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concept 5e Flashcards

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1
Q

system

A

the matter that is being observed in the reaction

total amount of reactants and products in a reaction

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2
Q

surroundings

A

or environment

everything outside of the system

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3
Q

categories of systems

A

isolated
closed
open

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4
Q

isolated system

A

system cannot exchange energy or matter with the surroundings
exp. an insulated bomb calorimeter

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5
Q

closed system

A

system can exchange energy but cannot exchange matter with the surroundings
exp. a steam radiator

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6
Q

open system

A

system can exchange both energy and matter with the surroundings
exp. pot of boiling water

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7
Q

energy exchange

A

exchange of heat and work

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8
Q

process

A

when a system experiences a change in one or more of its properties it undergoes a process
change can be change in concentrations of reactants or products, temp, or pressure
a change of the state of a system

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9
Q

First law of thermodynamics

A

states that the total energy of a system and its surroundings remains constant
delta U=Q-W
delta U is the change in internal energy of a system, Q is heat added, W is work done by the system

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10
Q

isothermal process

A

occur when the system’s temperature is constant
constant temp implies that total internal energy is constant
forms a hyperbolic curve on pressure-volume graph

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11
Q

adiabatic process

A

occur when no heat is exchanged b/w the system and the environment
thermal energy is constant
the change in internal energy is equal to work done on the system (-W)
hyperbolic curve on P-V graph

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12
Q

isobaric process

A

occur when the pressure of the system is constant
appears as a flat line on a P-V graph
common bc it is easy to control temp and pressure

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13
Q

isovolumetric (isochoric) process

A

processes experience no change in volume
the gas neither expands or compresses, no work is performed
the change in internal energy is equal to the heat added to the system
appears as a vertical line on P-V graph

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14
Q

spontaneous process

A

process that occur by itself without having to be driven by energy from an outside source
may not happen quickly and may not go to completion
tend to have high activation energies

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15
Q

state functions

A

properties that describe the system in an equilibrium state
cannot describe the process of the system, how the system got to its equilibrium
useful for comparing one equilibrium state to another

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16
Q

process function

A

the pathway taken from one equilibrium state to another described quantitatively
most important are work (W) and heat(Q)

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17
Q

values of state functions

A

pressure (P), density, temp (T), volume (v), enthalpy (H), internal energy (U), Gibbs free energy (G), entropy (S)
“when I’m under Pressure and feel Dense, all I want to do is watch TV and get HUGS”

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18
Q

standard conditions

A

standard values defined for measuring the enthalpy, entropy, and Gibbs free energy of a reaction
298 K (25 deg C)
1 atm pressure
1 M concentrations
dont confuse with standard temperature and pressure (STP)

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19
Q

standard temperature and pressure

A

temp is 273 K (0 deg C)
pressure is 1 atm
used for ideal gas calculations

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20
Q

standard state

A 
most stable form of a substance under standard conditions 
H2(g)
H20(l)
NaCl(s)
O2(g)
C(s,graphite)
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21
Q

phase diagrams

A

graphs that show the standard and nonstandard states of matter for a given substance in an isolated system
determined by temperatures and pressures

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22
Q

phase changes

A

changing between a solid, liquid, or gas
are reversible
an equilibrium of phases eventually reached at any given combination of temperature and pressure

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23
Q

phase equilibria

A

analogous to the dynamic equilibria or reversible chemical reactions
the concentrations of reactants and products are constant bc the rates of the forward and reverse reactions are equal

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24
Q

liquid phase

A

molecules are relatively free to move
some near the surface have enough kinetic energy to leave the liquid phase
each time a particle leaves the temp of remaining liquid decreases

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25
evaporation
vaporization the process where a molecule has enough kinetic energy to leave the liquid phase and escape into the gaseous phase endothermic process
26
boiling
specific type of vaporization that occurs only under certain conditions rapid bubbling of the entire solution with rapid release of liquid to gas particles only occurs above the boiling point of a liquid and involves vaporization through the entire volume of liquid
27
condensation
in a covered container the escaping molecules are trapped above the solution these molecules exert a countering pressure this forces some of the gas back into the liquid phase facilitated by lower temperature or higher pressure
28
boiling point
the temp at which the vapor pressure of the liquid equals the ambient pressure ambient pressure --> external, applied, incident, atmospheric
29
gas-liquid equilibrium processes
evaporation vaporization boiling condensation
30
solid phase
atoms or molecules are confined to specific locations each molecule can undergo motions about some equilibrium position these vibrations increase when heat is applied
31
energy microstates
availability of these microstates increase as the temperature of the solid increases that molecules have greater freedom of movement and energy disperses
32
fusion
or melting process from solid to liquid molecules in solid phase absorb enough energy the 3D structure of the solid will break down, and the molecules will escape into liquid phase temperature at which this occurs is called melting point
33
solidification
or crystallization, or freezing reverse process of fusion from liquid to solid temperature at which this occurs is called the freezing point
34
liquid-solid equilibrium processes
fusion or melting | solidification, crystallization, or freezing
35
sublimation
``` when solid goes directly into the gas phase dry ice (solid CO2) sublimes at room temperature and atmospheric pressure ```
36
deposition
reverse transition | from gaseous to solid phase
37
cold finger
device used in organic chem labs used to purify a product that is headed under reduced pressure this causes it to sublimate
38
gas-solid equilibrium processes
sublimation | deposition
39
lines of equilibrium
phase boundaries lines on a phase diagram indicate the temperature and pressure values for the equilibria b/w phases divide the diagram into 3 regions corresponding to the 3 phases represent phase transformations
40
triple point
point at which the three phase boundaries meet | this is the temp and pressure at which the three phases exist in equilibrium
41
critical point
the phase boundaries terminate at this point | this is the temp and pressure above which there is no distinction between phases
42
temperature (T)
related to the average kinetic energy of the particles of a substance was we scale how hot or cold something is
43
thermal energy (enthalpy)
average kinetic energy of the particles in a substance | we must also consider how much substance is present
44
heat (Q)
transfer of energy from one substance to another as result of their difference in temperature
45
heat vs temp
heat is specific form of energy that can enter or leave a system temp is a measure of the average kinetic energy of the particles in a system
46
zeroth law of thermodynamics
implies that objects are i thermal equilibrium only when their temperatures are equal this leads us to believe that heat is a process function not a state function
47
endothermic process
system absorbs heat | delta Q>0
48
exothermic process
system releases heat | delta Q<0
49
unit of heat
joule (J) or calorie (cal) | 1 cal=4.184 J
50
enthalpy
delta H | equivalent to heat (Q) under constant pressure
51
calorimetry
process of measuring transferred heat 2 basics types: constant-pressure calorimetry and constant-volume calorimetry coffee-cup calorimeter is an example of a constant-pressure calorimeter bomb calorimeter is an example of a constant-volume calorimeter
52
q=mc(deltaT)
the heat abosorbed or released in a given process | m is mass, c is specific heat of substance, delta T is the change in temp (in C or K)
53
specific heat
the amount of energy required to raise the temperature of one gram of a substance by one degree C or one K generally provided on test day except for H20-->c(H20)=1cal/g*K
54
heat capacities
product m*c (mass times specific heat)
55
constant-pressure calorimeter
an insulated container covered with a lid and filled with a solution in which a reaction or some physical process is occurring pressure is constant and temperature is measured as run progresses should be sufficient thermal insulation to ensure accurate temp measurement coffee-cup calorimeter
56
bomb calorimeter
or decomposition vessel type of constant-pressure calorimetry sample of matter placed in steel decomposition vessel filled with almost pure oxygen gas vessel then placed in insulated container holding mass of water vessel ignited with electric ignition mechanism material combusts in presence of oxygen and the heat that evolves is the heat of the combustion reaction
57
heating curves
as compound is heated, the temp rises until melting or boing point the it remains the same as compound is converted to next phase once sample is converted the temp begins to rise again until reaches next transition phase phase change regions do not change temp is we can't use q=mc(deltaT)
58
q=mL
used to find q during phase change | m is mass and L is latent heat, general term for enthalpy of isothermal process
59
enthalpy (H)
heat content of a system at constant pressure state function, so we can calculate the change in enthalpy for a system undergoing a process done by subtracting the H of reactants from the H of products positive delta H is an endothermic process, negative delta H is exothermic process
60
standard enthalpy of formation
the enthalpy required to produce one mole of a compound from its element in their standard states
61
standard heat of a reaction
enthalpy change accompanying a reaction being carried out under standard conditions calculated by taking the difference b/w the sum of the standard heats of formation for the products and subtracting the sum of standard heats of formation for the reactants
62
Hess's Law
states that enthalpy change of reactions are additive when thermochemical equations are added to get net equation the corresponding gets of reaction are also added to give net heat of rxn applies to all state functions including entropy and Gibbs free energy
63
bond dissociation energies
the average energy this si required to break a particular type of bond between atoms in the game phase endothermic process
64
bond formation energies
opposite of bond breaking has the same magnitude of energy but is negative rather than positive energy is released when bonds are formed
65
standard heat of combustion
enthalpy change associated with the combustion of a fuel
66
entropy
measure of the spontaneous dispersal of energy at a specific temperature how much energy is spread out, or how widely spread out energy becomes in a process deltaS=Qrev/T Qrev is the heat that is gained or lost, T is the temp in K when heat is distributed into a system entropy increases, when it is distributed out of a system entropy decreases
67
second law of thermodynamics
states that all spontaneous processes lead to an increase in entropy of the universe energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so
68
Gibbs free energy (G)
the energy of a system available to do work change in free energy can be determined for a given reaction equation from the enthalpy change, temperature, and entropy change negative deltaG denotes spontaneous reaction, positive deltaG denotes a nonspontaneous reaction
69
exergonic process
movement toward equilibrium position decrease in Gibbs free energy (deltaG<0) spontaneous
70
endergonic process
movement away from the equilibrium position increase in Gibbs free energy (deltaG>0) nonspontaneous
71
effects of deltaH, deltaS and T on spontaneity
+deltaH, +deltaS-->spontaneous at high T +deltaH, -deltaS-->nonspontaneous at all T -deltaH, +deltaS-->spontaneous at all T -deltaH,-deltaS-->spontaneous at low T
72
standard free energy
free energy change of reactions measured under standard conditions can be derived from standard free energy of formation of products minus reactants can also be determined from equilibrium constants
73
deltaGrxn=-RTlnKeq
free energy from equilibrium constant R is ideal gas constant, T is temp in K, Keq is the equilibrium constant evaluate free energy change and spontaneity of reaction higher Keq, more positive ln, more negative standard free energy change, more spontaneous
74
collision theory of chemical kinetics
states that the rate of a reaction is proportional the the number of collisions effective collision, that leads to product, occurs only if the molecules collide w/ each other in correct orientation w/ enough energy rate=Z*f (Z is # of collisions, f is fraction of effective collisions)
75
reaction concentrations
greater concentration of reactions, the greater the number of effective collisions, increase in reaction rate direct relationship
76
temperature
reaction rate will increase as the temperature increases all reactions are temperature dependent and experience an optimal temperature for activity usually b/w 30-40 deg C if it exceeds optimal temp the catalyst will denature and the rxn rate plummets
77
solvent
generally polar solvents are preferred their dipole polarized the reactants weakening the bonds and this allows for effective collisions and a reaction to occur
78
catalyst
increase the reaction rate by decreasing the activation energy for the reaction are not consumed in the reaction
79
rate law
for forward irreversible rxns the rate is proportional to the concentration of reactants aA+bB-->cC+dD rate=k[A]^a[B]^b
80
zero-order reaction
reaction in which the rate of formation of product C is independent of changes in concentration of any of the reactants rate=k[A]^0[B]^0=k on graph of concentration vs time curve, slope is linear and opposite (negative) of k
81
first-order reaction
rate that is directly proportional to only one reactant rate=k[A] or rate=k[B] shows a nonlinear graph of conc. vs time, this proves that the rate of formation of products is dependent on conc. of reactant plotting ln[A] vs. time reveals a linear graphe where k=-slope
82
second-order reaction
rate that is proportional to either the concentrations of 2 reactants or the square f the concentration of a single reactant rate=k[A][B] or rate=k[A]^2 or rate=k[B]^2 nonlinear graphs for conc. vs. time, decreases faster than first order (greater exponential decline) if plot 1/[A] vs. time get a linear graph where k=slope
83
higher-order reaction
very few noteworthiness reactions bc it is far more rate for 3 particles to collide simultaneously with correct orientation and sufficient energy to undergo a reaction
84
mixed-order reaction
non-integer orders (fractions) a reaction in which the reaction order changes over time in the rate law more specifically described as broken-order
85
law of mass action
states that if the system is at equilibrium at a constant temperature than the ratio is constant: Keq=[C]^c[D]^d/[A]^a[B]^b=[prod]/[react] defines the position of equilibrium
86
at equilibrium...
the rate of the forward reaction equals the rate of the reverse reaction entropy is at a maximum Gibbs free energy is at a minimum linking thermodynamic and kinetics
87
reaction quotient (Q)
has the same form as the equilibrium constant but the concentrations of products and reactants may not be at equilibrium when compared to Keq it dictates the direction a reaction will proceed spontaneously Q=[C]^c[D]^d/[A]^a[B]^b=[prod]/[react] during any point in the reaction
88
comparing Keq and Q
QKeq: reaction proceeds in the reverse direction toward reactants, greater conc. of products, deltaG>0 (endergonic)
89
properties of law of mass action
1. concentrations of pure solids and pure liquids don't appear in equil. constant expression 2. Keq (equil. constant) is temperature-dependent 3. the larger Keq, the equilibrium position lies further to the right (toward products), so more products at equil. than reactants 4. if equilibrium constant for rxn in one direction is Keq, for the reverse reaction is is 1/Keq
90
Le Chatelier's principle
states that if a stress is applied to a system the system shifts to relieve that applied stress reaction is temporarily moved out of equilibrium state and responds by reacting in whichever direction to reestablish equilibrium
91
change in concentration
add reactants: rxn shifts to right, toward product add products: rxn shifts to left, toward reactant remove reactants: rxn shift to left, toward reactant remove products: rxn shift to right, toward product
92
change in pressure
pressure increased: vol decreases, rxn shifts to right toward products pressure reduced: vol increases, rxn shifts to left toward reactant
93
change in volume
volume increased: pressure decreases, rxn shifts to left toward reactant volume decreased: pressure increases, rxn shift to right toward product
94
change in temperature
temperature increases: rxn shifts to left toward reactants | temperature decreases: rxn shifts to right toward products
95
kinetic product
the product of a reaction that is formed favorably at lower temperature bc thermal energy is not available to form the transition state required to create a more stable thermodynamic product has smaller overall different in free energy b/w products and reactants than thermodynamic product called the "fast" product minor product
96
thermodynamic product
the product of a reaction that is formed favorably at a higher temperature because thermal energy is available to form the transition stat of the more stable product has a larger overall different in free energy than kinetic product, but is the more stable product bc is has a lower free energy major product

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