MCAT Chemistry 2 Flashcards
1
Q
ex of a closed syst
A
steam radiator
2
Q
____________ is a meas of the averag ke of the particles in a syst
A
temperature
3
Q
G is ?
A
free energy
4
Q
bond e or bond dissoc e is?
A
averag of the energy req to break a part type of bond in one mole of gaseous molec
5
Q
give an ex of an isolated system?
A
insulated bomb reactor
6
Q
T or Fit is not poss to meas H directly
A
T
7
Q
system
A
partic part of universe being studied
8
Q
open system is?
A
can exch matter and e with surroundings.
9
Q
ΔG <0 means a proc can?
A
occur spont
10
Q
qrev/T =
A
ΔS
11
Q
conversion of C (diamond) to C (graphite) is __________ but its rate is slow
A
spont
12
Q
bond ____________ is always endothermic
A
breakage
13
Q
stand heat of format of a compound is?
A
enthalpy ch that would occur if one mole of a comp were formed direct from its elem in their standard states
14
Q
T or Fheat and temp are different
A
t
15
Q
boiling has a ____________ in entropy
A
increase
16
Q
state functions
A
prop that dep only on initial and final states of the system
17
Q
E or U is?
A
internal energy
18
Q
the react quotient Q is=
A
Q=[C]^c[D]^d/[A]^a[B]^b
19
Q
units of heat?
A
joules
20
Q
ΔH ΔS Outcome ?- +
A
spontaneous at all temperatures
21
Q
heat changes at const press, is
A
enthalpy
22
Q
isothermal process is?
A
temperat of systm is const
23
Q
a system undergoes a process when?
A
when one or more of its prop change
24
Q
ΔH ΔS + +
A
spontaneous only at high temperatures
25
-RTlnKeq =
ΔG
26
system can be?
isolatedclosedopen
27
ΔH ΔS - -
spontaneous only at low temperatures
28
t or fentropy is a state function
T
29
t or frate of a react dep on ΔG
f
30
for hesses law, if you multiply the prod and react by 3 ( or a cert number) then what must you do to ΔH?
also multiply it by 3
31
endothermic react?
absorb e
32
q=mc chTc is?
specific heat
33
q=mc chTq is?
heat
34
T or FΔHreact=Hprod-Hreact
T
35
enviroment
everthing outside system
36
freezing has a ___________ in entropy
decrease
37
a spontaneous react may or may not?
proceed to completion
38
bomb calorimeter, the overall system is?
adiabatic
39
standard free energy of format of a comp is?
free energy ch that ocurs when 1 mol of a comp in its stand form
40
ΔHf of an element in its standard state is?
zero
41
second law of thermodyn says that?
all spont proceed so that entropy of the systm plus its surr increases
42
temp is?
a measure of average ke of particles in a systm
43
name 7 state funct?
tempvolenthalpyentropyfree energyinternal energy E and U
44
H is?
enthalpy
45
what condit are heat changed measured under?
const vol calorimetryconst press calorimetry
46
enthalpy of a process does not depend on the?
path
47
rate of a reac depends on?
activat e
48
standard condit in thermody must not be confused with?
standard temperature and pressure in gas laws STP
49
closed system is?
can exch e but not matter
50
ΔH ΔS + -
nonspontaneous at all temperatures
51
isobaric process is?
press of systm is const
52
q react + q water + q steel=0 in a ?
bomb calorimeter
53
Goose hunters take shotguns reminds you of?
ΔG=ΔH-TΔS
54
discuss when water boils in terms of ch in G
TchS> chH only when T is above 373 K
55
reverse reaction has the same __________ as that of the forward react, but its ________ is oppos
enthalpy,but its sign is oppos
56
for simple diatomic molec such as H2, bond dissoc e can be easily measured?
spectroscopically using react E=hv E=bond dissoc e and v is freq of light absorbed
57
ΔS= ?/T
qrev/T
58
when a react commences ΔG=?
ΔG= ΔG° + RTlnQ
59
calorimetry measures?
heat changes
60
spectroscopy can be used to measure the bond dissoc e of?
simple diatomic molecules
61
standard free energy ΔG is ?
ΔG of a proc occur at 25 C and 1 atm . concentr of sol are 1 M
62
ΔS=Sfin- Sinit
Sfin- Sinit
63
standard heat of react is?
ΔHreact= (sum of ΔHf of prod)-(sum of ΔHf of react)
64
exothermic react?
release e
65
a system undergoes a ____________ when one or more of its properties change
process
66
standard state of a subs is the form a subst takes?
at 25 C and 1 ATM
67
Hess law says?
that enthalpies of react are additive
68
heat absorbed by a system is?
pos
69
Thermodynamics
- Is the study of energy and it's relationship to macroscopic properties of chemical systems. It's functions are based on probabilities and are only valid for systems that are composed of a large number of molecules. * I.e The rules of thermodynamics gov
70
System vs. Surroundings
A system is a part of a universe that is the macroscopic body under study, and the surroundings is everything else. Systems are based on mass and energy exchange with the surroundings. There are three systems: 1) open: exchange both mass and energy with surroundings.2) closed: Exchange energy but not mass3) isolated: Do not exchange energy or mass.
71
State Function
A state is the physical condition of a system described by a specific set of thermodynamic property values. Such properties that describe the state of a system are called State Functions.Two types of properties used to describe the state of a system: 1) Extensive: properties are propotional to the size of the system (eg. V + n)2) Intensive: are independent of the size of the system. ( eg. P + T )Seven State Functions: U (internal Energy), T, P, V, H (Enthalpy), S (entrophy) and G (Gibbs energy).
72
Heat
(q) *Heat has three forms:1) Conduction2) Convection3) Radiation*Always the movement of energy from hot to cold.
73
Work
work = P∆V
74
Convection
- Is the thermal energy transfer via fluid movements. - differences in density or pressure drive warm fluid to cold fluid*Oceans and air currents are commone examples of convection.
75
Radiation
- is thermal energy transfer via electromagnetic waves. (eg. heated metal red, orange, white, blue-white) - All objects above 0K radiate heat.P= σ∈AT⁴
76
The First Law of Thermodynamics
∆E = q + w*Warning: work ON the system is positive for "convention" a passage on MCAT may define work done BY the system as a postive in which case you use this formula.
77
The Second Law of Themodynamics
*See heat engines
78
Internal Energy Types
MCAT may refer to internal energy as "heat energy", "thermal energy," or even "heat.""Heat energy and thermal energy" are really the vibrational, rotational and translational parts of interanl energy. Called this because they effect temperature. "Heat" is a transfer of energy. Don't mistake.
79
Temperature
0 K = 273°C
80
Enthalpy
(H) = is a man- made property that accounts for this extra capacity to do PV work. Unlike functions such as pressure, volume , and temperature, enthalpy is not a measure of some intuitive property. Defined more so as an equation then a property:H≡ U + PV
81
Standard State
* Don't confuse with STP. STP is at 0°C whereas standard state is at 25°C and is arbritrarily assigned an enthalpy value of 0 J/mol
82
Reference Form
Define
83
Standard Enthalpy of Formation
Define
84
Exothermic vs. Endothermic
Define
85
Activation Energy
Define
86
Transition state
Define
87
Intermediates
Define
88
Catalyst
Define
89
Irreversible vs. reversible reactions.
Define
90
Third law of Thermodynamics
Define
91
Spontaneity
Define
92
Gibbs Free Energy
Define
93
Thermodynamics
- Is the study of energy and it's relationship to macroscopic properties of chemical systems. It's functions are based on probabilities and are only valid for systems that are composed of a large number of molecules. * I.e The rules of thermodynamics gov
94
What are two ways to transfer energy between systems
#NAME?
95
Conduction
- is thermal energy transfer via molecular collisions. - Requires direct physical contact- an objects ability to conduct heat is called it's thermal conductivity (k) **On page 46 go over again.Q / t = kA [ T(h) - T(c) ) / L ]
96
Stephan- Boltzman Law
P= σ∈AT⁴Where:A is the surface area of the objectT is the temperature σ is the Stefan- Boltzman constant (5.67 x 10⁻⁸ )∈ is the emissivity of the object's surface.
97
Zeroth Law of Thermodynamics
Discovered after 1st, 2nd, and 3rd Law of Thermodynamics. All three rely on this Law, as it is based off of temperature. States: Two systems in thermal equilibrium with a third system are in thermal equilibrium with each other. The law declares that the two bodies in thermal equilibrium share a termodynamic property, which must be a state function. This is temperature.
98
Standard Temperature and Pressure
(STP)
99
Mean free path
Define
100
Ideal gas
Define
101
Kinetic Molecular theory
Define
102
Ideal gas law
PV = nRTK.E.(avg) = 3/2 (RT)
103
Standard Molar volume
22.4L ---> At STP one mole of any gas (behaving ideally will occupy this amount.
104
Partial Pressure
Pa= Xa Ptotal
105
Dalton's Law
P (total) = P1 + P2 + P3 ...
106
Graham's Law
v1/v2 = √M2 / √M1
107
Effusion
effusion rate 1/ effusion rate 2 = √M2 / √M1
108
Diffusion
Define
109
Chemical Kinetics
Define
110
Collision Model
Define
111
Activation Energy
Define
112
The effect of temperature on the rate of reaction
The rate of a reaction increases with temperature
113
Intermediates
Define
114
Rate Law
Define
115
Order of each respective reactant
Define
116
Overall Order of a reaction
Define
117
Rate Determining Step
Define
118
Catalyst
Define
119
Heterogeneous catalyst
Define
120
Homogeneous Catalyst
Define
121
Chemical Equilibrium
Define
122
Equilibrium Constant
K
123
The law of Mass Action
Define
124
Reaction Quotient
Define
125
Le Chatelier's Principle
Define
126
Metals
- large atoms that tend to lose electrons and form positive ions (cations) to obtain a noble gas configuration. - Can bond with other metal atoms (metallic bonding) - can also form ion bonds with negative ions (anions) *Usually non-metals. -Metals are
127
Ductile
- easily stretched. - occurs in metals- if you stretch or hammer e- the metals can adapt (which is why you can stretch metal without breaking them)
128
Luster
#NAME?
129
Non- metals
- form negative ions (anions) to obtain a noble gas configuration. - form covalent bonds with other non-metals- form ionic bonds with positive ions (usually metals)- poor conductors of heat and electricity- non-malleable, brittle and possess low to moder
130
Metalloids
#NAME?
131
Transition metals
- wont be focused on for MCAT, just know that they exist and that this might be a passage based question. - 1/2 vs fully filled orbitals occur here often.
132
Alkaline Earth Metals
- Harder metallic sollids have 2 valence electrons (can easily form +2 cations)- have a higher melting point than solf metallic solids- Heavier are more reactive than lighter ones.
133
Halogens
F₂: gas at RT and 1 atmCl₂: gas at RT and 1 atmBr₂: liquid at RT and 1 atmI₂: solid at RT and 1 atm* all highly reactive unless in diatomic form. -7 valence e-
134
Noble gases
- 8 valence e-- complete octet so very stable and non-reactive.
135
How do you determine the number of valence electrons in a main group atom?
- look at the group numberif group 1= 1 e- on valence shell (+1)group 2= 2 e- on valence shell (+2)recall: First shell closest to nucleus = ground stateLast shell closest to outside world = Valence e- shell. *Valence e- are involved in chemical bonding
136
Atoms
Define
137
Protons
Define
138
Electrons
Define
139
Nucleus
Define
140
Neutrons
Define
141
Elements
Define
142
Mass Number
(A)
143
Atomic Number
(Z)
144
Isotopes
Define
145
Atomic Weight
Define
146
Molar Mass
(MM or M)
147
Atomic Mass Units
(amu)
148
Mole
Define
149
Avogadro's Number
Define
150
Periodic Table
Define
151
Period
Define
152
Groups or Families
Define
153
Metals
Define
154
Nonmetals
Define
155
Metalloids
Define
156
Transition Metals
Define
157
Inert Gases
Define
158
Shields
Define
159
Effective Nuclear Charge
(Z eff)
160
Periodic trends
Define
161
Atomic Radius
Define
162
Ionization Energy
Define
163
Second Ionization Energy
Define
164
Electronegativity
Define
165
Electron affinity
Define
166
Metallic Character
Define
167
Bonds
Define
168
Covalent Bonds
Define
169
Bond Length
Define
170
Bond Energy
*Or bond dissociation energy
171
Compound
Define
172
Empirical Formula
Define
173
Molecules
Define
174
Molecular formula
Define
175
Ionic Compounds
Define
176
Acids
Define
177
Binary Molecular compounds
Define
178
Physical reaction
Define
179
Chemical reaction
Define
180
Runs to completion
Define
181
Fundamental Reaction types
1) Combination: A+B ---> C2) Decomposition: C ---> A + B3) Single Displacement: A + BC --> B + AC4) Double Displacement: AB +CD ---> AD + CB
182
Principle Quantum Number
(n)*Note: quantum numbers are the equivalent of a mailing address for an electron. - Description: Energy level and average distance from nucleus- Possible values: n= 1, 2, 3 etc.
183
Shell
Define
184
Valence electrons
Define
185
Azimuthal Quantum number
(l)*Also known as angular momentum quantum numberDescription: Orbital shape (s, p, d, f)Possible values: l=0, to ...n-1l=0 is the s-subshelll=1 is the p-subshelll=2 is the d-subshelll= 3 is the f-subshell
186
Subshell
eg. s, p, d and fs= sphericalp= dumbbell shapedd= clover leaf
187
Magnetic Quantum number
m- Description: Orbital orientation (px, py, pz)- Possible values: m= -l to +lfor l=0 (spherical s orbital) there is only one orientation.for l=1 (dumbbell p orbital_ there are three orientations along each of the x, y, and z axes.
188
Atomic Orbital
Space around a nucleus in which e- have the ability to exist.- max 2 e- in each orbital.
189
Electron Spin quantum Number
m (s)- Description: Describes spin of electron- Possible values: m(s) = +½ or -½Either clockwise or counter clockwise.
190
Pauli exclusion Principle
#NAME?
191
Heisenberg Uncertainty Principle
- tells us that the exact position and momentum of such an electron cannot be measure simultaneously. *Ie the more you know about the momentum of an electron, the less you know about it's exact position.
192
Aufbau Prinicple
*Write out on a piece of paper before writing mcat.1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f 6s 6p 6d7s 7p
193
Electron Configuration
Know 5 points on page 29 of lecture manual.
194
Ground State
Define
195
Hund's Rule
Define
196
Planck's Quantum theory
Define
197
Photoelectric effect
- an important experiment demonstrating the particle nature of light. An electron can be ejected from the surface of a metal by shining light of a certain frequency. Incoming light (photon) must have enough energy to eject the electron.E(photon) = h(Plank
198
Wavelength vs. Frequency vs. Velocity
Wavelength = distance btwn two consecutive crests of troughs. measured in meters. Frequency= the number of crests that pass through a given point per unit of time (Hertz)Velocity= the distance a wave travels through space/ unit time (m/s)Frequency= (velocity) / (wavelength)
199
Wave particle duality
all matter, including light can be thought of as both a wave and as a particle.
200
Electromagnetic Radiation
(EMR) is simple energy. as energy travels forward electric and magnetic fields are created perpendicular to each other and to the direction of travel.
201
Order of Colours
RedOrangeYellowGreenBlueViolet
202
Absorption Line spectrum vs. Emmision Line spectrum
ALS= Dark lines on a light backgroundELS= Light lines on a dark back ground.
203
How do you determine the max number of e- an atom can have?
#ERROR!:parse
204
Mass Percent (of a solute)
Mass(solute)/Mass(solution)
205
Archimedes Principle
Any object wholly or partially immersed in a fluid is buoyed up by a force equal to the weight of the fluid displaced by the object
206
Buoyant Force
Buoyant Force=Density(medium)*Volume(object)*g
207
Apparent immersed weight
Weight-Weight of Displaced Fluid
208
Molar Volume of an Ideal Gas at STP
22.4 Liters
209
Empirical Formula
Simplest whole number ratio of atoms in a molecule
210
Mass Percent (of an element)
Mass percent=[(mass atoms)/(mass compound)]*100%
211
Molecular Formula
Requires multiplying the empirical formula by the whole number ratio of the molecular mass:empirical mass; requires knowing the molecular mass of the compound
212
Hygroscopic salt
Absorbs water; examples include calcium chloride and magnesium sulfate
213
KOH
Potassium Hydroxide; undergoes a combustion reaction with carbon dioxide to form solid potassium bicarbonate
214
Molarity
Moles of Solute/Liters of Solution
215
Molality
Moles of Solute/Kilograms of Solvent; does not change with temperature so it's used to calculate the boiling-point elevation and freezing-point depression of solutions containing non-volatile impurities
216
Dilution:Fold vs. Parts
Fold (based on the total volume); parts (based on the volume added)
217
Beer's Law
The specific wavelength of light absorbed and the intensity of the absorbance varies with the solute and its concentration respectively
218
Absorbance
Absorbance=εcl, where ε refers to the absorption constant (a constant for the solution at λ max), c refers to concentration, and l refers to path length (or the width of the cuvette).
219
Limiting Reagents
The reactant with the lowest ratio of actual moles to needed moles
220
Precipitation Reactions (or metathesis or double-displacement reactions)
Involves two aqueous salts that react to form spectator ions and a solid salt precipitate
221
Acid-Base Reactions
A reaction between an acid (proton donor) and a base (proton acceptor) that results in the formation of a neutral salt and water
222
Composition Reaction
The number of reactants exceeds the number of products; entropy decreases and bonds are formed
223
Decomposition Reactions
Reactant(s) decompose to form multiple products; entropy increases and bonds are broken
224
Oxidation-ReductionReactions
Transfer of electrons from one atom to another ; the oxidation state must change in a redox reaction
225
Oxidation
Loss of electrons
226
Reduction
Gain of electrons
227
Reducing Agent
The atom (or compound) that is losing electrons
228
Oxidizing Agent
The atom (or compound) that is gaining electrons
229
Combustion Reactions
A special type of oxidation-reduction reactions where the oxidizing agent is the oxidizing agent and the products are oxides (water and CO2)
230
Molarity
(M) * units: mols/LM= (moles of solute) / (Volume of solution)
231
Mole Fraction
(m) *has no unitsm= (moles of solute) / (kilograms of solvent)
232
Solving for X
X= (moles of solute) / (total moles of all solutes and solvent)
233
Parts per million
PPM= (mass of solute) / (total mass of solution) x10⁶*Note: PPM is NOT the number of solute molecules per million molecules. It is the mass of the solute per mass of solution times on million.
234
Solution Formation
Define
235
Vapor Pressure
Define
236
Nonvolatile Solute
Define
237
Raoult's Law
Define
238
Volatile Solute
Define
239
Solubility
Define
240
Precipitation
Define
241
Saturated
Define
242
Solubility product K(sp)
Define
243
Spectator Ions
Define
244
The common Ion effect
Define
245
Solubility Factors
Define
246
Solution
Define
247
Solute vs. Solvent
Define
248
Colloids
Define
249
London Dispersion Forces
Define
250
Solvation
Define
251
Hydration
Define
252
electrolyte
Define
253
Arrhenius Acid vs. Base
Acid = Is anything that produces hydrogen ions in aqueous solution. [ H+ ]Base = is anything that that produces hydroxide ions in an aqueous solution. [ OH- ]*Only aq solutions.
254
Bronsted and Lowry
DefinesAcid: as anything that DONATES a PROTON.Base: as anything that ACCEPTS a PROTON.
255
Lewis
More general then Bronsted/ Lowry or Arrhenius.Acid: Anything that ACCEPTS a pair of ELECTRONSBase: anything that DONATES a pair of ELECTRONS. *Includes B+L acids and bases and more because it also accepts molecules that have incomplete octets of electrons eg. BF3 or AlCL3. Also includes simple cations (Smaller the cation the higher the charge, the stronger the acid) *Excludes alkali and heavy alkaline earth metals.*
256
Acid + Base vs. Conjugate Acid + Base
Acid + Base= Refers to the Reactants [A]Conjugate Acid + Base= Refers to the Product [HA][HA] is the conjugate ACID of BASE [A-][A-] is the conjugate BASE of ACID [HA]*The stronger the acid/Base the weaker the conjugate base/Acid, HOWEVER, weak acids may have EITHER weak OR strong conjugate bases.
257
Polyprotic
Define
258
Diprotic
Define
259
Amphoteric
Define
260
How does molecular structure affect Acid strength
Define
261
Hydrides
Define
262
Autoionization of water
Define
263
Acid Dissociation Constant
K(a)
264
Finiding the PH
Define
265
Titration
Define
266
Titration Curve
Define
267
Equivalency point
or Stoichiometric point
268
Half equivalence point
Define
269
Buffered
Define
270
Henderson- Hasselbalch Equation
pH= pK(a) + log ( [A-] / [HA] )
271
Indicator
Define
272
Endpoint
Define
273
Polyprotic titrations
Define
274
Strong Acids
- Hydroiodic Acid (HI)- Hydrobromic Acid (HBr)- Hydrochloric Acid (HCl)- Nitric Acid (HNO₃)- Perchloric Acid (HClO₄)- Chloric Acid (HClO₃)- Sulfuric Acid (H₂SO₄)
275
Strong Bases
- Sodium Hydroxide (NaOH)- Potassium Hydroxide (KOH)- Amide Ion (NH₂-)- Hydride Ion ( H-)- Calcium Hydroxide Ca(OH)₂- Sodium Oxide (Na₂O)- Calcium Oxide (CaO)
276
______ occurs at the cathode in an electrochemical cell. Electrons flow _____ the cathode.
reduction, toward
277
_____ occurs at the anode in an electrochemical cell. Electrons flow ____ the anode.
oxidation, from
278
The moles of solute over the kg of solvent
molality
279
Normality is the molar _____ per unit volume
equivalents
280
A ____ function is one in which the result in dependent only on the initial and final measurements and independent of the path needed to get there
state
281
_____ describes the the cagelike assortment of solvent particles around solute particles
solvation
282
Hess's Law
The enthalpy of reaction is equal to the difference between the ∆Hformationproducts-∆Hformationreactants
283
Isobaric process
One that occurs under constant pressure
284
Isochoric process
One that occurs under constant volume and no volume-pressure work is done
285
Constant-volume calorimeters (bomb calorimeters) are used to measure
the total heat absorbed or given off by a reaction
286
Adiabatic process
One that occurs without the transfer of heat
287
Collision theory of chemical kinetics
The rate of reaction is proportional to the number of collisions that occur between reaction molecules per second
288
Disproportionation
When a species is oxidized and reduced in the same reaction
289
Partial pressure equation
Ppartial = Ptotal * X(mole fraction)
290
Kinetic Molecular Theory of Gases
Gases have completely elastic collisions with themselves and the walls of their container. Gas particles have negligible volume, negligible attractive forces and exhibit random motion
291
Ideal gases
Monoatomic, no attractive forces or volume
292
Henry's Law:
The partial pressure of a gas above a solution is directly proportional to the partial pressure of the gas dissolved in the solution
293
Graham's Law:
Particles diffusion and effusion rates are inversely proportional to the squareroots of their molecular weights
294
Effusion
The passage of gas from high to low pressure through a small opening
295
Charles and Gay-Lussac's Law
At constant pressure, the volume of an ideal gas is directly proportional to its temperature
296
Boyle's Law
At constant temperature, the pressure of an ideal gas is inversely proportional to its volume
297
Molecular orbital
A regions where bonding or antibonding orbitals overlap, resulting in a low-energy bonding orbital or high-energy unstable antibonding orbital
298
Avagadros Principle
Different gases at the same temperature, pressure and volume, contain the same number of particles
299
Electrochemical reaction
One which needs or produces electricity
300
Nernst Equation
V=Vo-(0.6/n)logQ
301
VSEPR
Shows the 3D geometrical shape of a molecule that is based on the electronic interactions between bonding and non-bonding electrons
302
Isoelectronic
When two atoms have the same electron configuration
303
Magnetic Quantum number
M(L). Range from -L-L. Determines what orbital within the subshell the electron is likely to reside in.
304
Azimuthal quantum number
L. Ranges from 0-(n-1). Determines what subshell the electron is likely to be found in.
305
Aufbau Principle
Electrons fill an atom in order of increasing energy level
306
Exceptions to the aufbau principle
Cu and Cr. Cu takes an electron from a s orbital and places it in a d-orbital, completing its d-orbital. Cr takes an electron from an s orbital and places it in an d orbital, giving it the maximum number of unpaired electrons.
307
Hunds Rule
Electrons fill orbitals such that a maximum number of unpaired electrons results
308
Pauli Exclusion Principle
No two electrons in an atom can have the same set of 4 quantum numbers
309
Arrhenius Definition
Acid dissociates to form H+ in aqueous solution. Base dissociates to form OH- in aqueous solution.
310
Colligative properties
Those that depend only on the number of solute particles and not on the chemical makeup of the particles. (Boiling point elevation, freezing point depression, osmotic pressure, partial pressure)
311
Raoults Law
The vapor pressure above a solution is directly proportional to the mole fraction in solution. Ppartial = Ptotal*X
312
Energy of emitted electron =
hf- electron binding energy
313
malleability and other metallic characteristics ____ down and to the left on the periodic table
increase
314
most likely to have multiple oxidation states
transition metals
315
ionization energy
: E needed to remove an electron from a gaseous state
316
atomic mass increases ____ and ____ on the periodic table
right and down
317
which ions are typically the largest
negative ones (anions)
318
density _______ with increasing atomic radius
decreases
319
diatomic gases will have _____ densities than monoatomic gases
larger
320
T/F: solutions containing ions of transition metals are frequently colored
T
321
Units of ideal gas law
J/K*mol
322
Kinetic energy of n moles of gas
(3/2)nRT
323
When an outside force of other gases causes the molecule to move
Brownian forces
324
Variations from ideal gas conditions occur at _____
low volumes, low temperatures, high pressures
325
The a and b in the van der Waals equation represent...
a = attractive forces between particles, b=volume of molecules themselves
326
Ways of determining the reaction rate
Measure the inital reaction rate for a variety of reactant concentrations. Graph the concentration of the reactants as a function of time. Find the mechanism of the reaction.
327
N+L rule
determines which electron configuration comes next
328
The rate law constant (k) is affected by ____ and ____
temperature, catalyst
329
At equilibrium, the rate of forward reaction is ______ to/than the backwards reaction
equal
330
If a collision between two reactants does not form product, it can be assumed that....
the positioning was not right and or the energy of collision was not sufficient
331
As temperature increases, the range of reaction kinetic energy at which particles collide _____. The rate of collisions peaks at a higher kinetic energy than at _____ temperatures
increases, lower
332
Do liquids and solids enter into equilibrium equations?
no
333
Rates _____ increase with increasing temperature in endothermic and exothermic reactions
always
334
The emissivity value of blackbodies
1
335
Is it more efficient to heat quickly or slowly?
quickly
336
The area under the graph of a P vs. V graph is equal to _____
the work done on the gas
337
Free adiabatic expansion
No change in temperature since no work is done and no heat is lost (ideal gas)
338
Intensive properties
Not dependent on the amount of substance
339
Extensive properties
Dependent on the amount of substance
340
Percent yield
actual/theoretical * 100
341
Rapid cooling of polymers will result in an _____ _____ whle slow cooling will result in _____ _____
amorphous solid, crystalline solid
342
Diamond is a _____ _____ solid
network covalent
343
Heisenburg uncertainty _____ with increasing mass
decreases
344
Enthalpy will differ from the energy of reaction when...
The change in entropy is high
345
Heat of formation
Amount of heat needed to produce 1 mole of a given product
346
The heat of formation of natural elements is...
zero
347
Gases have ____ entropy than phases of the same substance
higher
348
Example of 2nd law of thermo
Heat always flows from high to low T
349
Bond formation is ____thermic and _____ entropy
exo, decreases
350
Any machine is less than ___% efficient because of ____ loss
50, heat
351
A substance at absolute zero has this entropy value
zero
352
Constant ____ and ____ are necessary in classfying spontaneity of reactions
temperature and pressure
353
When the velocity of a fluid increases, the pressure _____
decreases
354
When temperature decreases in a fluid, velocity _____, and thus pressure _____
increases, decreases
355
When density of a fluid decreases, volume ______ and pressure _____
increases, decreases
356
The pressure of a fluid against a flat surface is equal in momentum/∆t*A
Define
357
Ideal fluid characteristics
Volume flow rate is constant, no viscosity, incompressible
358
The flow of an ideal fluid _____ be determined by pressure changes alone
cannot
359
Possible phases of a solution
gas, liquid, solid
360
In ideal dilute solutions, solute particles ____ interact, and the mole fraction of the solvent approaches ____
never, 1
361
Ideal solutions obey _____ law, where solute and solvent molecules are similar and interact with each other similarily
Roults
362
Non crystalline homogenous solution
Colloid
363
Colloids do not ____ in solution. Cannot be _____ by centrifugation. May be separated by addition of _____, by raising_____ or by this technique.
settle, separated, electrolytes, temperature, dialysis
364
Scattering of light results in the appearance of the lights path
Tyndall effect
365
Colloid particles are ____ than thos of the solvent but not _____ enough to precipitate out
larger, large
366
Ionic compound naming:
-ate, -ite, hypo, per
367
conjugate base of carbonic acid
bicarbonate
368
Hydration involves
breaking of water-water H-bonds, formation of water-solute bonds
369
Ion number
number of water molecules that bind to an ion in an aqueous solution
370
If a compound is composed of two non-metals, it must be a ____ compound
molecular
371
A negative heat of solution indicates that the solute-solvent bonds are ____ than the solute-solute bonds.
stronger
372
Sublimation occurs when the vapor pressure of the solid is _____ than the partial pressure above it
greater
373
Boiling occurs when the vapor pressure of the liquid is greater than the _____ pressure above it
total
374
Vapor pressure ____ with temperature
increases
375
Atmospheric gas will condense when ____ pressure of the atmospheric gas is at least as great as the ____ pressure of the liquid at that temperature
partial, partial
376
Steam refers to
water vapor above 100ºC
377
A non volatile solute has ____ vapor pressure
zero
378
Ionic compounds with greater charges are typically ____ soluble in water
less
379
Deviation from Ksp in real exeriments is usually witness because of ____ ____ and _____ reactions that take ions out of solution
ion pairing, hydrolysis
380
Salt crystals nucleate gas bubbles causing them to ____
coalesce
381
Amorphous solids have ____ melting points
poorly defined
382
Gases nearly always form ____ phase(s)
one
383
Coffe cup calorimeters are ____ systems. They are used to measure ____ change.
open, enthalpy
384
When an impurity is added to a solid, its melting point ____ and ____
decreases, broadens
385
When an impurity is added to a liquid, its boiling point ____
increases
386
vant hoff factor for non-electrolytes
1
387
Strong acids
HCl, HNO3, H2SO4, HBr, HI, HClO4
388
Amphiprotic
Capable of gaining or losing a proton
389
Strong acids have ____ conjugate bases
weak
390
Strong bases have ____ conjugate bases
weak
391
weak conjugate acids can have ____ conjugate bases
weak
392
Strong acids are stronger than ____
H3O+
393
Strong bases are stronger than OH-
Define
394
Strong bases
H-, Na2O, N(3-)
395
Percent ionization of an acid is dependent on
temperature, identity of acid, concentration of acid
396
In living organisms, small pH changes can great ____ rate chagnes
large
397
Atoms with unfilled d orbitals can _____ visible light and move to these orbitals, causing brilliant colors in transition metals
absorb
398
HCl reacts with CO3- to form
CO2
399
Reaction involving the exchange of bonds between the chemical species. Often called double-displacement.
metathesis reaction
400
The strength of oxyacids _____ with the central atoms oxidation state
increases
401
Which acid is stronger HClO or HIO
HClO
402
Which acid is stronger HCl or HI
HI
403
Acidity increases with bond _____ and decreases with bond ____ and and increases with the conjugate base ____
polarity, strength, stability
404
A hydride contains ____ and one other element
H-
405
A substance with a high boiling point will generally have a ____ vapor pressure
low
406
The melting point of water will _____ under low pressure
increase
407
The melting point of most liquids will ____ under low pressure
decrease
408
Osmolarity
Π=iMRT
409
Isobaric
constant pressure
410
Pico (p)
Define
411
Alkaline Earths
Group IIa
412
Ksp of MX
Define
413
Homogenous Catalyst
same phase
414
Heisenberg uncertainty principle
Define
415
Reaction Entropy
Define
416
Anode Galvanic
negative electrode (oxidation)
417
Arrhenius equation
k = Ae^(-Ea/RT)
418
Metals
Elements that are characteristically electropositive, malleable, and ductile. These elements tend to be found on the left side of the periodic table, be lustrous, and have relatively low ionization energies and electron affinities
419
Halogens
Contains nonmetals, 7 valence electrons in it's outermost energy level. Very reactive
420
Equilbrium constant is ______ dependent
temperature
421
Effusion
one gas moves through air
422
Standard delta G =
-RT ln(K)
423
Titration
Analytical procedure in which a solution of known concentrations is slowly added to a solution of unknown concentration to the point of molar equivalency, thereby providing the concentration of the known solution.
424
Valence Electrons
Electrons occupying the outermost electron shell of an atom, participating in chemical bonds. Atoms with the same number of valence electrons tend to have similar properties (families in the Periodic Table).
425
STP
Standard Temperature and Pressure. 273 Kelvin (0 Celsius), 1 atmosphere (760 torr, 760 kPA).
426
Lewis definition
Acids defined as electron-pair acceptors and bases as electron-pair donors.
427
Half-Cell
An electrode immersed in an electrolytic solution that is the site of either oxidation or reduction in a galvanic (voltaic) cell
428
Noble gases
Contains nonmetals that are non-reactive. Full outermost energy level except helium which has 2.
429
Which group is most metallic?
group IA
430
Best insulator has _____ specific heat
highest
431
Speed is dependent on
kinetic energy only
432
autophagy
A process that describes lysosomes using their hydrolytic enzymes to recycle the cell's own organic material
433
Phase Diagram
A pressure vs. temperature plot showing the conditions under which a substance exists in equilibrium between different phases or in which the substance exists in pure phase.
434
Ion
A single or polyatomic particle with an electric charge.
435
Normality
Gram equivalent weight of solute per liter of solution, often denoted by N.
436
Disproportionation
Redox reaction, in which the same species is both oxidized and reduced.
437
Faraday's Constant
Denoted by F, 9.65x10⁴ coulombs/mol e⁻. Commonly used in the formula It = nF (I = Current, t = time (s), n = mol e⁻).
438
Indicator
A chemical species that changes color during dissociation, used to signal the end point of a titration.
439
Redox Half-Reaction
Hypothetical equation showing only the species that is oxidized or reduced in a redox reaction and the correct number of electrons transferred between the species in the complete, balanced equation.
440
solvation
sol, a chemical process in which solvent molecules and molecules or ions of the solute combine to form a compound
441
Amphoteric
having characteristics of both an acid and a base and capable of reacting as either
442
System
The part of the universe under consideration that is separated by some real or imaginary boundary from its surroundings
443
atomic radius
one-half the distance between the nuclei of two atoms of the same element when the atoms are joined
444
mole
the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon-12
445
bond length
the average distance between the nuclei of two bonded atoms
446
Decomposition reaction
(chemistry) separation of a substance into two or more substances that may differ from each other and from the original substance C>>>>A+B
447
Molecule
The smallest unit of a substance, composed of two or more atoms joined in covalent bonds, that still retains all the chemical properties of that substance
448
What charge to group 7 elements have?
-1
449
Entropy when energy into the system
entropy increases
450
Chemical Similarity between Atoms
stay in same column
451
Money Metals
Au, Pt, Ag, Hg, Cu, Ni
452
Speed of light (c)
3.00 x 10^8 m/s
453
Ion Product
Product of the molar concentrations of dissociated ions in solution at any point in the reaction other than equilibrium or saturation, where each ion is raised to the power of its stoichiometric coefficient. Denoted IP.
454
Molecular Formula
A chemical formula showing the actual number of atoms present in a certain compound.
455
Limiting Reagent
Reactant of a chemical equation that, given nonstoichiometric amounts, determines the amount of product that can form; the reactant that runs out first.
456
Molecular Orbital
Region in a molecule where atomic orbitals overlap, resulting in either a stable low-energy bonding orbital or an unstable high-energy antibonding orbital.
457
Acid Dissociation Constant
An equilibrium expression used to measure weak-acid strength, given by the ratio of the product of the products' molar concentrations to the product of the reactants' molar concentrations, with each term raised to the power of its stoichiometric coefficient. Denoted Ka.
458
Percent Yield
A ratio, calculated as a percentage, of the actual mass of product yielded to the theoretical yield of product mass.
459
Isothermal System
Process in which the system either gains or loses energy to maintain a constant temperature.
460
Isolated System
A system that can exchange neither energy nor matter with its surroundings.
461
Exothermic
A reaction that proceeds with the net release of energy (heat) into the surroundings
462
Magnetic quantum number
specifies the specific orbital in which the electron is most likely to be found., Third quantum number, designated as ml. Describes a particular orbital within a subshell where an electron is very likely to be found. Possible values are integers in the -1 to 1 range, including 0.
463
Chemical Kinetics
the area of chemistry that is concerned with reaction rates and reaction mechanisms
464
Rate-Determining Step
The slowest step in a reaction mechanism that determines the overall rate of the reaction
465
Effective Nuclear Charge
The resulting positive nuclear charge an outer electron senses after accounting for the shielding effect of inner core electrons. Abbreviated Zeff. Increases from left to right and from bottom to top on the periodic table
466
Concentration
tells you how much solute is present compared to the amount of solvent
467
Theoretical Yield
The expected amount of product yielded in a reaction according to reactants' stoichiometry
468
Atomic Absorption Spectrum
The spectrum of certain absorbed wavelengths of light corresponding to an atom's spectrum of emitted frequencies of light
469
Diamagnetic
An atom or a substance that contains no unpaired electrons and is consequently repelled by a magnet
470
Reduction
A reaction in which a species gains electrons
471
Atom
The basic building block of all matter in the universe. An atom is made up of three main components: protons, neutrons, and electrons
472
Azeotrope
A liquid mixture of two or more substances that has a constant boiling point greater than or less than the boiling points of its constituents. The vapor of this unique mixture has the same composition as the liquid state, making difficult to separate the constituents
473
Isochoric Process
A process in which volume remains constant and no net pressure-volume work is done
474
Bronsted-Lowry definition
Common definition of acids as proton (H+) donors and bases as proton acceptors
475
What charge do group 1 elements tend to have?
1
476
From left to right across the periodic table, metallic characteristics (increase, decrease)
decrease
477
Metallic Solid
molecules held in place by delocalized bonding
478
State Variables
conditions that must be specified to establish the state of the system, pressure, volume, temperature, and amounts of substances
479
Glavanic Reduction potentials
higher reduction potential is cathode, lower reduction potential is anode
480
Hund's Rule
orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin
481
Combination Reaction
A reaction in which two or more reactants combine into a single product.
482
Pfund Series
Set of spectral lines resulting when a hydrogen atom undergoes a transition from energy levels n>5 to n=5.
483
Ideal Gas Law
A unification of Boyle's Charles and Gay-Lussac's, and Avogadro's Principle into a formula describing the behavior of ideal gases: PV=nRT. (Pressure*Volume = moles*Avogadro's Number*Temperature).
484
Reducing Agent
A species that is oxidized in the process of reducing another species.
485
Lyman Series
Set of spectral lines appearing in the UV region when a hydrogen atom undergoes a transition from energy levels n>1 to n=1.
486
Specific Heat
The amount of heat required to raise one gram of a substance by 1 degree Celsius;heat capacity
487
Molality
The ratio of the number of moles of solute dissolved in one kilogram of solvent. molality (M = moles solute/kg of solution)
488
Adiabatic Process
A process in which no heat is transferred to or from the system by its surroundings
489
Reaction Quotient
A ratio of the concentrations of the products to the concentrations of the reactants at any point during the reaction aside from equilibrium, where each reactant and product in the expression is raised to the power of its stoichiometric coefficient. Commonly denoted by Q
490
Solution Equilibrium
When a solute is dissolved in a solvent, it will dissociate until reaching an equilibrium point at which the rate of dissociation equals the rate of precipitation of the solute, regardless of any additional solute introduced into the mixture
491
Solubility Product Constant
Product of the molar concentrations of dissociated ions in solution at saturation, where each ion is raised to the power of its stoichiometric coefficient. Denoted Ksp.
492
Aqueous Solution
a solution in which water is the solvent
493
Hydrogen Bonding
involves lone pairs of electron on an electronegative atom of one molecule and a polar bond to hydrogen in another rmolecule. They are confined tomolecules that contain O, N, and F atoms
494
Ideal Gas Postulates
1. molecules are very small compared to the distance between them.2. molecules are constantly moving3. pressure of the gas- collisions of the molecules with container walls4. molecules do not experience intermolecular forces5. KEave proportional to T
495
Osmotic Pressure
increase in pressure due to a solvent crossing a membrane into a more concentrated solution ΠV = nRT
496
Assumptions of ideal gases
No volume, no repulsive forces, elastic collisions, kinetic energy is proportional to temperature
497
Le Châtlier's Principle
When a system in equilibrium is placed under one of several stressors, it will react in order to regain equilibrium. In other words, act on a system, it will work back towards equilibrium.
498
Vapor Pressure
Partial pressure of a vapor when it is in equilibrium with its solid or liquid phase.
499
Single Displacement Reaction
Chemical reaction in which an atom or ion of one compound is replaced by another atom or ion.
500
Balmer Series
A set of spectral lines that appear in the visible light region when a hydrogen atom undergoes a transition from energy levels n>2 to n=2
