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General Chemistry Flashcards

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1
Q

Formal charge equation

A

Formal charge = V - N_nonbonding - 1/2*N_bonding

where V is the number of valence electrons

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2
Q

Exceptions to the octet rule

A

Incomplete octet: He, H, Li, Be, and B

Expanded octet: Any element in or beyond period three

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3
Q

Electron geometry

A

The spatial arrangement of all pairs of electrons around the central atom. Ex: H2O, NH3, and CH4 all have the same electron geometry because they all have 4 pairs of electrons surrounding the central atom

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4
Q

Molecular geometry

A

Describes the spatial arrangement of only the bonding pairs on electrons. Ex: water, ammonia, and methane differ in molec geo despite the same electron geometry. Water: bent/angular; ammonia: trigonal pyramidal; methane: tetrahedral

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5
Q

Coordination number

A

The number of atoms that surround and are bonded to the central atom. Relevant in determining molecular geometry

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6
Q

Dipole moment equation

A

p = q*d

where q is the magnitude of the charge and d is the displacement vector separating the partial charges

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7
Q

Coordinate Covalent Bond

A

Results when a single atom provides both bonding electrons while the other does not contribute any.

Ex: NH3 capturing H+ to form NH4+

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8
Q

Electrolytes

A

Contain equivalents of ions from molecules that dissociate in solution. Strength depends on its degree of dissociation or solvation

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9
Q

Name ending “-ous” vs “-ic”

A

-ous denotes lesser charge, while -ic denotes greater charge

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10
Q

Oxyanion name endings “-ite” vs “-ate”

A

-ite denotes compounds with lesser amount of oxygen than those with -ate ending

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11
Q

3-D arrays of charged particles when in solid form

A

Ionic compounds

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12
Q

Gram equivalent weight

A

Mass of a substance that can donate one equivalent of the species of interest

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13
Q

Neutralization reaction

A

Reaction of an acid and base to form a salt and water

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14
Q

Combustion Reactions

A

Occur when a fuel and an oxidant (typically oxygen) react to form water and CO2

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15
Q

Phosphate

A

PO4^3-

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16
Q

Phosphide

A

P^3-

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17
Q

Hypochlorite

A

ClO-

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18
Q

Ammonium

A

NH4+

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19
Q

Bicarbonate

A

HCO3-

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20
Q

Nitrite

A

NO2-

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21
Q

Homogenous catalyst

A

Catalyst that is the same phase as the reactants

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22
Q

Heterogenous catalyst

A

Catalyst that is a different phase than the reactants

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23
Q

For a second-order reaction, which graph has a linear slope?

A

1/[A] vs time. Slope = k

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24
Q

Slope of [reactant] vs time graph equals ____ for a zeroth order reaction

A

k

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25
Q

Which order reaction has a linear graph of ln([A]) vs time, and what is the value of the slope?

A

First-order; -k

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26
Q

Comparing reaction coordinate diagrams of different reactions under the same conditions, the reaction with the ____est ____ ___ will proceed the fastest, regardless of entropic change.

A

lowest activation energy

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27
Q

Endergonic vs Endothermic

A

Endergonic = nonspontaneous = energy of the products is greater than the energy of the reactants;

In contrast, endothermic means that heat had to be absorbed in order to achieve the reaction

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28
Q

Exergonic vs exothermic

A

Exergonic = energy of the products is lower than the energy of the reactants = sponaneous;

In contrast, exergonic simply means that heat was released from the system during the reaction

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29
Q

T/F: The presence of a catalyst affects K_eq of a reaction

A

False. A catalyst does not affect equilibrium position

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30
Q

Law of Mass Action

A

Expression for equilibrium constant, K_eq = [C]^c[D]^d / ([A]^a[B]^b)

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31
Q

Reaction Quotient

A

Same calculation as K_eq, but is calculated using the instantaneous concentrations of each species.

Q = [C]^c[D]^d / ([A]^a[B]^b)

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32
Q

Q>K_eq

A

There are more products, currently, than there would be at equilibrium. Reaction proceeds left to decrease amount of products

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33
Q

Q=K_eq

A

System is at dynamic equilibrium

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34
Q

Q

A

Currently less products than there would be at equilibrium. Reaction proceeds forwards

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35
Q

Which direction does the reaction proceed when pressure is increased on side A (arbitrary)? Decreased?

A

Increased: The reaction will proceed to the side with the fewest moles of gas

Decreased: The reaction will proceed to the side with the greater number of moles of gas

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36
Q

Kinetic products vs Thermodynamic products

A

Kinetic products are higher in free energy and less stable than thermodynamic products and can form at lower temperatures.

Kinetic pathway = FAST, higher energy intermediates

Thermodynamic pathway = slower, but more spontaneous, lower energy intermediates

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37
Q

Conditions that favor kinetic product over thermodynamic product?

And thermodynamic over kinetic?

A

Favor kinetic: lower temperatures, low heat transfer

Favor thermodynamic: higher temperature, high heat transfer

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38
Q

Combustion of ___-chain hydrocarbons yields greater energy

A

long

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39
Q

Triple point

A

The temperature and pressure at which all three phases exist in equilibrium

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40
Q

For a spontaneous forward reaction, Keq __ Q and deltaG _ 0

A

Keq must be GREATER than Q and deltaG<0 in order for a reaction to proceed forward spontaneously. Keq must be greater because a lower Q indicates that concentration or products are relatively low compared to equilibrium, so the reaction will proceed forward to correct this

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41
Q

Equation for entropy change

A

delta S = heat transfer / temperature = q/T

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42
Q

Standard free energy change equation

A

deltaG_0rxn = -RT*ln(K_eq)

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43
Q

Isolated systems

A

Neither matter nor heat can be exchanged with the environment

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44
Q

Closed systems

A

Matter cannot be exchanged with the environment, but heat can be

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45
Q

Adiabatic

A

Processes that exchange no heat with the environment

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46
Q

State functions

A

Path-independent functions that describe the physical properties of an equilibrium state. Ex: pressure, density, internal energy Gibbs free energy, and entropy

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47
Q

Standard state

A

The condition of an element in its most prevalent form under standard conditions (298 K, 1 atm, 1 M concentration)

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48
Q

Critical point

A

Temperature at which the liquid and gas phases of a substance are indistinguishable

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49
Q

Enthalpy

A

Measure of the potential energy of a system found in intermolecular attractions and chemical bonds

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50
Q

Hesse’s Law

A

Total change in potential energy of a system is equal to the changes of potential energies of the individual steps of the process

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51
Q

Entropy

A

Measure of the degree to which energy has been spread throughout a system or between a system and its surroundings. Unit = heat transferred / (mol*K)

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52
Q

When is entropy maximized?

A

At equilibrium

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53
Q

Equation for generalized enthalpy of a reaction

A

delta H = H_prod - H-reac

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54
Q

Equation for Gibbs free energy from reaction quotient

A

deltaG = deltaG_rxn + RT*lnQ

deltaG= RT*(Q/K_eq)

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55
Q

Equivalent units to 1 atm:

A

760 mm Hg = 760 torr = 101.3 kPa

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56
Q

Ideal Gas Law

A

PV = nRT

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57
Q

Avogadro’s Number

A

6.023 * 10^23 atoms / mol

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58
Q

Boyle’s Law, Charles’ Law, and Gay-Lussac Law

A

Boyle’s: P and V are inversely related
Charles’: Direct relationship between temp and volume
G-L: Direct relationship between temp and pressure

(when all other variables held constant)

Combine to form the “Combined Gas Law”

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59
Q

Dalton’s Law of Partial Pressure

A

partial pressure = mole fraction * total pressure

P_A = X_A * P_tot

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60
Q

Electromagnetic spectrum (order of waves)

A

Radio, microwaves, infrared, visible, ultraviolet, X-Rays, gamma

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61
Q

Visible range of wavelengths

A

400 nm (violet) to 700 nm (red)

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62
Q

Ideal gas constant (R)

A

.0821 Latm/(molK) = 8.312 J/(mol*K)

63
Q

Equation to show the relationship between density and molecular weight of an ideal gas

A

M = density_STP * 22.4 L/mol

64
Q

Kinetic Molecular Theory (state the assumptions)

A

Gases are made up of particles with volumes that are negligible compared to their container volume
Gas atoms or molecules exhibit no intermolecular attractions or repulsions
Gas particles are in continuous, random, motion, undergoing collisions with other particles and the container walls
Collisions between the particles and with the wall are elastic (conservation of momentum and energy)
Average KE of gas particles is proportional to the absolut temperature of the gas )in K) and it is the same for all gases at given temp

65
Q

Equation for the KE of an ideal gas

A

KE = 3/2 * k_B T = 1/2m*v^2

k_B is Boltzmann’s constant: 1.38*10^-23 J/K

66
Q

Root-mean-square speed

A

Defines average speed of gases in terms of average molecular speed.

u_RMS = sqrt(3RT/M)

67
Q

Graham’s Law

A

r1/r2 = sqrt(M2/M1)

Shows that gases with smaller particle diffuse faster

68
Q

Effusion

A

The flow of gas particles under pressure from one compartment to another through a small opening

69
Q

Van der Waals Equation of State

A

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

a indicates the attractions between the molecules of the nonideal gas
b indicates the size of the molecules

70
Q

Amphoteric Species

A

Molecules that can act as both acids and bases. Acts acidic in an acidic environment and acts as a base in a basic one

71
Q

Nomenclature of oxyacids

A

Compounds with fewer oxygens in them: -ite ending in the anion form, -ous ending in the protonated form

Compounds with more oxygens in them: -ate ending in the anion form, -ic acid in the protonated form

72
Q

Amphiprotic

A

Able to gain or donate an H+. Special class of amphoteric species

73
Q

Water dissociation constant

A

K_w = [H3O+]*[OH-] = 1e-14

74
Q

pH equation

A

= -log ([H+]) = log( 1/[H+] )

same for pOH but with OH- in place of H+

75
Q

Based on strength of acids and bases, why is the bicarbonate an ideal buffering system?

A

CO3^2- is a weak base and its conjugate acid HCO3- is a weak acid, too.

76
Q

Hydrolysis of a salt

A

Adding water to a salt to result in the reformation of the acid and base that reacted to form the salt and water in the first place.

77
Q

Strong acid + strong base

A

= salt + water

This is a neutralization reaction

78
Q

Strong acid + weak base

A

Forms a weak acid, but no water because the base is not a hydroxide. Some of the original weak base is formed, which also results in the formation of hydronium, making the result slightly acidic

79
Q

Strong base + weak acid

A

Formation of a salt, the basic ion of which then deprotonates water to form OH- and make to solution basic

80
Q

Weak acid + weak base

A

The acidity of the resulting solution depends on the relative strengths of the reactants. If the Ka of the acid is greater than the Kb of the base, then the result will be acidic

81
Q

Relationship between Ka, Kb, and Kw

A

Ka * Kb = Kw = 1e-14 (at 298 K)

82
Q

Polyvalent species

A

Acids or bases that, for each mole, liberate more than one mole of acid or base equivalent. Under B-L definition, they are polyprotic

83
Q

Gram equivalent weight

A

The mass of a compound that yields one equivalent (one mole of charge)

84
Q

Titrant

A

The species of known pH that, in very small volumes, gets added to a larger solution

85
Q

Titrand

A

The solution of unknown concentration that gets acid or base added to it until reaching the equivalence point

86
Q

Equivalence point in titrations

A

The point at which the number of acid equivalents initially present in the titrand is equal to the number of equivalents of base that have been added. This does not always occur at pH=7, although it is at 7 when either a strong acid or strong base is titrating the strong version of the other.

87
Q

What can the following equation be used for?

Na*Va = Nb * Vb

A

Calculation of the concentration of a titrand from the volume of base added, the volume of original acid, and the normality of the base added.

Shows that the equivalents of acid and base are equal at the equivalence point

88
Q

Buffers

A

Composed of a weak species and a salt of its conjugate

89
Q

Consider the buffer of CH3COOH + NaCH3COOH.

Describe how the system mitigates addition of OH-

A

As OH- is added, H+ from CH3COOH react with it and form water. As a result, more CH3COO- is in solution, but this does not cause the solution to become nearly as basic as the initial addition of OH- would have.

90
Q

Hendersen-Hasselbalch equation (in both forms: pH and pOH)

A

pH = pKa + log(A-/HA)

pOH = pKb + log (B+/BOH)

91
Q

Indicators

A

Weak organic acids or bases that change color when they become protonated or deprotonated. They change color at the endpoint

92
Q

Buffering capacity

A

The ability of a system to resist changes in pH

93
Q

Anode

A

The electrode in an electrochemical cell at which oxidation occurs.

Think: an ox

94
Q

Cathode

A

The electrode in an electrochemical cell at this reduction occurs.

Think: red cat

95
Q

Electromotive force (what corresponds with positive and negative?)

A

The voltage or electrical potential of the cell.

Positive emf: negative delta G, so the reaction is spontaneous

Negative emf: positive delta G, so the cell must absorb energy to cause the reaction (nonspontaneous)

96
Q

Movement of electrons in electrochemical cells

A

From the anode to the cathode

97
Q

Galvanic cells

A

Also called voltaic cells. Always spontaneous (energy is released to the environment, emf is positive). All non-rechargeable batteries are voltaic cells. Electric potential energy gets converted to kinetic in the form of electron movement as the system approaches equilibrium.

98
Q

Salt bridge

A

Permits the flow of anions and cations. Contains an inert electrode that will not react with the electrodes or ions in solution (usually KCl or NH4NO3).

Cations flow towards the cathode to make up for the positive charge lost by the reduction of a charged metal ion to its atom.

Anions flow towards the anode to neutralize the positive charges introduced by the oxidation of a metal atom to its ion.

99
Q

Galvanization

A

This refers to the precipitation process of the cathode itself. Also called plating

100
Q

Cell diagram

A

Shorthand notation representing the reactions in an electrochemical cell:

anode (s) | anode solution || cathode solution | cathode (s)

The double line represents a salt bridge or some other type of barrier

101
Q

Electrolytic cells

A

Non-spontaneous redox reactions occur here, thus requiring the input of energy (electrolysis). Achieves chemical decomposition

102
Q

Electrolysis

A

Redox reactions drive by external voltage source

103
Q

How and why to obtain pure Na+ and pure Cl-:

A

How: electrolysis of molten NaCl into Cl(g) and Na(l)

Why: these elements do not occur naturally in their elemental forms because they are so reactive

104
Q

Charge per electron

A

-1.6 * 10^-19 C

105
Q

Faraday constant

A

Equal to the amount of charge contained in one mole of e’

F = 96,485 C (rounds to 10^5 C/mol e-)

106
Q

Electrodeposition equation

A

Determines the number of moles of element being deposited on a plate or the amount of gas that is liberated by electrolysis:

mol M = It/nF

107
Q

Concentration cell

A

Allows a spontaneous redox reaction to occur between two half cells, but the electrodes are chemically identical. Current is generated by an essential concentration gradient. Current stops flowing when the concentrations in the half cells are equal. Voltage is calculated using the Nernst equation

108
Q

A rechargeable cell/battery can function as both a(n) ____ cell and a(n) _____ cell

A

Galvanic; electrolytic

109
Q

Lead storage battery

A

A voltaic cell with a Pb anode and a porous PbO2 cathode connected by a conductive material (concentrated with 4 M H2SO4). Both reactions at the anode and cathode result in plating with lead sulfate and dilute the acid electrolyte when discharging. These kinds of cells have some of the lowest energy-to-weight ratios, which means that A LOT of battery material to produce a certain output compared to a other batteries

110
Q

Energy density

A

A measure of a battery’s ability to produce power as a function of its weight

111
Q

Nickel-Cadmium batteries

A

Rechargeable cells.

Anode: made of Cd(s)
Cathode: Ni(III)OH_3
Connected by KOH or other conductive material

Must be vented to prevet build up of H and O gas during electrolysis. Have higher energy density than Pb-acid batteries, and tend to provide higher surge currents

112
Q

Surge currents

A

Periods of large currents (amperage). Preferable in appliances like remote controls that demand rapid responses.

113
Q

In a voltaic cell, the anode is considered _____ in charge because it is the _____ of electrons

A

negative; source

114
Q

In a galvanic cell,

Electrons move from ___ electric potential to ____ electric potential, while current (the flow of positive charge) flows from ___ to ___.

A

Electrons: Low to high;

Current: High to low

115
Q

Which is the positive half-cell in an electrolytic cell? Why?

A

The anode is considered positive because it is attached to the positive pole of the external voltage source and attracts anions from the solution.

116
Q

Which is the negative half-cell in an electrolytic cell? Explain

A

The cathode is considered the negative half cell because it is attached to the negative end of the battery/external power source and attracts cations from the solution.

117
Q

The anode always attracts ____ and the cathode always attracts _____.

A

Anode: anions
Cathode: cations

118
Q

Isoelectric focusing

A

A technique used to separate amino acids or polypeptides based on their isoelectric points. The positively charged amino acids will migrate towards the cathode and negatively charged ones toward the anode.

119
Q

Ni-Cd cells

A

A rechargeable battery.

When discharging: consists of Cd anode and NiO(OH) cathode in concentrated KOH solution.

When charging: The Ni(OH)2 and Cd(OH)2 plated electrodes are dissociated to restore the original Cd and NiI(OH) electrodes and concentrate the electrolyte.

Higher energy density than lead-acid batteries

120
Q

Nickel-metal hydride (NiMH) batteries

A

Had basically replaced Ni-Cd batteries because they have higher energy density, are more cost-effective, and slightly less toxic

121
Q

When is a rechargeable battery acting electrolytically?

Describe the flow of electrons and current

A

During charging.

Electrons flow from high electric potential to low electric potential. Current flows from low to high.

122
Q

Reduction potential

A

Quantifies the tendency for a species to gain electrons and be reduced. A species wants to be more and more reduced the higher the reduction potential

123
Q

Standard reduction potentials

A

(E*_red) - They are calculated by comparison to the standard hydrogen electrode (SHE) under the standard conditions

124
Q

Standard hydrogen electrode

A

Reduction potential of 0 V

125
Q

Standard emf (def and equation)

A

The difference in standard reduction potential between the two half-cells

E_cell = E_red - E_oxid

126
Q

Sign of E_cell and delta G in galvanic cells

A

E_cell: positive

delta G: negative

127
Q

Sign of E_cell and delta G in electrolytic cells

A

E_cell: negative (because E_cat

128
Q

Dehydrogenases

A

Enzymes that catalyze oxidation in which the electrons lost are in the form of Hydrogen

129
Q

Changes in the oxidation numbers of transition metals usually correspond with changes in _______.

A

color

130
Q

Net ionic equations show only….

A

The ions that participate in the redox reaction

131
Q

Metathesis reactions

A

Reactions that involve the switching of counterions. AKA double-displacement reactions. They are NOT usually redox reactions

132
Q

Disproportionation reaction

A

AKA dismutation. A specific type fo redox reaction in which an element undergoes both reduction and oxidation to produce its products.

Ex: catalysis of peroxides in the peroxisome by catalase. It catalyzes the reaction => 2 H2O2 –> 2 H2O + O2
in which oxygen in the reactant has the oxidation state -1, and it is then both oxidized and reduced

133
Q

Catalase

A

An enzyme found in the peroxisome that is used to protect cells from excessive oxidation by free radicals or reactive oxygen species using the following disproportionation reaction:

2 H2O2 –> 2 H2O + O2

134
Q

Superoxide dismutase

A

An mutase in the human body that disproportions oxygen free radicals using a following disproportionation mechanism in which radical oxygen is both oxidized to elemental O2 and reduced to O in hydrogen peroxide

135
Q

Iodometric titration

A

Relies on the titration of free iodine radicals

136
Q

Dominant species when the pH of a solution is greater than pKa:

A

Due to the H-H equation, when pH is greater than pKa, this indicates that log(A-/HA)>1, so A->HA. This means that the compound is mainly dissociated

137
Q

T/F: Cooperative enzymes have one active site, regardless of allosteric regulation

A

False - if they engage in cooperative binding, this means that they MUST have more than one active site

138
Q

Diamagnetic atoms/ions

A

Characterized by the electron configuration having only complete pairs. Thus, the electrons will be repelled by a magnetic field

139
Q

Paramagnetic atoms/ions

A

Unpaired electrons are present, so they will be attracted to an external magnetic field

140
Q

Dilution Factor

A

DF = V_T1 / V_F1

T1: volume transferred
F1: final volume

If more than one dilution, then multiply dilution factors

141
Q

T/F: Mass changes due to radioactive decay are negligible

A

True!

142
Q

Electron capture

A

Form of radioactive decay in which the nucleus “captures” an electron, turning a proton into a neutron. Mass number remains, but atomic number decreases by 1

143
Q

Hydration/Solvation

A

The attractive force of an ion-molecule causing a thin shell of water molecules to surround it

144
Q

Describe the change in entropy when water is forced to interact with nonpolar substances

A

Water molecules are unable to interact with such hydrophobic molecules, so the water molecules must form a highly ordered solvation shell around them in order to minimize interaction with those reisdues. This high-order causes a decrease in entropy, therefore is it energetically unfavorable!

145
Q

Common Ion Effect

A

Pre-existing presence of an ion in solution reduces the molar solubility of a substance containing that ion. The same amount of solute being added to a previously mixed solution of that solute+solvent will result in more precipitate/closer to saturation than if it were added to a solution of pure water.

146
Q

Describe the location of equivalence points and pKas on a titration curve

A

Equivalence point: the center of the steep portion of the graph

pKa point: approximately the center of the more horizontal buffered region

147
Q

Pyranose

A

6-membered cyclic sugar with 5 C’s and 1 O

148
Q

Furanose

A

5-membered cyclic sugar with 4 Cs and 1O

149
Q

Mass Spectrometry

A

Ionizes molecules and detects their molecular weight - This is done by accelerating ions toward a magnet, which then deflect according to mass. Molecular ions can be broken into fragments, each of which can be analyzed by another round of MS. The ions are then detected and plotted in mass abundance vs the mass-to-charge ratio (m/z). Mass spec can be used to identify the mass of a molecule’s fragments by taking the difference between m/z peaks.

150
Q

Reaction of LiAlH4 and an ester

A

Produces a primary alcohol

151
Q

Equation for specific rotation

A

a_specific = a_obs / (concentration*path length)

a_spec = a_obs / (c*l)

152
Q

Groups of the Representative Elements

A

Groups 1 and 2, and 13-18

153
Q

Trend in first ionization energy

A

Increases going to the right across a period (so reactivity decreases)

Decreases going down a group (so reactivity increases)

MCAT (6 decks)

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