvents:

A

I- > Br- > Cl- > F-

88
Q

Physical properties

A

Characteristics of processes that don’t change the composition of matter (ie. mp, bp, solubility, odor, color, and density)

89
Q

Chemical properties

A

Have to do with the reactivity of the molecule with other molecules.

90
Q

Conformational isomers

A

Different Newman projections; strain maxed in eclipsed conformations

91
Q

Cyclic strain types

A

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

92
Q

Configurational isomers

A

Can only be interchanged by breaking and reforming bonds; ie. enantiomers and diastereomers

93
Q

Properties of alcohols

A

Higher BP than alkanes; weakly acidic hydroxyl hydrogen

94
Q

Synthesis of alcohols

A
  • 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)
95
Q

Reactions of alcohols

A

Substitution reactions after protonation and leaving group (tosylate) conversion

96
Q

Levels of organic oxidation-reduction

A

0) alkanes
1) alcohols, alkyl halides, amines
2) aldehydes, ketones, imines
3) carboxylic acids, anhydrides, esters, amides
4) carbon dioxide

97
Q

Organic redox definitions of oxidation and reduction

A

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

98
Q

Organic oxidizing agents

A

Have high affinity for electrons (O2, O3, Cl2) or unusually high oxidation staes (Mn7+ in MnO4- or Cr6+ in CrO4 2-)

99
Q

Organic reducing agents

A

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)

100
Q

Strong and weak oxidizers

A

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

101
Q

Alcohols as protecting groups

A

Can be used as protecting groups for carbonyls –> rxn with a dialcohol forms an acetal which can be removed with aqueous acid

102
Q

Phenol acidity

A

Hydrogen of alcohol is particularly acidic because the anion is stabilized by resonance with the ring

103
Q

Quinone synth? Hydroxyquinone importance/synth?

A

Oxidation of phenol (cyclohexane w vertical DBs and carbonyls at bottom and top corner)

further oxidation –> hydroxyquinone which have bio activity

104
Q

Ubiquinone

A

Also called coenzyme Q, a vital electron carrier associated with complexes I, II, and III of the ETC

105
Q

Q cycle

A

Reduction of ubiquinone to ubiquinol, rinse and repeat. In the ETC

106
Q

Aldehyde properties

A

Higher bp than alkanes due to polarity, but not as high as alcohols because no H-bonding

107
Q

Synthesis of aldehydes

A
  • Oxidation of primary alcohols

- Ozonolysis of alkenes

108
Q

Reactions of Aldehydes

A

-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

109
Q

Carboxylic acid properties

A

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

110
Q

Synthesis of Carboxylic acids

A
  • oxidation of primary alcohols with KMnO4

- hydrolysis of nitriles

111
Q

Reactions of Carboxylic acids

A
  • formation of soap by reacting carb. acids with NaOH, arrange in micelles
  • nucleophilic acyl substitution (incl. ester formation and reduction to alcohols)
  • decarboxylation
112
Q

Lactam

A

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)

113
Q

Lactone

A

Cyclic esters; named by carbon bound to oxygen and by length of chain; ie. alpha-acetolactone, beta-propiolactone

114
Q

Reactivity of Carb. Acid Derivatives

A

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

115
Q

Synthesis of Anhydrides (normal and cyclic)

A
  • Dehydration of two carboxylic acids

- Intramolecular reaction of two carboxylic acids (cyclic)

116
Q

Synthesis of Amides

A
  • ammonia and anhydride

- ammonia and ester

117
Q

Reactions of Amides

A
  • acid catalyzed hydrolysis

- reduction to an amine with LiAlH4

118
Q

Reactions of Esters

A
  • transesterification (alcohol and ester)
  • acid catalyzed hydrolysis to carb. acid
  • reduction to alcohols with LAH
  • saponification with strong base
119
Q

Nitrogen containing compounds

A
  • 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)
120
Q

Strecker synthesis

A

Synth of amino acids with aldehyde, ammonium chloride, and potassium cyanide

121
Q

Gabriel (Malonic Ester) Synthesis

A

Synth of amino acids with potassium phthalamide and diethyl bromomalonate

122
Q

Phosphoric acid (bio app?)

A

A phosphate group aka Pi; at phys pH includes both HPO4 2- and H2PO4 -

123
Q

Pyrophosphate (PPi)

A

P2O7 4-; released during formation of phosphodiester bonds in DNA; unstable in aqueous solution and is hydrolyzed to form two molecules of Pi

124
Q

Organic Phosphates

A

Nucleotides w phosphate groups ie. ATP, GTP, those in DNA

125
Q

Extraction

A

Separates dissolved substances based on differential solubility in aqueous vs. organic solvents

126
Q

Filtration

A

Separates liquids from solids

127
Q

Distillation

A

Separated liquids based on boiling point, which depends on IMF. Types are simple, fractional, and vacuum.

128
Q

Simple distillation

A

Can be used to separate two liquids with boiling points below 150˚C and at least 25˚ apart.

129
Q

Vacuum distillation

A

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.

130
Q

Fractional distillation

A

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.

131
Q

Recrystallization

A

Separates solids based on differential solubility in varying temperatures

132
Q

Electrophoresis

A

Used to separate biological macromolecules based on size and/or charge

133
Q

IR Spectroscopy

A

Measures molecular vibrations of characteristic functional groups.

134
Q

UV Spectroscopy

A

Passing UV light through a chemical and plotting absorbance vs. wavelength; useful for studying compounds containing double bonds and heteroatoms with lone pairs.

135
Q

TLC vs reverse phase

A

in TLC - polar plate

in reverse phase - nonpolar plate

136
Q

Average velocity

A

v = ∆x/∆t (m/s)

137
Q

Acceleration

A

The rate of change of an object’s velocity

a = ∆v/∆t (m/s^2)

138
Q

Linear motion equations

A

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

139
Q

Projectile motion (vertical and horizontal components of velocity)

A 
vertical = vsinø 
horizontal = vcosø
140
Q

Static friction

A

the force that must be overcome to set an object in motion

0 ≤ fs ≤ usN

141
Q

Kinetic friction

A

opposes the forces of objects moving relative to each other

fk = ukN

142
Q

Work

A

W = Fdcosø (if a force is perpendicular to the displacement, there is no work)

Joules (N/m)

143
Q

System work

A

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.

144
Q

Power

A

The rate at which work is performed; W/∆t (Watts, J/s)

145
Q

Kinetic energy

A

1/2mv^2

146
Q

Potential energy

A

mgh

147
Q

Total mechanical energy

A

E = U + K

148
Q

Work-energy theorem

A

When there are no nonconservative forces acting on a system, the total mechanical energy remains constant.

149
Q

Newton’s first law

A

Objects will be stationary unless acted on by a force

150
Q

Newton’s second law

A

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)

151
Q

Newton’s third law

A

Every action has an equal and opposite reaction

152
Q

Newton’s law of gravitation

A

all objects experience attraction to each other

Fg = Gm1m2/r^2

153
Q

Uniform circular motion

A 
Ac = v^2/r
Fc = mv^2/r
154
Q

Linear expansion

A

the increase in length by most solids when heated; “when temperature increased, the length of a solid increase a LoT”

∆L = aL∆T

155
Q

Volume expansion

A

The increase in volume of fluids when heated

∆V = BV∆T

156
Q

Conduction

A

heat transfer involving direct molecular collisions

157
Q

Convection

A

heat transfer by physical motion of a liquid

158
Q

Radiation

A

heat transfer by EM waves

159
Q

Specific heat

A

Q = mc∆T; use to calculate Q when the object does not change phase

160
Q

Heat of transformation

A

The quantity of hear required to change the phase of 1g of substance

Q = mL

161
Q

First law of thermodynamics

A

∆U = Q - W

162
Q

Second law of thermodynamics

A

Entropy of the system and the surroundings increases or remains unchanged

163
Q

Specific gravity

A

psubstance/pwater

164
Q

Density of water

A

10^3 kg/m^2 or 1 g/cm^2 or 1g/ml

165
Q

Pressure

A

F/A (pascals = N/m^2)

166
Q

Absolute pressure

A

P = Po + pgz

167
Q

Gauge pressure

A

Pg = P-Patm

168
Q

Continuity equation

A

A1V1 = A2V2

169
Q

Bernoulli’s Equation

A

P + 1/2pv^2 + pgh = constant

170
Q

Archimedes’ Principle

A

Fbuoy = pfluidgvsubmerged

171
Q

Pascal’s Principle

A

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

172
Q

Coulomb’s law

A

kq1q2/r^2 = F

173
Q

Electric field

A

Lines point toward negative

E = F/q = Q/r^2 (N/C or V/M)

174
Q

Electric potential energy

A

U = q∆V = qEd = kQq/r (J)

175
Q

Dipole

A

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.

176
Q

Electrical potential

A

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)

177
Q

Potential difference (voltage)

A

∆V= W/q = kQ/r

178
Q

Current

A

the flow of electric charge

I = Q/∆t (A or C/s)

179
Q

Ohm’s law

A

V = IR

180
Q

Resistance

A

opposition to the flow of charge

R = resistivity*L/A (ohms Ω)

*resistance increases with increasing temperatures

181
Q

Kirchoff’s laws

A

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

182
Q

Resistors in series

A

Add

183
Q

Resistors in parallel

A

1/Rt = 1/R1 + 1/R2… add as reciprocals and and take reciprocal of sum

184
Q

Power dissipated by resistors

A

P = IV = V^2/R = I^2R

185
Q

Capacitance

A

The ability to store charge per unit voltage

C = Q/V

186
Q

Capacitors in parallel and series

A

Add in parallel, reciprocals in series

187
Q

Energy stored by Capacitors

A

U = 1/2QV = 1/2CV^2 = 1/2Q^2/C

188
Q

Wave formulas

A

f = 1/T; v = wavelength * f

189
Q

Standing waves in strings

A

wavelength = 2L/n; f = nv/2L

Both ends are always nodes

190
Q

Standing waves in open pipes

A

wavelength = 2L/n; f = nv/2L

Both ends are are antinodes

191
Q

Standing waves in closed pipes

A

wavelength = 4L/n; f = nv/4L

The closed end of the pipe is always a node and open end is always an antinode

192
Q

Sound propogation

A

Moves through deformable medium by the oscillation of particles parallel to the direction of the wave’s propogation

193
Q

Intensity

A

I = P/A (W/m^2)

194
Q

Sound level

A

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
195
Q

Doppler effect

A

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

196
Q

Refraction

A

n = c/v

197
Q

Speed of light

A

c = 3.00 * 10^8 m/s

198
Q

Snell’s law

A

n1sinø1 = n2sinø2

when n2>n1, light bends toward the normal; when n1 > n2, light bends away from the normal

199
Q

Diffraction

A

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..)

200
Q

Optics eqution

A

1/o + 1/i = 1/f = 2/r

201
Q

Concave mirrors

A

If an object is placed inside the focal length of a concave mirror instead, the image formed is behind the mirror, enlarged, and virtual

202
Q

Convex mirrors

A

Regardless of the position of the object, a onvex mirror only forms virtual upright image

203
Q

Converging lenses

A

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

204
Q

Diverging lenses

A

Concave; image is always virtual and between the object and lens

205
Q

Magnification

A

-1/o

206
Q

Photoelectric effect

A

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.

207
Q

Mass defect

A

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

208
Q

Half-life

A

n = noe^-wavelength*t

209
Q

Alpha decay

A

Release of an alpha particle He4/2

for new element - mass is -4, protons is -2

210
Q

Beta-minus decay

A

Release of an electron

for new element - protons is +1

211
Q

Beta -plus decay

A

Release of a positron

for new element - protons is -1

212
Q

Log identites

A

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

213
Q

Hill’s criteria

A

Helps determine causal relationships; only temporality is necessary

MCAT (5 decks)

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