General Chemistry Flashcards
(153 cards)
1
Q
Formal charge equation
A
Formal charge = V - N_nonbonding - 1/2*N_bonding
where V is the number of valence electrons
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
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
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
5
Q
Coordination number
A
The number of atoms that surround and are bonded to the central atom. Relevant in determining molecular geometry
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
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+
8
Q
Electrolytes
A
Contain equivalents of ions from molecules that dissociate in solution. Strength depends on its degree of dissociation or solvation
9
Q
Name ending “-ous” vs “-ic”
A
-ous denotes lesser charge, while -ic denotes greater charge
10
Q
Oxyanion name endings “-ite” vs “-ate”
A
-ite denotes compounds with lesser amount of oxygen than those with -ate ending
11
Q
3-D arrays of charged particles when in solid form
A
Ionic compounds
12
Q
Gram equivalent weight
A
Mass of a substance that can donate one equivalent of the species of interest
13
Q
Neutralization reaction
A
Reaction of an acid and base to form a salt and water
14
Q
Combustion Reactions
A
Occur when a fuel and an oxidant (typically oxygen) react to form water and CO2
15
Q
Phosphate
A
PO4^3-
16
Q
Phosphide
A
P^3-
17
Q
Hypochlorite
A
ClO-
18
Q
Ammonium
A
NH4+
19
Q
Bicarbonate
A
HCO3-
20
Q
Nitrite
A
NO2-
21
Q
Homogenous catalyst
A
Catalyst that is the same phase as the reactants
22
Q
Heterogenous catalyst
A
Catalyst that is a different phase than the reactants
23
Q
For a second-order reaction, which graph has a linear slope?
A
1/[A] vs time. Slope = k
24
Q
Slope of [reactant] vs time graph equals ____ for a zeroth order reaction
A
k
25
Which order reaction has a linear graph of ln([A]) vs time, and what is the value of the slope?
First-order; -k
26
Comparing reaction coordinate diagrams of different reactions under the same conditions, the reaction with the ____est ____ ___ will proceed the fastest, regardless of entropic change.
lowest activation energy
27
Endergonic vs Endothermic
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
28
Exergonic vs exothermic
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
29
T/F: The presence of a catalyst affects K_eq of a reaction
False. A catalyst does not affect equilibrium position
30
Law of Mass Action
Expression for equilibrium constant, K_eq = [C]^c*[D]^d / ([A]^a*[B]^b)
31
Reaction Quotient
Same calculation as K_eq, but is calculated using the instantaneous concentrations of each species.
Q = [C]^c*[D]^d / ([A]^a*[B]^b)
32
Q>K_eq
There are more products, currently, than there would be at equilibrium. Reaction proceeds left to decrease amount of products
33
Q=K_eq
System is at dynamic equilibrium
34
Q
Currently less products than there would be at equilibrium. Reaction proceeds forwards
35
Which direction does the reaction proceed when pressure is increased on side A (arbitrary)? Decreased?
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
36
Kinetic products vs Thermodynamic products
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
37
Conditions that favor kinetic product over thermodynamic product?
And thermodynamic over kinetic?
Favor kinetic: lower temperatures, low heat transfer
Favor thermodynamic: higher temperature, high heat transfer
38
Combustion of ___-chain hydrocarbons yields greater energy
long
39
Triple point
The temperature and pressure at which all three phases exist in equilibrium
40
For a spontaneous forward reaction, Keq __ Q and deltaG _ 0
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
41
Equation for entropy change
delta S = heat transfer / temperature = q/T
42
Standard free energy change equation
deltaG_0rxn = -RT*ln(K_eq)
43
Isolated systems
Neither matter nor heat can be exchanged with the environment
44
Closed systems
Matter cannot be exchanged with the environment, but heat can be
45
Adiabatic
Processes that exchange no heat with the environment
46
State functions
Path-independent functions that describe the physical properties of an equilibrium state. Ex: pressure, density, internal energy Gibbs free energy, and entropy
47
Standard state
The condition of an element in its most prevalent form under standard conditions (298 K, 1 atm, 1 M concentration)
48
Critical point
Temperature at which the liquid and gas phases of a substance are indistinguishable
49
Enthalpy
Measure of the potential energy of a system found in intermolecular attractions and chemical bonds
50
Hesse's Law
Total change in potential energy of a system is equal to the changes of potential energies of the individual steps of the process
51
Entropy
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)
52
When is entropy maximized?
At equilibrium
53
Equation for generalized enthalpy of a reaction
delta H = H_prod - H-reac
54
Equation for Gibbs free energy from reaction quotient
deltaG = deltaG_rxn + RT*lnQ
deltaG= RT*(Q/K_eq)
55
Equivalent units to 1 atm:
760 mm Hg = 760 torr = 101.3 kPa
56
Ideal Gas Law
PV = nRT
57
Avogadro's Number
6.023 * 10^23 atoms / mol
58
Boyle's Law, Charles' Law, and Gay-Lussac Law
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"
59
Dalton's Law of Partial Pressure
partial pressure = mole fraction * total pressure
P_A = X_A * P_tot
60
Electromagnetic spectrum (order of waves)
Radio, microwaves, infrared, visible, ultraviolet, X-Rays, gamma
61
Visible range of wavelengths
400 nm (violet) to 700 nm (red)
62
Ideal gas constant (R)
.0821 L*atm/(mol*K) = 8.312 J/(mol*K)
63
Equation to show the relationship between density and molecular weight of an ideal gas
M = density_STP * 22.4 L/mol
64
Kinetic Molecular Theory (state the assumptions)
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
Equation for the KE of an ideal gas
KE = 3/2 * k_B *T = 1/2*m*v^2
k_B is Boltzmann's constant: 1.38*10^-23 J/K
66
Root-mean-square speed
Defines average speed of gases in terms of average molecular speed.
u_RMS = sqrt(3RT/M)
67
Graham's Law
r1/r2 = sqrt(M2/M1)
Shows that gases with smaller particle diffuse faster
68
Effusion
The flow of gas particles under pressure from one compartment to another through a small opening
69
Van der Waals Equation of State
(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
Amphoteric Species
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
Nomenclature of oxyacids
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
Amphiprotic
Able to gain or donate an H+. Special class of amphoteric species
73
Water dissociation constant
K_w = [H3O+]*[OH-] = 1e-14
74
pH equation
= -log ([H+]) = log( 1/[H+] )
same for pOH but with OH- in place of H+
75
Based on strength of acids and bases, why is the bicarbonate an ideal buffering system?
CO3^2- is a weak base and its conjugate acid HCO3- is a weak acid, too.
76
Hydrolysis of a salt
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
Strong acid + strong base
= salt + water
This is a neutralization reaction
78
Strong acid + weak base
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
Strong base + weak acid
Formation of a salt, the basic ion of which then deprotonates water to form OH- and make to solution basic
80
Weak acid + weak base
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
Relationship between Ka, Kb, and Kw
Ka * Kb = Kw = 1e-14 (at 298 K)
82
Polyvalent species
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
Gram equivalent weight
The mass of a compound that yields one equivalent (one mole of charge)
84
Titrant
The species of known pH that, in very small volumes, gets added to a larger solution
85
Titrand
The solution of unknown concentration that gets acid or base added to it until reaching the equivalence point
86
Equivalence point in titrations
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
What can the following equation be used for?
Na*Va = Nb * Vb
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
Buffers
Composed of a weak species and a salt of its conjugate
89
Consider the buffer of CH3COOH + NaCH3COOH.
Describe how the system mitigates addition of OH-
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
Hendersen-Hasselbalch equation (in both forms: pH and pOH)
pH = pKa + log(A-/HA)
pOH = pKb + log (B+/BOH)
91
Indicators
Weak organic acids or bases that change color when they become protonated or deprotonated. They change color at the endpoint
92
Buffering capacity
The ability of a system to resist changes in pH
93
Anode
The electrode in an electrochemical cell at which oxidation occurs.
Think: an ox
94
Cathode
The electrode in an electrochemical cell at this reduction occurs.
Think: red cat
95
Electromotive force (what corresponds with positive and negative?)
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
Movement of electrons in electrochemical cells
From the anode to the cathode
97
Galvanic cells
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
Salt bridge
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
Galvanization
This refers to the precipitation process of the cathode itself. Also called plating
100
Cell diagram
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
Electrolytic cells
Non-spontaneous redox reactions occur here, thus requiring the input of energy (electrolysis). Achieves chemical decomposition
102
Electrolysis
Redox reactions drive by external voltage source
103
How and why to obtain pure Na+ and pure Cl-:
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
Charge per electron
-1.6 * 10^-19 C
105
Faraday constant
Equal to the amount of charge contained in one mole of e'
F = 96,485 C (rounds to 10^5 C/mol e-)
106
Electrodeposition equation
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
Concentration cell
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
A rechargeable cell/battery can function as both a(n) ____ cell and a(n) _____ cell
Galvanic; electrolytic
109
Lead storage battery
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
Energy density
A measure of a battery's ability to produce power as a function of its weight
111
Nickel-Cadmium batteries
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
Surge currents
Periods of large currents (amperage). Preferable in appliances like remote controls that demand rapid responses.
113
In a voltaic cell, the anode is considered _____ in charge because it is the _____ of electrons
negative; source
114
In a galvanic cell,
Electrons move from ___ electric potential to ____ electric potential, while current (the flow of positive charge) flows from ___ to ___.
Electrons: Low to high;
Current: High to low
115
Which is the positive half-cell in an electrolytic cell? Why?
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
Which is the negative half-cell in an electrolytic cell? Explain
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
The anode always attracts ____ and the cathode always attracts _____.
Anode: anions
Cathode: cations
118
Isoelectric focusing
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
Ni-Cd cells
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
Nickel-metal hydride (NiMH) batteries
Had basically replaced Ni-Cd batteries because they have higher energy density, are more cost-effective, and slightly less toxic
121
When is a rechargeable battery acting electrolytically?
| Describe the flow of electrons and current
During charging.
Electrons flow from high electric potential to low electric potential. Current flows from low to high.
122
Reduction potential
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
Standard reduction potentials
(E*_red) - They are calculated by comparison to the standard hydrogen electrode (SHE) under the standard conditions
124
Standard hydrogen electrode
Reduction potential of 0 V
125
Standard emf (def and equation)
The difference in standard reduction potential between the two half-cells
E_cell = E_red - E_oxid
126
Sign of E_cell and delta G in galvanic cells
E_cell: positive
| delta G: negative
127
Sign of E_cell and delta G in electrolytic cells
E_cell: negative (because E_cat
128
Dehydrogenases
Enzymes that catalyze oxidation in which the electrons lost are in the form of Hydrogen
129
Changes in the oxidation numbers of transition metals usually correspond with changes in _______.
color
130
Net ionic equations show only....
The ions that participate in the redox reaction
131
Metathesis reactions
Reactions that involve the switching of counterions. AKA double-displacement reactions. They are NOT usually redox reactions
132
Disproportionation reaction
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
Catalase
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
Superoxide dismutase
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
Iodometric titration
Relies on the titration of free iodine radicals
136
Dominant species when the pH of a solution is greater than pKa:
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
T/F: Cooperative enzymes have one active site, regardless of allosteric regulation
False - if they engage in cooperative binding, this means that they MUST have more than one active site
138
Diamagnetic atoms/ions
Characterized by the electron configuration having only complete pairs. Thus, the electrons will be repelled by a magnetic field
139
Paramagnetic atoms/ions
Unpaired electrons are present, so they will be attracted to an external magnetic field
140
Dilution Factor
DF = V_T1 / V_F1
T1: volume transferred
F1: final volume
If more than one dilution, then multiply dilution factors
141
T/F: Mass changes due to radioactive decay are negligible
True!
142
Electron capture
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
Hydration/Solvation
The attractive force of an ion-molecule causing a thin shell of water molecules to surround it
144
Describe the change in entropy when water is forced to interact with nonpolar substances
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
Common Ion Effect
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
Describe the location of equivalence points and pKas on a titration curve
Equivalence point: the center of the steep portion of the graph
pKa point: approximately the center of the more horizontal buffered region
147
Pyranose
6-membered cyclic sugar with 5 C's and 1 O
148
Furanose
5-membered cyclic sugar with 4 Cs and 1O
149
Mass Spectrometry
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
Reaction of LiAlH4 and an ester
Produces a primary alcohol
151
Equation for specific rotation
a_specific = a_obs / (concentration*path length)
a_spec = a_obs / (c*l)
152
Groups of the Representative Elements
Groups 1 and 2, and 13-18
153
Trend in first ionization energy
Increases going to the right across a period (so reactivity decreases)
Decreases going down a group (so reactivity increases)
