Chemistry: Inorganic Chemistry Flashcards
1
Q
A (physical/chemical) property tells how a substance changes into new substances.
A
Chemical
2
Q
(T/F): A chemical property is the same thing as a chemical change.
A
FALSE
3
Q
What are five signs of chemical change?
A
Color change, odor, temperature change, evolution of gas (bubbles), precipitate
4
Q
(T/F): Change in temperature and odor can also be physical changes.
A
TRUE. Heating/painting
5
Q
(T/F): The ability of a substance to support burning is flammability.
A
FALSE. This is heat of combustion; flammability describes ability of a chemical to burn or ignite
6
Q
If property doesn’t depend on the amount of matter present, it is an (intensive/extensive).
A
Intensive
7
Q
Volume, weight, mass, length, and number of things are (intensive/extensive) properties.
A
Extensive
8
Q
(T/F): Both homogenous mixtures and pure substances have uniform composition and properties throughout.
A
TRUE. Pure substances always have the same composition, but mixtures have different compositions across samples
9
Q
Material composed of two or more substances that can be separated by physical methods
A
Mixture
10
Q
Substance that cannot be broken down into chemically simpler components
A
Element
11
Q
Substance that can be broken down into chemically simpler compounds only by chemical methods
A
Compound
12
Q
Filtration, distillation, and crystallization are all (physical/chemical) methods of separation.
A
Physical
13
Q
A sharp melting/boiling point (not a range) indicates a (pure/impure) substance.
A
Pure
14
Q
The passing of a visible beam through a sample indicates a (homogenous/heterogenous) mixture.
A
Heterogenous as it indicates small suspended particles scatter the light (suspended particles = hetero)
15
Q
When a certain soft metal is burned in oxygen, lime is produced (with no other products). This indicates lime is a (element/compound).
A
Compound
16
Q
(Mass number/atomic weight) refers to the number of protons and neurons while (mass number/atomic weight) refers to the average of masses of all isotopes
A
Mass number, atomic weight
17
Q
Whose model showed that electrons surround a nucleus
A
Rutherford
18
Q
Whose model showed that farther orbits had higher energy, and that photons are emitted when going down an energy level
A
Bohr
19
Q
An atom or molecule with an unpaired electron
A
Free radical
20
Q
If there are no unpaired electrons, the atom is (diamagnetic/paramagnetic).
A
Diamagnetic
21
Q
Paramagnetic atoms (repel/attract) magnetic fields.
A
Attract
22
Q
C = 1s^2 2s^2 2p^2 is a (diamagnetic/paramagnetic) atom
A
Paramagnetic
23
Q
Rule that states that every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied
A
Hund’s rule
24
Q
(T/F): According to Hund’s rule, the electrons in singly occupied orbitals may have different spins.
A
FALSE; all must have the same spin
25
Principle that states that for paired electrons in an orbital, one must have a spin of +1/2 and the other -1/2
Pauli exclusion principle
26
Principle that states that it is impossible to know the momentum and position simultaneously
Heisenberg uncertainty
27
Quantum number that describes the main electron energy level or shell number and possible values?
n (principal); 1,2,3,...
28
(T/F): The shell number matches the row of the periodic table
FALSE; this does not apply to d and f orbitals (row number-1 for d, 4 & 5 only for f)
29
Quantum number that describes the 3D shape of the orbital and possible values?
l (azimuthal); 0, 1, 2, n-1 (0 = s, 1 = p, 2 = d,...)
30
Quantum number (orientation) that describes the orbital sub-type and possible values?
m_l (orbital sub-type), integers (-L to +L, middle orbital is 0)
31
Quantum number that labels electron spin and possible values?
m_s (spin); +1/2 and -1/2
32
Formula for maximum e- in shell (in terms of n)
2n^2
33
What is the number of orbitals in the 4th shell
16 (1s, 3p, 5d, 7f)
34
Formula for maximum e- in subshell
4(L) + 2
35
Ions that have (lesser/greater) charge have -ous, and those with (lesser/greater) charge have -ic.
Lesser, greater
36
Polyatomic anions that contain oxygen
Oxyanions
37
Ions that have (less/more) oxygen end with -ate and those with (less/more) oxygen end with -ite
Ate, ite
38
In extended series of oxyanions, prefix hyper/per- is added to the one with (most/least) oxygen, and hypo- is added to the one with the (most/least) oxygen.
Most
39
Polyatomic anions that gain H+ to form anions of lower charge add the word ______ in front.
Hydrogen/dihydrogen
40
Group 1 is called the ______ family.
Alkali metals (H, Li, Na, K, Rb, Cs, Fr)
41
Group 2 is called the ______ family.
Alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra)
42
Groups 3-12 are called the _________ family.
Transition metals
43
Under the 5 family system, groups 13-16 consist of which families of elements (3)?
Post transition metals, metalloids, non-metals
44
Under the 9 family system, groups 13-16 are what families (4)?
Earth metals (B), tetrels (C), pnictogens (N), chalcogens (O)
45
Group 17 is called the ______ family.
Halogens (F, Cl, Br, I, At, Ts)
46
Group 18 is called the _______ family.
Noble gas (He, Ne, Ar, Kr, Xe, Rn, Og)
47
The last two periods of the periodic table are called the _________. What elements (Z) are part of these rows?
Rare earth metals; 58-71 & 90-103
48
Actual amount of positive (nuclear) charge experienced by an electron in a polyelectronic atom
Zeff = (Z - S where S is shielding constant)
49
Pull between nucleus and valence electrons
Zeff (effective nuclear charge)
50
Zeff (increases/decreases/is the same) across a period and (increases/decreases/is the same) down a group
Increases, is the same
51
Energy required to remove an electron from a neutral atom in its gaseous phase
Ionization energy
52
The lower the ionization energy, the more readily the atom becomes a (cation/anion)
Cation
53
Ability of an atom's inner electrons to shield its positively-charged nucleus from its valence electrons; results in decrease of IE down a group
Electron shielding
54
(T/F): Electron affinity is conceptually the opposite of electronegativity
FALSE; ionization energy is the opposite of EN
55
Ability of an atom to accept an electron (quantitative measurement of energy change when electron is added)
Electron affinity
56
∆H rxn for atoms with high electron affinity is (less than/greater than) 0 when gaining e-
Less than (note that EA is still reported as positive value)
57
Atom's tendency to attract and form bonds with (gain) electrons; force atom exerts on e- in a bond
Electronegativity
58
Of the noble gases, which ones have EN?
Kr and Xe
59
One-half the distance between the nuclei of two atoms
Atomic size
60
Which trend is the only one that decreases across a period and increases down a group?
Atomic size
61
Are anions bigger or smaller than neutral atoms?
Bigger (due to e-e repulsion)
62
Mercury is better suited than water for use in a barometer chiefly because mercury has _______.
High density and little evaporation
63
Hypothetical charge of an atom if all of its bonds to different atoms were fully ionic
Oxidation state
64
Charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally
Formal charge
65
Formula for formal charge
Ve - nonbonding e - (total e shared in bonds)/2
66
Bond formed via the sharing of electrons between two elements of similar EN
Covalent
67
Refers to whether a covalent bond is single, double, or triple
Bond order
68
Higher bond order corresponds to (higher/lower) bond strength, (higher/lower) bond energy, and (higher/lower) bond length.
Higher, higher, lower
69
How many pi bonds does a triple bond have
2 (one sigma, 2 pi)
70
(Nonpolar/polar) bonds have ∆EN of 0.5-1.7
Polar
71
A bond that has an ∆EN higher than 1.7 is?
Ionic
72
Bond wherein a single atom provides both bonding electrons, as most often found in Lewis acid base chemistry
Coordinate covalent bonds
73
H bonds are (strong/weak) attractions between EN atoms of one compound and H of another compound.
Weak
74
Bond formed via the transfer of one or more electrons from an element with relatively low IE to element with high EA
Ionic
75
Large, organized arrays of ions
Crystalline lattices
76
Formula for compounds that shows simplest whole-number ratio of atoms
Empirical
77
Formula for compounds that shows exact # of atoms of each element
Molecular
78
In (electronic geometry/molecular shape), bonded and lone pairs are treated the same, but lone pairs take up more space than bonds in (electronic geometry/molecular shape).
Electronic geometry, molecular shape
79
Possible shapes for an atom with 2 e- groups + bond angle
Linear; 180º
80
Possible shapes for an atom with 3 e- groups + bond angle
Trigonal planar, bent; 120º
81
Possible shapes for an atom with 4 e- groups + bond angle
Tetrahedral, trigonal pyramidal, bent; 109.5º
82
Possible shapes for an atom with 5 e- groups + bond angle
Trigonal bipyramidal, sawhorse, t-shaped, linear; 90º & 120º
83
Possible shapes for an atom with 6 e- groups + bond angle
Octahedral, square pyramidal, square planar, t-shaped, linear; 90º
84
Van de Waals forces includes which intermolecular forces?
London dispersion and dipole dipole
85
Type of force acting between atoms and molecules that are normally electrically symmetric and thus only create temporary dipoles
London dispersion forces
86
Attractive forces between the positive end of one polar molecule and the negative end of another polar molecule
Dipole dipole interactions
87
Strongest intermolecular force in which hydrogen is in proximity to a highly EN element (N,O,F)
Hydrogen bond
88
Why do we assume that gas particles experience no intermolecular forces?
Because gas particles are usually a long distance from one another
89
Ionic compounds conduct electricity when molten (liquid) or in aqueous solution (dissolved in water), because _______________.
Their icons are free to move from place to place
90
Boiling occurs at (one/multiple) temperatures, while evaporation occurs at (one/multiple) temperatures.
One, multiple
91
Which element has the highest melting point?
Carbon
92
Maximum concentration of a solute that can dissolve in a solvent at a given temperature
Solubility
93
Units of solubility
mol/L; g/L
94
Factors that affect solubility (4)
Concentration of solute, temperature of the system, pressure (for gasses), polarity of solute and solvent
95
Describe the curve that represents the relationship between solubility and temperature of most solids
Exponential upward
96
Solubility equilibrium is achieved when the rate of ________ is equal to the rate of _________.
Dissolution, precipitation
97
Solubility equilibrium equation (how much a salt will dissolve)
Ksp = [A+][B-], usually has only one value for a given salt at a given temperature
98
A Ksp that is greater than 1 indicates a (soluble/insoluble) salt.
Soluble (same applies when comparing Ksps of salts that produce same number of ions)
99
Given CaF2 dissolved in water, eq. [Ca2+] was found to be 2 x 10^-4. What is the Ksp at 25ºC?
Ksp = [Ca2+][F-]^2
Ksp = [2 x 10^-4][4 x 10^-4]^2
Ksp = 3.7 x 10^-11
100
Equilibrium constant for complex formation; is usually (greater/lower) than Ksp
Formation or stability constant (Kf), greater
101
Effect of how the solubility of a compound in a solution that already contains one of the ions in the compound
Common ion effect
102
When a central cation is bonded to the same ligand in multiple places to sequester toxic metals
Chelation
103
Molarity of the solute at saturation
Molar solubility
104
Cation bonded to at least one ligand which is the e- pair donor; held together by coordinate covalent bonds
Complex ions
105
Formation of complex ions (increase/decrease) solubility
Increase
106
Given CaF2 dissolved in water, eq. [Ca2+] was found to be 2 x 10^-4. What is the molar solubility of CaF2
CaF2: Ca^2+ = 1:1
also 2 x 10^-4
107
For acids, those with more oxygen are called (-ous/-ic)
-Ic
108
Average kinetic energy depends on (temperature/number of moles/both).
Temperature only (2 moles at 546K > 1 mol at 293K)
109
(T/F): Average speed of particles depends on molar mass.
TRUE
110
Theoretical gas whose molecules occupy negligible space and whose collisions are perfectly elastic
Ideal gas
111
Ideal gases have (high/low) temperatures and (high/low) pressures.
High, low
112
Volume of 1 mol of gas at STP
1 mol gas
113
Ideal gas law formula
PV = nRT (R = 8.314 J/mol K)
114
1 atm = ____ mmHg = ___ torr = ____ kPa = ___ psi
760, 760, 101.3, 14.7
115
Formula for density of gases in terms of ideal gas law
PM/RT
116
Combined gas law formula
P1V1/T1 = P2V2/T2
117
In the combined gas law, what variable is constant?
n
118
Avogadro's principle formula
n1/V1 = n2/V2 (constant T&P)
119
Boyle's Law formula
P1V1 = P2V2 (constant n&T)
120
Charles's Law formula
V1/T1 = V2/T2 (constant n&P)
121
Gay-Lussac's Law formula
P1/T1 = P2/T2 (constant n&V)
122
A 5 liter bulb and a 3 liter bulb are attached by a closed stopcock. The larger bulb contains helium gas at 540 torr, while the smaller bulb contains neon gas at 320 torr. What is the pressure in the larger bulb after the stopcock is opened and equilibrium is achieved?
Use Dalton's law of partial pressures to get the individual gasses' pressures at the total volume, then add: 457.5 torr
123
To change a gas into a liquid, (reduce/increase) pressure and (reduce/increase) temperature.
Increase, reduce
124
(T/F): Real gases deviate from ideal behavior at lower temperature and high pressure
TRUE.
125
At MODERATELY high P, low V, low T, real gases occupy (less/more) volume than predicted by ideal gas law due to intermolecular attractions
Less
126
At EXTREMELY high P, low V, low T, real gases will occupy (less/more) volume than predicted by ideal gas law because the particles occupy physical space
More
127
Dalton's Law (total pressure)
PT = PA + PB + PC + ...
128
Dalton's Law (partial pressure)
PA = XaPT (X = mol fraction)
129
Fruits are canned while hot because ______.
The pressure inside will decrease (lower than atmospheric) once the fruit cools.
130
All diatomic gases
H2 N2 F2 O2 I2 C2 B2 (Have No Fear Of Ice Cold Beer)
131
Henry's Law
[A] = kH x PA (or [A]1/P1 = [A]2/P2 = kH)
[A] - concentration of A, kH = Henry's constant, PA = partial pressure of A
132
Loss of water (or a solvent) of crystallization from a hydrated or solvated salt to the atmosphere on exposure to air
Efflorescence
133
Formula for root-mean-square speed of a gas
u_rms = sqrt(3RT/M)
134
Graham's law of diffusion formula
r1/r2 = √(M2/M1)
where r = rate, M = mass (note inverse rel.)
135
Movements of gas from one compartment to another through a small opening under pressure
Effusion
136
Pressure exerted by a vapor in thermodynamic equilibrium with the condensed phases (solid or liquid)
Vapor pressure
137
Boiling point denotes the temperature at which ______________.
Vapor pressure equals the pressure of the gas above it (e.g. atmospheric)
138
Vapor pressure is the pressure at which there is an equal amount of molecules ______ and ______.
Evaporating, condensing back to the liquid
139
Having a low vapor pressure denotes a (lower/higher) BP.
Higher; need more energy to equate vapor pressure to atmospheric pressure
140
The freezing point denotes the temperature at which _________.
Solid and liquid phases coexist in equilibrium (at 1 atm)
141
In a phase diagram, an exponential curve suggests _________.
The transition from liquid to gas
142
If you want to remove carbon dioxide from a solution, you ____ the solution.
Boil, carbon dioxide has very low BP
143
Heating of solids to a high temperature for the purpose of removing volatile substances, oxidizing a portion of mass, or rendering them friable
Calcination
144
Type of reaction where two or more reactants form one product
Combination
145
Type of reaction where a single reactant breaks down
Decomposition
146
Type of reaction that involves a fuel (usually an HC) and O2
Combustion
147
Products of combustion
CO2 and H2O
148
Type of reaction wherein an atom/ion in a compound is replaced by another atom/ion
Single-displacement
149
Type of reaction where elements from two compounds swap places
Double-displacement (metathesis)
150
Type of reaction wherein acid + base --> salt + H2O
Neutralization
151
Neutralization is a type of what reaction
Double replacement
152
Avogadro's number
6.022 x 10^23 = 1 mol
153
Represents the fraction of collisions that have enough energy to overcome the activation barrier
Arrhenius equation (k = A x e^(-Ea/RT))
154
Rate law
rate = k [A]^x[B]^y
155
Kinetic products are (higher/lower) in free energy than thermodynamic products can form in (higher/lower) temperatures
Higher, lower
156
Thermodynamic products are (faster/slower) and (more/less) spontaneous
Slower, more
157
Formula for heat energy change (∆H)
∆H = ∆H (bonds broken in reactants) - ∆H (bonds made in products)
158
Energy required to reach transition state
Activation energy
159
The activation energy of an exergonic reaction is (less/more) than its reverse.
Less
160
The change in internal energy ∆U is given by
∆U = q + w (+w for work done on system, -w for work done on surroundings)
161
Change in Gibbs free energy is defined as
∆G = ∆H - T∆S
162
When ∆G is negative, a process will proceed (spontaneously/unspontaneously) and is (exergonic/endergonic)
Spontaneously, exergonic
163
When ∆G = ____, system is in equilibrium and the concentrations of products and reactants is constant
0
164
Gibbs free energy in terms of activation energy is
Ea - Ea(rev)
165
Le Chatelier's Principle states that if a stress is applied to a system, the system shifts to _______ the applied stress.
Relieve
166
A measure of the degree to which energy has been spread throughout a system or between a system and its surroundings.
Entropy
167
Increasing pressure shifts the equilibrium towards the side with __________.
Fewer total molecules (A molecule is a mole on the equation (# terms x coefficients))
168
Increasing the reactants shifts the equilibrium to the (left/right).
Right
169
If the system's reaction enthalpy is ∆H = 250 kJ, and temperature is decreased in the system, the equilibrium shifts to the (left/right)
Left
170
Raising the temperature of a solution of a gas in a liquid (increases/decreases) solubility thereby (increasing/decreasing) equilibrium concentration.
Decreases, decreases (because more entropy & freedom)
171
Solubility of a gas in liquid (increases/decreases) with increasing temperature.
Decreases
172
According to Henry's Law, the solubility of gases (increases/decreases) as the partial pressure of the gas above a solution increases.
Increases
173
Process where a solute in gaseous, liquid, or solid phase dissolves in a solvent to form a solution
Dissolution
174
According to the modern theory, when sodium chloride is dissolved in water, it forms ________ sodium and chloride ions.
Hydrated
175
Are ethyl and methyl miscible or immiscible in water? Why?
Miscible; attraction of OH to water
176
In dissolution, the breaking of the solute-solute bonds & solvent-solvent bonds is (endothermic/exothermic), and the bonding of the solute with the solvent is (endothermic/exothermic).
Endothermic, exothermic
177
In an endothermic dissolution, (less/more) energy is released when the solute-solvent bond compared to the energy needed to break the solute particles.
Less
178
In endothermic reactions, solubility is (increased/decreased) when the temperature is increased.
Increased; promote dissolution with additional heat
179
In exothermic reactions, solubility is (increased/decreased) when the temperature is increased
Decreased; inhibit dissolution to counter more heat loss with additional heat
180
The dissolution of gasses is (exothermic/endothermic).
Exothermic; increase in temp --> increase in KE --> weaker bonds
181
Equilibrium constant or reaction quotient formula
Keq = ([C]^c [D]^d)/([A]^a [B]^b); exponents are the coefficient
182
Reaction quotient and equilibrium constant have the same formula, but (reaction quotient/equilibrium constant) will fluctuate as the system reacts, whereas the (reaction quotient/equilibrium constant) is based on equilibrium concentrations
Reaction quotient, equilibrium constant
183
(T/F): Q < Keq means a ∆G < 0
TRUE; whereas an equal Q = Keq means equilibrium
184
Basis of Arrhenius definition of acids & bases
H+/OH- in H2O
185
Basis of Bronsted & Lowry definition of acids & bases
H+ donor/acceptor
186
Basis of Lewis definition of acids & bases
E pair donor/acceptor
187
Species that can behave as an acid or base
Amphoteric base
188
Acid with multiple ionizable H atoms
Polyprotic acid
189
Strong acids have (high/low) solubility in water.
High
190
Formula for pH
pH = -log [H+], so [H+] = 10^-PH
191
Formula for acid dissociation constant
Ka = [H+][A-]/[HA]
192
What are the seven strong bases?
LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2, Mg(OH)2
193
Is ammonia a strong or weak base?
Weak; thus, NH4OH is also weak
194
What are the seven strong acids?
HCl, HBr, HI, H2SO4, HClO4, HClO3, HNO3
195
Is the bisulfate ion (HSO4-, not found as molecule) a strong or weak acid?
Weak
196
A (strong/weak) acid + a (strong/weak) base will give an alkaline solution.
Weak, strong
197
In every solution, pOH + pH = 14 because the product of [H+] and [OH-] must always equal the equilibrium constant for the ionization of water, which is ________.
1 x 10^-14
198
Formula for pH
-log ([H+])
199
Formula for pOH
-log ([OH-])
200
If I dilute 5 mL of 0.15 M NaCl to a final volume of 5 L, what is the final concentration of NaCl?
0.00015 M
201
Formula for radioactive decay
N(t) = N0 (1/2)^(t/t.5)
where N0 - initial amt, Nt = final amt, t = time, t.5 = half life
202
Equivalent mass is the mass of an acid that yields _________ or mass of a base that reacts with ________.
1 mole of H+ (thus is equal to #H or OH of the compound)
203
GEW (gram equivalent weight) formula (basically equivalent weight
Molar mass / mol H+ or e-
204
Equivalents formula
Mass of compound/GEW
205
Normality formula
Eq/L (Eq - equivalents of solute, V - volume of solvent in liters) or M*Eq (M - molarity)
For acids, Eq (n) is the # of H+ available from a formula unit
206
Can donate or accept multiple equivalents
Polyvalent
207
A student wants to prepare 500 ml of 8 N H2SO4 from a solution which is 8 molar. What volume of the concentrated acid is needed?
250 mL (8N H2SO4 = 4 M H2SO4 * 2 Eq)
208
Law of equivalence (used to neutralize the whole solution)
N1V1 = N2V2 where 1 is an acid and 2 is a base, or used to know after dilution
209
Molarity formulas
mol/L = normality/mol H+ or e-
210
Colligative properties depend upon the _________ of solute molecules or ions, but not upon the identity of the solute.
Concentration
211
Adding a solute will (lower/increase) equilibrium vapor pressure. Why?
Lower, because less solvent molecules will be able to come to the surface and evaporate
212
Formula for new vapor pressure given addition of a solute
Raoult's Law: Psolvent = Xsolvent P0solvent
213
Adding a solute will (lower/increase) boiling point. Why?
Elevate, because it will require more energy to reach the same vapor pressure needed to equate the external surroundings (given more solute = lower vapor pressure)
214
Formula for new boiling/freezing point given addition of a solute
∆T = iKbm or ∆T = iKfm
i - van't Hoff factor (# of particles into which solute dissociates), m - molality, Kb - molal boiling point constant (for water, 0.5121ºC/m), Kf - molal freezing point constant (for water, -1.86ºC/m)
215
Adding a solute will (lower/increase) freezing point. Why?
Lower, because the addition of solute makes the mixture more disordered (and thus with higher entropy) and requires lower temperature to make ordered (frozen)
216
Pressure required to achieve osmotic equilibrium
Osmotic pressure (by adding this pressure to a solution separated to a pure solvent by a semipermeable membrane, the escaping tendency of the solution in spite of high solute concentration will be raised to the level of the pure solvent)
217
Number of individual particles in solution
Osmolarity (1M NaCl = 2 osmol/L)
218
T/F: Osmotic pressure is the pressure required to sustain osmosis.
FALSE: stop osmosis
219
Formula for osmotic pressure
II = nRT/V or MRT (T is in K, M is total ionic concentration of solute, i.e. 2x molarity for NaCl in sol. as it dissociates)
220
Acidity is increased when the _________.
H+ is more available
221
Carboxylic acids have greater acidity compared to alcohols because of __________.
Electron delocalization in the carboxylate ion
222
Things that affect acidity include:
Availability of proton, stability of conjugate base
223
Higher s orbital character is more (acidic/basic).
Acidic because more s character (more unsaturated) means lower energy (stable)
224
Things that affect basicity include:
Availability of lone pair, stability of conjugate ion
225
If resonance of the lone pair is possible, it is (less/more) basic.
Less
226
EWG increase (basicity/acidity) and decrease (basicity/acidity).
Acidity, basicity
227
Higher negative charge indicates more (acidity/basicity) while more positive charge indicates more (acidity/basicity).
Basicity, acidity
228
What are the most basic amines? Why?
Secondary amines. While primary amines are more solvated due to more hydrogen bonds, it does not have EDG (alkyl groups) which increase basicity.
229
