125 terms

Chem 4

Chem 4
What do we know about molecules in the same phase?
1) They are in the same state (liquid, solid, gas)
2) They have the same chemical composition (i.e. molecular formula is the same)
3) They are structurally homogenous (they have the same molecular structure...not isomers)
In its solid form, Sulfur forms S8 and S6 rings, among others. Are these in the same phase?
No...although they are both in the same state (both are solid) and they both have the same empirical formula (S or S2) they do NOT have the same molecular structure
Do different phases of the same substance have the same specific heat capacity?
No...1 substance in its solid or liquid state will have a much lower specific heat than the same substance in its gas state bc in its gas state a lot of the heat put into the system does work rather than increasing the temperature of the substance
When you see melting point/freezing point, what 4 things should you immediately think?
1) In general
2) Ex- Higher melting point
1) General:
-does it take more/less heat/energy to melt it
-does it melt at a higher or lower temperature
-if the temperature of the solution is in between its melting the point and the melting point of what you're comparing it to, which will be a solid and which will be a liquid
-Why would it have a higher/lower mp/bp than the other thing? 2) More heat/energy is required to make it melt...it melts at a higher temp...it stays in its solid state when the substance it is compared to is in its liquid state...greater intermolecular forces...shape...MW
When you see boiling point/condensation point, what 5 things should you immediately think?
1) In general
2) Ex.- Lower boiling point
1) General:
-does it take more/less heat/energy to boil it?
-does it boil at a higher/lower temp.?
-does it have a higher/lower vapor pressure at the same temp.?
-if the temp. of the solution is btwn its boiling pt. and the boiling pt. of the thing you're comparing it to, which is a liquid and which is a gas?
-Why would it have a higher/lower melting point than the other thing?
2) Less heat/energy is required to make it vaporize...it vaporizes (boils) at lower temperatures...it has a higher vapor pressure at the same temperature (why?)...this will be in the gas state when the thing it is compared to is in its liquid state...weaker intermolecular forces...shape...MW
ΔH of fusion...how does the temperature change at this point?
-For solid -> liquid ΔH is positive bc it indicates the amnt. of energy needed to be put IN to increase convert from solid state to liquid state
-For liquid -> solid ΔH is negative bc it represents the amnt. of energy that is removed when the substance converts from its liquid state to its solid state
-Temp. does NOT change bc the energy goes into breaking the strong intermolecular interactions rather than increasing the average kinetic energy of the molecules...once the phase has changed, any further energy added will increase the kinetic energy of the molecules and thus raise the temp.
ΔH of vaporization...how does temp. change at this point?
-For gas-> liquid ΔH is negative bc it represents the amnt. of energy removed when it converted from gas to liquid
-For liquid -> gas ΔH is positive bc it represents the amnt. of energy needed to be added in order to convert from liquid to gas
-Temp. does NOT change bc the energy goes into breaking the strong intermolecular interactions rather than increasing the average kinetic energy of the molecules...the DEGREE OF ORGANIZATION of the solution changes...once the phase has changed, any further energy added will increase the kinetic energy of the molecules and thus raise the temp
How does the temp. of the air just above a liquid compare to the temp. of the liquid while the liquid is boiling?
-the temperature of both will be equal bc the liquid is currently undergoing a phase change and thus all the heat being added into the system is being used to break intermolecular interactions and is not being used to increase the velocity of the individual molecules...even if some heat did increase the velocity of the particles, that increase in kinetic energy would immediately be dissipated from its collisions with other particles, including the liquid, thus resulting in increased disruption of intermolecular forces (i.e. phase change)
Why do things boil? What things would you expect to increase boiling point?
-vapor pressure is a measure of how much pressure the molecules at the surface of a liquid are exerting on the atmosphere...it represents how much the liquid molecules want to leave the liquid and get into the gas phase
-as you increase the temp., you increase the kinetic energy of the particles and thus they want to get out of the liquid phase even more and they have more force to do so...thus, the vapor pressure increases
-as soon as the vapor pressure exceeds the ambient pressure (atmospheric pressure or whatever other air pressure is exerted on it) the substance boils
-Things that increase bp: increased ambient pressure...higher intermolecular forces...increased mass (same energy results in lower velocity of individual particles)...adding a non-volatile solute...temperature does NOT raise bp (temp. brings it to its boiling point but it doesn't RAISE its boiling point)
What are the following phase changes called?
A) Solid -> Liquid
B) Gas -> Liquid
C) Solid -> Gas
D) Liquid -> Gas
E) Gas -> Solid
F) Liquid -> Solid
A) Melting
B) Evaporation
C) Sublimation
D) Boiling
E) Deposition
F) Freezing
What does volatile mean? Non-volatile?
Volatile means that a substance's vapor pressure exceeds atmospheric pressure at normal temperatures...thus, at very low temperatures, the substance boils...non-volatile is the opposite
Formula for the effects on boiling point from adding a non-volatile solute
ΔT = (kb)(m)(i) ...kb = boiling point constant of the solvent (units = Kelvin/molal)...m = molality (mol solute/mass of solvent in kg)...i = # of things present when it dissolves in solution (e.g.- glucose would be i = 1...Pb(NO3-)2 would be i = 3)
Why is the boiling point elevated when you add non-volatile solute?
-when you add a non-volatile solute to a solution, some of it is dissolved in the liquid deep down but some of it is dissolved and hangs out at the surface...remember that things boil bc their vapor pressure exceeds atmospheric pressure...you can imagine vapor pressure as a force exerted on the atmosphere due to the "push" of tons of molecules in the liquid state with a high kinetic energy...the more molecules you have "pushing" together near the surface, the higher the vapor pressure will be...the non-volatile solute decreases bp by hanging out at the surface of the liquid...bc the solute particles are also at the liquid's surface, there are fewer solvent particles that can fit at the surface to help "push" against the atmosphere...this is why vapor pressure is decreased and that is why the boiling point is increased
m = (mol of solute)/(mass of solvent)...be careful to notice that the denominator is mass of SOLVENT only...not the entire solution
Formula for the effects on freezing point from adding a non-volatile solution
ΔT = (kf)(m)(i)...kf = freezing point constant of the solvent (units = Kelvin/molal)...m = molality (mol solute/mass of solvent)...i = # of things present when it dissolves in solution (e.g.- glucose would be i = 1...Pb(NO3-)2 would be i = 3)
Why is the freezing point depressed when you add non-volatile solute?
if you add a non-volatile solution to any liquid it will lower the freezing point (meaning that it will require MORE energy to be removed from the system in order to freeze it...this is the same thing as saying it has a lower melting point meaning that it will melt at lower temp. bc LESS energy is required to be put in to break intermolecular forces)...the reason the freezing point lowers is bc the solute that is introduced decreases the ability of the solvent particles to interact and form intermolecular forces...intermolecular forces are the most stable (and the most organized) in the solid state...when something else (the solute) is inserted, it disrupts the intermolecular forces and thus the solvent can not have these intermolecular interactions with itself as easily as before...thus, more energy must be removed to slow down the particles EVEN more, so that the intermolecular interactions can be strong enough to move the solvent into its solid state
What determines the phase (solid, liquid, gas)? What determines the temperature?
-phase is determined by the degree of organization...at solid phase, the degree of organization is the highest...at liquid phase, the degree of organization is slightly less (fewer intermolecular interactions occur, and particles have more freedom to move around)...at gas phase, the degree of organization is the lowest!
-temperature is determined by the average kinetic energy of the individual molecules...KE = 3/2(k)(T)
Osmotic pressure (Π)
-a measure of the tendency of water to move from a high water concentration to a low water concentration (usually across a semi-permeable membrane)
-the side that will receive the water has the HIGHER osmotic pressure...thus, the lower concentration of water has the higher osmotic pressure
A cell is put into a hypertonic solution. Which side of the cell has the higher osmotic pressure.
The extracellular side has the higher osmotic pressure bc the extracellular side has the higher solute concentration (lower water concentration)...thus, water will flow out of the cell
Formulas for Π
Π = iMRT...i = # of particles once dissolved...M = molarity...R = gas constant (.0821 Latm/(K*mol)...T = absolute temp. (kelvin)
Π = (pgh1) - (pgh2)
How do you calculate ΔH from a heating curve?
ΔH of vaporization = the heat added to change from liquid to gas...just measure the portion of the x-axis for the flat line btwn the liquid and gas phases
ΔH of fusion = the heat added to change from solid to liquid...measure the portion of the x-axis for the flat line btwn the solid and liquid phases
At constant pressure Q = ΔH, so basically, ΔH is equal to the x-axis on one of the heating curve graphs, including the sloped sections
Draw a phase diagram
-Solid is the upper left portion...then liquid, then gas
-Pressure is on the y-axis...and temp. is on the x-axis
Under what conditions is sublimation possible? Under what conditions is deposition possible?
Sublimation happens when a solid converts DIRECTLY into gas...this happens at low temperature and low pressure...deposition ALSO happens at low temperature and low pressure; however, it happens by increasing the pressure or decreasing the temperature as opposed to sublimation which results from a decrease in pressure or an increase in temp. (at super low pressure)
What is the triple point?
This is the point at which the 3 phases intersect...to the left of this point on the graph is where sublimation and deposition can occur
What is the critical point?
It is the point on the graph at which critical pressure and critical temperature overlap...it is a point at the end of the line that separates the gas phase from the liquid phase...beyond this point (at higher temperatures and/or higher pressures) the substance is called supercritical fluid (plasma)
Supercritical fluid
-anything above the critical point
-has both properties of liquid and gas...it can dissolve substances in the same way liquids can...it can effuse through solids the same way that gases can
-the density of a supercritical fluid can be changed a ton by small differences in temperature and/or pressure
Critical Temp.
-gases become more difficult to liquify above certain temperatures bc of the high kinetic energy of the particles
-above critical temp., no matter how much pressure is exerted, the gas can not be turned into NORMAL liquid...instead, it converts into a semi-liquid, semi-gas state called supercritical fluid
Critical Pressure
-the minimum pressure required to liquify a gas at its critical point
-due to the high temperature of the gas, this liquification is actually just a conversion of the gas into its supercritical fluid state
What phases are (or are not) at the...
A) Triple point
B) Critical Point
A) Gas, liquid, solid...NOT supercritical fluid
B) Gas, liquid, supercritical fluid...NOT solid
-a homogenous mixture of two or more distinct compounds that are all in the same phase
-a mixture of various substances that are all in the liquid phase can be considered a solution...ALSO, a solution can be a mixture of various substances that are all in the GAS phase
Solvent, Solute, Solution
Solute is what you have less of...solvent is what you have more of...solution is the mixture of both (if they are both in the same phase)
-something is dissolved in an aqueous solution due to water's intermolecular interactions with it...in order for something to dissolve in water, it must have greater intermolecular interactions with the water than it does with itself...or, as is the case with most salts, the increase in entropy can supply the energy needed to make an endothermic reaction spontaneous
Hydration number
the number of molecules of water with which an ion can interact with in an aqueous solution of given concentration
-contains water
How would you expect the hydration number to change with...
A) Different concentrations of the ion
B) Different sized ions
Common Ion effect
You add compound AB to a solution and it dissolves into A and B...you then add compound BC to the solution and it dissolves into B and C. Bc you have added more B to the solution, this increases the ion product and causes the equation to run to the left according to Le Chatelier's principle...thus, adding common ions to a solution DECREASES solubility of the compound
Definition of a salt
a neutral compound made up of 2 or more IONS (permanently charged species)
Why do things dissolve?
-the intermolecular forces between the solute and the solvent determine whether or not something will dissolve
WATER examples:
-if the intermolecular forces are STRONGER btwn the solvent (water) and the solute (e.g. NaCl) than btwn water and water, then 100% of it will dissolve...however, saturation can still be reached at which point no more solute can be dissolved
-if intermolecular forces btwn solvent (water) and solute molecules (alcohols) are relatively equal to forces btwn water and water, then different amnts. of it will dissolve depending on how large the solute is (ethanol dissolves better than heptanol)...also, depending on how much the entropy increases when the solute dissolves will determine whether or not the solute spontaneously dissolves or not
-if intermolecular forces btwn. the solvent (water) and solute molecules (alkanes) are weaker than the intermolecular forces btwn water and water molecules, then it will not dissolve
What kinds of functional groups would a molecule need to be soluble in water?
1) ionic...ionic forces are stronger than H-bonds and thus ALL ionic compounds dissociate in water bc the water would rather interact with the extremely charged ions than H-bond with other water molecules
2) Alcohols (need abt. 1 -OH for every 5 carbons in order to be soluble)...this is bc it can H-bond...other elements capable of forming H-bonds are fluorine and nitrogen
3) Any oxygen or Nitrogen group (carbonyls; amines; amides)
Why do micelles form?
-bc water molecules can not interact with hydrophobic molecules, the water molecules are forced to take up a super structured form which is energetically unfavorable due to the low entropy...thus, the hydrophobic portions turn in and form a micelle in order to decrease the area that interacts with the water..this is energetically favorable bc of the increase in entropy due to the fewer water molecules that highly ordered near the surface of the hydrophobic molecules
What is a colloid?
-colloids are NOT solutions
-colloids are solvents that contain undissolved solute particles that are too small to be separated by filtration, but are much larger than the dissolved solute particles in a true solution
-the undissolved solute particles ARE still evenly dispersed throughout the solvent so it LOOKS like a solution
What are some properties of colloids?
-colloids scatter light whereas solutions do NOT
-examples of colloids= paint; dust in the air; etc.
(moles solute)/(liters of solution)
(moles solute)/(kg of solvent)
Why does molarity change with temperature but molality does not?
Molarity is a measure of moles/liter...thus, when temperature is increased, the volume will increase and this will decrease molarity...however, molality is a measure of moles/kg and the mass of an object does NOT change with temperature
Mole fraction
(moles solute)/(total moles solution) ...total moles of solution = moles of solute + moles of solvent
Mass percent
(mass of solute)/(mass of solution)*100
ppm...can you guess what ppb is?
(mass of solute)/(total mass solution)10^6...ppb = same10^9...parts per million does NOT mean how many solute particles there are per 1 million total particles...it is simply the mass percent *10^4...the purpose of multiplying by 1 million is to make extremely small concentrations easier to work with
N = (# of equivalents)/(liters solution) ....# of equivalents = the # of hydrogen ions the solute would produce OR accept...ex- the number of equivalents HCl can produce is 1 and thus, molarity = normality for HCl...however, the number of equivalents that H2SO4 produces is 2...thus, 2*molarity = normality...NaOH = 1 equivalent...CuSO4 = 2 equivalents
What is required to create a solution?
-intermolecular interactions between the solute molecules must be broken
-intermolecular interactions btwn the solvent molecules must be broken
-finally, new intermolecular interactions are formed between the solute and solvent particles
What can you tell about the ability of something to dissolve based on the enthalpy value?
-ΔH = when the intermolecular forces btwn solute and solvent particles are GREATER (stronger/more stable) than the sum of the original intermolecular forces btwn particles of each of the pure liquids, then net energy is released and the process is deemed to be EXOTHERMIC
+ΔH = when the intermolecular forces btwn the solute and the solvent particles are LESS than the sum of the original intermolecular forces btwn each of the pure liquids, then the reaction is endothermic and HEAT is required to dissolve the solute in the solvent
What can you tell about the enthalpy based on whether something dissolves in a solution or not?
Nothing...unless you know the change in temperature of the solution as well...if the solution heats up, then you know that energy was released bc the new interactions formed between solute and solvent are more stable than the old interactions btwn solute and solute and solvent and solvent...thus the ΔH is negative...however, if the temp. decreases, then you know that energy must have been required from the system to allow for the thing to dissolve and thus the ΔH must be positive
Heat of solution (another form of enthalpy)
A certain mixture has a large negative heat of solution. Describe the relative strength of...
A) The intermolecular forces between solvent molecules
B) The intermolecular forces between solute molecules
C) The intermolecular forces between solute and solvent molecules
A) the forces are weaker than the forces btwn solute and solvent molecules
B) the forces are weaker than the forces btwn solute and solvent molecules
C) the forces are stronger than the forces of both the solute-solute molecules and solvent-solvent molecules put together
What happens to entropy when a solution forms?
Entropy always increases
Vapor Pressure
when the gas phase and liquid phase are in equilibrium with one another, the vapor pressure is equal to the partial pressure of the gaseous form above the liquid
How is vapor pressure affected by the following? And why?
A) Temperature
B) Addition of a non-volatile solute
C) Addition of a volatile solute
A) As temp. increases, vapor pressure increases...the average kinetic energy of the molecules in the liquid increases and this increase in energy allows the molecules to exert a greater pressure on the ambient pressure above them...this is why things will boil at higher temperatures...things only boil when the vapor pressure exceeds the ambient pressure and you can reach that by increasing the temp.
B) Decreases vapor pressure...bc there are more molecules near the surface exerting a smaller push against the ambient pressure
C) Increases vapor pressure...bc there are more molecules on the surface that are creating a stronger "push" against the ambient pressure...the increased vapor pressure of the volatile solute increases the vapor pressure of the solvent due to its stronger "push"
What does solubility mean?
it refers to how many particles of the original compound are able to dissociate and dissolve in a solvent
the part that does not dissolve in a particular solvent...this may be because the substance is insoluble or because there is more solute available in the solvent than can be dissolved in the solvent and thus the extras precipitate out
Saturated solution
a solution that has the solute completely dissolved to the max...if any more solute is added to the solution it will precipitate out
Unsaturated solution
a solution that can still accomodate more solute...if more solute is added, it will still dissolve
Super-saturated solution
-solution that can accomodate MORE solute than it normally can according to the solute's NORMAL solubility...this may be a result of an increase in temperature
-solutes that don't dissolve all the way are essentially like endothermic reactions...the solute-solvent interactions are not as strong as the solvent-solvent and solute-solute interactions...thus, bc it is an endothermic reaction, by adding more HEAT it will drive the reaction to the right and create more product (dissolve more of the solute)
"Like dissolves like"
polar substances are soluble in polar solvents and non-polar substances are soluble in non-polar solvents
Why are you able to leave out the pure solids and pure liquids from an equilibrium equation?
-you can only leave them out because the CONCENTRATION does not change...for example, pure water can be left out of an equilibrium equation bc in a particular reaction, the amount of water used up compared to the amount of water there is available is so small that there is practically no change in the amnt. of water, and thus, there is no change in the [H2O]...since it is ONLY the surface of a solid that is able to react with another species, even as some of the solid is used up in the reaction and there is less TOTAL solid available, there is still the SAME amount of solid that is ACTUALLY available to react (bc the surface area does not change...at least not much)
What is Ksp?
Ksp = the equilibrium constant for how much of a particular solute dissolves in a solvent...it is the exact same as Keq, Ka, Kb, etc.
How to calculate Ksp
1) Leave out pure liquids and pure solids (thus, ALL your Ksp equations should be 1 line only!!! if you have a denominator, it's WRONG!)
2) Temperature is the ONLY thing that changes Ksp..or pressure, when dealing with solubility of a gas
3) Ksp can only be measured for a saturated solution...this is because saturation is the point at which the dissolution reaction has reached equilibrium
When can equilibrium constants be measured?
the ONLY times when equilibrium constants (Ksp, Keq, Ka, Kb, etc.) can be measured is when the reaction/solution is at equilibrium...at ANY other point in time, the measurement will be inaccurate
What happens to the Ksp in a supersaturated solution? How?
-Ksp increases...it increases because more is able to dissolve due to an increase in temperature
What is the difference between solubility and the solubility constant?
The solubility constant (Ksp) is a measure of how much of a substance dissolves into its various individual pieces in a solvent...the solubility is how many of the individual pieces can dissolve in the solvent...ex- CaCl2 -> Ca2+ + 2Cl- ...Ksp = [Ca2+][2Cl-]^2 ...Ksp = [s][2s]^2 ...s = solubility...Ksp = solubility constant
Ion product
a.k.a. solubility product...ion product is to Ksp what Q is to Keq...if the ion (solubility) product is greater than Ksp then the reaction will go to the left and precipitate will form...and if the ion (solubility) product is less than Ksp, then the reaction will go all the way to the right and NO precipitate will form...if the value is equal to Ksp, then the solution is exactly saturated and still no precipitate forms
How to calculate solubility
1) Write out the Ksp expression
2) Substitute the value for Ksp into the expression
3) Substitute x into the equation for each ion (use 2x, 3x, etc. for each mole of an ion produced
4) solve for x...your answer, "x", is the solubility of that particular species
Describe the Common Ion Effect and define the term "spectator ion"
What is the solubility for the following compound dissolved in water? CaCl2
Ksp = [Ca][Cl]^2 ... Ksp = [s][2s]^2 ... Ksp = 4s^3...what this question is asking is how much of the parent molecule was dissolved? thus, the parent molecule dissolved into 4s^3 solute particles, but there was still only "s" parent particles that actually dissolved...thus, you solve for "s"
What will happen if a common ion is added to a saturated solution? What will happen if a spectator ion is added to the same solution?
-if a common ion is added to a saturated solution, a precipitate will form...this is bc the common ion increases the # of dissolved ions in the solution and makes the ion product GREATER than the Ksp...thus, the rxn goes to the left and a precipitate forms
-however, if a spectator ion is added there is no change...as many spectator ions can be dissolved in the solvent as is possible for their own Ksp value without forming a precipitate
What are some ways that you can drive a solubility reaction forward so that you can dissolve more of the original solid and form more of the dissolved products?
1) Increase temperature
2) Take away products (according to Le Chatelier's principle, more products will then be formed)...but how can you remove products that are soluble in a solution? you can remove them from the solution by several methods...Add acid to protonate one of the products (some species will become neutral and will thus precipitate out of an aqueous solution and others, like (CO3)2- will be unstable and will break apart into a gas (H2O and CO2) and thus leave the solution...or find a compound with 2 spectator ions that has a higher Ksp (why?) than the Ksp of a product that would be formed with one of your products from the original dissolution
3) Increase the amount of solvent...increasing the amount of solvent will decrease the CONCENTRATIONS of both the dissolved ions and thus allow more of the solid to dissolve to reach its Ksp again
You want to dissolve more NaCl in an aqueous solvent (Ksp = 4.510^-5) so that you can harvest a solution with a lot more Na+ ions than you would get from simply dissolving NaCl in water. You can't use heat because you can't effectively store your solution with its high Na+ content since you can't have it on the heat all the time. You choose to add another compound (LiNO3) because you know that the solubility constant for LiCl is 310^-5.
1) By adding this compound, will you actually get a solution more saturated with Na+ ions than you normally would?
2) How would you be able to answer that question without knowing the Ksp values?
Yes...you know that LiNO3 will be able to dissolve just fine in the solution without affecting its own ability to dissolve or the ability of NaCl to dissolve bc there are no common ions...however, LiCl is a precipitate that CAN form and it has a smaller Ksp value than NaCl and since the NaCl solution is saturated, you know that the ion product for LiCl is GREATER than the Ksp for LiCl (due to the higher number of Chloride ions than there should be)...thus, the reaction will drive to the left and form LiCl precipitate...this will then cause a chain reaction bc by decreasing the amnt. of Cl- ions present, this will increase the amnt. of NaCl dissolved and that will again increase [Cl-] which will continue formation of more LiCl precipitate...the reason you can figure this out without the Ksp values is by comparing what you know about the ionic character of NaCl compared to LiCl and the condosity of LiCl...Na is bigger than Li and has greater metallic character so NaCl will have more ionic character than LiCl...plus, we know that NaCl dissolves better in H2O than LiCl does based on what we know about the condosity of a known concentration of LiCl
Which components of a compound will immediately tip you off so that you can know that it is either soluble or insoluble on the MCAT?
Soluble: contains nitrate, ammonium, and ALL alkali metals (Group IA)
Insoluble: contains carbonate, phosphate, silver (Ag), mercury (Hg) and lead (Pb)...UNLESS one of these is paired with something from the soluble list; then it IS still soluble
5 General Characteristics of Gases
1) Gases are far less dense than liquids or solids
2) They contain much weaker intermolecular attractions than do liquids or solids
3) Gas molecules are very far apart
4) Gas molecules move very fast and collide billions of times per second
5) Gases are ALWAYS miscible (with other gases) regardless of their polarity...think about how many different gases exist in the atmosphere
Why don't gas molecules eventually lose their energy as a result of constantly colliding with one another?
They must be perfectly elastic collisions in order for the energy to be conserved
Effect of temperature on gas solubility
Gas solubility is the exact opposite of solubility for liquids and solids...by increasing the temperature, the solubility of the gas DECREASES bc the vapor pressure will increase, and thus, more of the gas particles will leave the liquid and go into the gas phase...by decreasing the temperature, the solubility of the gas INCREASES bc the kinetic energy of the gas particles is lower and thus the liquid particles have greater attractive forces on it that hold it in the liquid phase...also, the vapor pressure is lower, so fewer molecules are able to leave the liquid phase and thus the solubility is higher
Effect of pressure on gas solubility
By increasing the vapor pressure of gas X over a liquid, the solubility of gas X in that liquid INCREASES...(this is why they pressurize soda pop cans with excess CO2)
The Ideal Gas Law
PV = nRT ...R= ideal gas constant = 0.0821 Latm/molK OR 8.31 J/mol*K
What assumptions are made concerning Ideal Gases?
1) Gas molecules have no volume
2) No intermolecular forces exist between gas molecules
3) All collisions are perfectly elastic
4) Average KE is exactly proportional to temperature
*Important! the first 2 assumptions the most responsible for the differences btwn what PV = nRT predicts and how real gases actually behave
Standard Temperature and Pressure...a set of standard conditions true of any ideal gas said to be "at STP"...for the MCAT assume that all gases are ideal and start out at STP
What are the values for each of the variables at STP?
P = 1 atm
V = 22.4 Liters
n = 1 mole
R = 0.0821 Latm/molK or 8.31 J/mol*K
T = 273 K (0*C)
What is the difference between STP and Standard conditions?
Standard conditions usually indicates 25*C plus several other agreed-upon conditions at which thermodynamic data are always measured
What is the new volume of a gas that is increased to have 3 moles of gas and the new pressure on it is 2 atm?
Manipulate the equation! There is twice the pressure as normal which means that volume will be cut into half, but there are also 3 times as many moles which means that the volume is 3 times as great...thus, the volume ends up being 3/2 times as great
Another way to solve ideal gas problems
P1V1/T1 = P2V2/T2
In what situations would ideal gas behavior most differ from real gas behavior?
1) When the temperature is extremely low
2) the pressure is extremely high
-both of these situations cause the gas molecules to be very close together which allows for greater intermolecular interactions and several other things occur that cause a great deviation from ideal behavior
How can you determine what caused a deviation from ideal gas behavior? (volume or intermolecular forces?)
1) If PV/RT > 1, then the cause is most likely due to the molecular volume assumption (gas molecules actually do have volume)
2) If PV/RT < 1, then the cause is most likely due to the intermolecular forces assumption (the gas does experience intermolecular interactions)
Formula for Dalton's Law of Partial Pressures
Ptotal = P1 + P2 + P3...
How will the following be affected if you add 1 mole of gas A to a mixture of gases A, B, and C?
1) Total pressure
2) Partial pressure of gas A
3) Partial pressure of gas B and gas C
4) # of moles of gas B and gas C
5) Mole fraction of gas B
6) Solubility of gas A in the liquid it was hovering above
7) Solubility of gas B and gas C in the liquid it was hovering above
1) total pressure increases
2) Partial pressure of gas A increases
3) Partial pressure of gas B and gas C remain UNCHANGED
4) # of moles of B and C obviously doesn't change
5) Mole fraction of gas B decreases
6) Solubility increases bc total pressure increased
7) Solubility of gases B and C do NOT change bc their partial pressures are still the same
What is diffusion?
-the process by which gas molecules spread from areas of high concentration to areas of low concentration due to the random motion imparted to them as a result of their kinetic energy and collisions with other molecules
What is effusion?
the diffusion of gas particles through a pin hole
What is the definition of a pin hole?
a hole smaller than the average distance a gas molecule travels between collisions
Graham's Law
Rate1/Rate2 = √(MW2)/√(MW1)...the rate can be the rate of either effusion or diffusion...notice that the rate of effusion or diffusion is INVERSELY proportional to the molecular weight of the gas
You have 2 different containers. In container A you have 8 moles of H2(g) and 2 moles of Cl2(g). In container B you have 2 moles of H2(g) and 8 moles of Cl2(g). Which container has higher pressure?
Neither...they both have the exact same pressure...according to the ideal gas law ALL gases behave the exact same way regardless of what gas it is
How can you change the boiling point of a particular species? What are some ways that might trick you into thinking it increases boiling point but really don't?
1) Increase/decrease ambient pressure...this makes it harder/easier for the species (that has the SAME vapor pressure) to break free of the liquid state
2) Add non-volatile/volatile solute (non-volatile solute crowd the surface and don't allow for as great of a net vapor pressure on the atmosphere....volatile solutes crowd the surface as well, but they help PUSH against the atmosphere along with the solvent and thus increase the vapor pressure bc of a net increase in the PUSH force)
1) NOT by changing intermolecular forces, MW, or branching...you can NOT change anything about the species itself
2) NOT by changing temperature...an increase in temperature increases the vapor pressure, but the boiling point remains the same...the solution simply starts to boil when the vapor pressure is equal to the ambient pressure
If you change the boiling point of a substance, do you change its vapor pressure as well?
Not necessarily...when you add volatile/non-volatile substances you DO change vapor pressure...but when you simply change bp by increasing/decreasing ambient pressure, nothing happens to the vapor pressure
Also, when the boiling point changes, it only changes bc you are changing circumstances OTHER than the vapor pressure that allow the vapor pressure to either be closer to overcoming the ambient pressure or farther from it
What is the difference between evaporation and boiling?
1) Evaporation is a constantly occurring phenomenon...it is when particles of the solvent simply leave the liquid phase and enter the gas phase...this occurs at ALL temperatures but MORE particles do this at higher temperatures
2) Boiling is when you see bubbles...this happens when the vapor pressure increases so much that it is equal to the ambient pressure...at this point, not only is it leaving the surface via normal evaporation, but there is so much vapor pressure, that it enters the gas phase in the middle of the liquid (which is why the gas bubble forms!!)
If the freezing point depression constant (Kfp = 1.86*C/m for water) what is the the freezing point of a mixture of 1.6 moles of glucose in 1000 g of H2O?
A) .003 *C
B) -.003*C
C) 3*C
D) -3*C
ΔT = Kfp(molality)(i) .... i for glucose = 1 ... molality = moles of solute/mass of solvent in kg = 1.6 molal ... 1.6(1.86)(1) = 3 ...but since the freezing point is being depressed, the ΔT = -3*C...thus D is correct
Is ΔH of fusion greater for a container of ice containing 10 gallons of ice compared to another container with only 1 gallon of ice? If you were to compare the energy diagram for the 2 different containers, which would have a longer horizontal line for the phase change from solid to liquid?
1) No...ΔH of fusion = J/kg...thus, it doesn't matter how much of it you have...ΔH of fusion only measures how much energy is required to break apart the intermolecular interactions btwn molecules in a solid (or how much energy is released when the interactions form)
2) However, more heat IS necessary to convert the solid ice into liquid ice simply bc there are a lot more intermolecular interactions that must be broken apart...thus, the 10 gallon container would have a much longer line on the energy diagram to illustrate how much more heat is required to be put into the system to melt all the ice...this is bc the x-axis on the phase change energy diagrams is heat which is simply joules...thus, it makes sense that more energy is required to break apart more intermolecular interactions
If you have 500 mL of 1 molar H2SO4, what is the normality of the solution?
H2SO4 produces 2 mol of equivalents (H ions)...thus, normality = 2N
N = nM ...N = normality ...n = number of equivalents the solute produces in solution ...M = molarity of the solute
You have 100 grams of compound A (MW = 50g/mol) in a 200 mL solution. What is the normality of the compound. Assume that it is a diprotic acid.
2 moles of compound A in 200 mL of solution...thus, the molarity is 2/.2 = 10 M...we know that it is a diprotic acid, and therefore has 2 equivalents for every mole, thus we know that the normality is N = nM = 20 N
You have a 2 N solution of H3PO4. What is the molarity?
N = nM.... 2/3 = M = .66 M
You dissolve a solute in a particular solution and the solution and the beaker warm up. What does this tell you about...
A) ΔH of solution
B) Solute-solute intermolecular forces
C) Solvent-solvent intermolecular forces
D) Solute-solvent intermolecular forces
E) stability of the new solution compared to the stability of either of the 2 pure mixtures
A) ΔH is negative bc heat is released from the mixing into the environment
B, C, and D) The sum of the solute-solute and solvent-solvent intermolecular forces is WEAKER than the solute-solvent intermolecular forces
E) the new solution is more stable than either of the 2 pure mixtures
A particular solvent experiences highly attractive intermolecular forces with itself and is then put into a solvent. What will be required to make the solute dissolve?
Heat must be added...bc the solute-solute intermolecular forces are stronger than the solute-solvent intermolecular forces, this reaction is ENDOTHERMIC...heat must be added to get the solute to dissolve
1) Is a mixture of polar solute in water an exothermic or endothermic reaction?
2) Is a mixture of oil in water an exothermic or endothermic reaction?
1) Exothermic...the new intermolecular forces btwn solute and solvent are more stable...which is why it dissolves! a little bit of energy is released...could potentially warm up the beaker if -ΔH is large enough
2) Endothermic...the new intermolecular forces are very unstable...thus, a lot of heat is required to make the non-polar mix dissolve... thus, ΔH is positive
What is the solubility of CaCl2 in terms of S?
Ksp = [Ca2+][2Cl-]^2 ...replace the dissolved ions with S ... Ksp = [S][2S]^2 = 4S^3 ....if you were to use Ksp to solve for S, S would be equal to the number of particles of CaCl2 that would dissolve in solution
What is the vapor pressure of a boiling pot of water that is 1 meter tall? Calculate vapor pressure for the bubbles forming at the bottom of the pot.
Vapor pressure must exceed ambient pressure in order for the substance to leave the liquid state and enter the gas state...thus, we know that vapor pressure is at least equal to the pressure at the bottom of the pot....P at bottom of pot = 101.325 kPa + 1000kg/m^3(10m/s^2)(1 m) = 111.325 kPa
What is the name for Ksp?
solubility product...equilibrium constant for solubility of a compound...if the ion product is greater than the solubility product, a precipitate will form
You are doing an experiment on 2 moles of gas molecules. However, your results are not matching the results you had expected when you calculated your estimates using PV = nRT. What is most likely causing the discrepancy?
PV/RT > 1...bc you have 2 moles of gas molecules...this means that your gas sample is not following ideal gas behavior which is why your results are off...specifically what is wrong, is that the gas molecules actually DO have volume and that assumption is not able to hold true in this situation bc PV/RT > 1
You are doing an experiment on a sample of a particular gas. You calculate the air pressure to be 2 atm in a 4 L space. The temperature of the gas is 25*C. However, you notice discrepancies in your expected data with the real data. What is most likely the problem?
PV/RT < 1... 2(4)/(8.34)(298) < 1...this tells us that the gas is not following ideal gas behavior and the reason why is bc the assumption that the gas molecules are NOT experiencing any intermolecular interactions is the assumption that is violated bc PV/RT < 1
Osmoles/L ... Osmolarity = Molarity*i ... i = number of things the solute dissolves into...NaCl = i = 2...glucose = i = 1
1) What happens to the temperature of water when you add salt?
2) What happens to ice when you add salt to it?
3) What happens to the melting point of water when you add salt?
Explain the opposing forces of vapor pressure and partial pressure
-When vapor pressure = partial pressure, the system is at EQUILIBRIUM
-when vapor pressure > partial pressure, the liquid boils
-when vapor pressure < partial pressure, the liquid stays in liquid form
Rank the Ksp of the following molecules in order of smallest to largest:
1) Ammonia
2) Propane
3) 2-bromopropane
Ammonia > 2-bromopropane > propane
1) How does pressure affect Ksp?
2) How does temperature affect Ksp?
1) Pressure ONLY affects the Ksp of GASES...if pressure increases, the Ksp of the gas will increase
2) Temperature affects Ksp of solutes that are GASES, LIQUIDS, OR SOLIDS....if the solute is a liquid or SOLID, increasing temperature will INCREASE Ksp....if the solute is a gas, increasing temperature will DECREASE Ksp
When do you use .0821 for R and when do you use 8.314 for R? What are the units of each?
1) .0821 J/mol*K ... J = Fd = PV = (F/A)(V)...you use this when you pressure is in Pascals (bc pascals are N/m^2) and when your volume is in m^3
2) 8.314 atmL/molK... obviously you use this when your pressure is in atmospheres and your volume is in liters
If you add 50 grams of KNO3 (MM = 101 g) to 1 liter of solution, what is the molarity of the solution? The density of KNO3 is 2 g/mL.
calculate moles of KNO3 = .5 moles...you calculate the number of moles of the compound ADDED, without considering what dissolves in the solution....moles/liters of solution...you can ignore the change in total solution volume caused by adding the KNO3...??? really
You have a 2.3 M solution of Mg(NO3)2. If you were to have 2 liters of this solution, how many NO3 particles would you have in the solution?
The MOLARITY of the solution is what remains constant, NOT the number of moles of solute...2.3 mol/L, thus with 2 liters you have 4.6 moles of solute...4.6 moles of solute dissolves to give you 4.6 moles of Mg and 9.2 moles of NO3...(9.2 * 6e^23) = 5.4e^24 NO3 ions in solution
Calculating change in vapor pressure after adding a volatile solute
Raoult's Law = (mole fraction of the solvent)(VP of the solvent) + (mole fraction of solute)(VP of solute)
Calculating change in vapor pressure after adding a non-volatile solute
Raoult's Law = (mole fraction of the pure solvent)*(VP of the pure solvent)...mole fraction decreases when you add more solute
Henry's Law
Vapor pressure of solute = (mole fraction of solute)*(Henry's Law Constant)
What is the relationship between heat of solution and vapor pressure?