CH 12 - Solutions Flashcards
1
Q
What is a thirsty solution?
A
a solution that draws more water to itself.
Sea water draws water from the body leading to dehydration.
2
Q
Solution
A
a homogeneous mixture of two or more substances or components.
3
Q
Solvent
A
the majority component in a solution
4
Q
Solute
A
the minority component in a solution
5
Q
In most solutions the particles of the solute interact with the particles of the solvent through ________.
A
IMFs
6
Q
Substances tend to __________.
A
combine into uniform mixtures unless it is highly unfavorable energetically.
7
Q
Aqueous Solution
A
water is the solvent and a solid, liquid, or gas is the solute.
8
Q
Aqueous Solution examples
A
- Salt(solid) in seawater
- ethyl alcohol(liquid) in alcohol
- CO2(gas) in club soda
solid, liquid, and gas in water
9
Q
3 types of common solutions
A
- Gaseous Solutions(gas with gas)
* Air - Liquid Solutions(liquid solvent with solid, liquid, or gas solute)
- seawater, vodka, club soda
- Solid Solutions(solid with solid)
- brass(copper and zinc) and other alloys
10
Q
Solubility
A
the amount of the substance that will dissolve in a given amount of solvent.
11
Q
What does solubility govern?
A
the ability of a substance to dissolve in another substance.
12
Q
Is grease soluble in water?
A
Solubility is nearly 0 so no grease will not dissolve in water.
13
Q
The formation of a solution ____________ the potential energy of its constituent particles.
A
does not necessarily lower
- most other interactions do(cations and anions want lower PE!)
14
Q
The tendency for ideal gases to mix does not _______ PE but is related to ______,
A
lower
entropy
15
Q
Entropy
A
a measure of the energy randomization or energy dispersal in a system.
16
Q
Entropy:
What happens when Neon and Argon mix from two separated containers into one?
A
even though the two gases have low PE they mix and spread out, along with their kinetic energy, over the larger volume.
The mixture has greater energy dispersal, or greater Entropy, than the separated components.
17
Q
Why do two ideal gases mix?
A
The pervasive tendency for energy to spread out, or disperse, whenever it is not restrained from doing so is the reason.
Entropy NOT PE in effect here.
- You heat an iron rod on one side but the entire rod becomes evenly hot
18
Q
Does energy spontaneously concentrate?
A
No.
It will spontaneously spread out and disperse.
19
Q
What is the fundamental criterion that ultimately determines the spontaneity of any process?
A
The dispersal of energy(entropy)
20
Q
The formation or prevention of solutions is affected by _______.
A
IMF interaction
21
Q
Solvent-Solute interactions > solvent-solvent and solute-solute interactions
A
solution forms
22
Q
Solvent-Solute interactions = solvent-solvent and solute-solute interactions
A
Solution Forms
23
Q
Solvent-Solute interactions < solvent-solvent and solute-solute interactions
A
may or may not form depending on how large the relative disparity is.
Yes a solution can form even if an energetically uphill battle.
Too large and no solution.
24
Q
Solvent-solute interactions
A
the interactions between a solvent particle and a solute particle
25
Solvent-solvent interactions
the interactions between a solvent particle and another solvent particle
26
Solute-Solute interactions
the interactions between a solute particle and another solute particle
27
Miscible
the ability of two substances to mix without separating
28
Miscibility is governed by _______.
the IMFs between the solute-solvent, solute-solute, and solvent-solvent interactions
29
Miscibility:
solute-solvent, solute-solute, and solvent-solvent interactions are the same same type of IMFs(polar with polar etc) then ______.
they are the same magnitude and the substances are soluble in all proportions.
30
Miscibility:
solute-solvent, solute-solute, and solvent-solvent interactions are more strong with their own type then ______.
they may still be soluble even if energetically uphill assuming the difference isnt too large.
- H2O H bonds are too strong to be overcome by weak hexane IMFs and no solution would be formed.
31
Common Laboratory Solvents:
Polar
```
o H2O(water)
o CH3COCH3(Acetone)
o CH3OH(methanol)
o CH3CH2OH(ethanol)
```
32
Common Laboratory Solvents:
Nonpolar
```
o C6H14(Hexane)
o CH3CH2OCH2CH3(diethyl ether) has a small dipole moment and can be considered intermediate between polar and nonpolar
o C7H8(Toluene)
o CCl4(carbon tetrachloride)
```
33
3 steps in the formation of a solution
- Separating the solute into its constituent particles
- separating the solvent particles from each other to make room for the solute particles
- mixing the solute particles with the solvent particles
34
3 steps in the formation of a solution:
1. Separating the solute into its constituent particles
Endothermic(deltaH is +)
energy is required to overcome the forces that hold the solute particles together
35
3 steps in the formation of a solution:
2. separating the solvent particles from each other to make room for the solute particles
Endothermic(deltaH is +)
energy is required to overcome the IMFs among the solvent particles
36
3 steps in the formation of a solution:
3. mixing the solute particles with the solvent particles
Exothermic(deltaH is -)
Energy is released as the solute particles interact(through IMFs) with the solvent particles.
37
Hess law and *the enthalpy of solution* (deltaHsoln)
the sum of the changes in enthalpy from each step of the formation of a solution.
deltaHsoln = deltaHsolute(endothermic +) + deltaHsolvent(endothermic +) + deltaHmix(exothermic -)
The overall reaction can be endo or exothermic based on the magnitudes of the individual terms
38
Hess law and *the enthalpy of solution* (deltaHsoln):
3 outcomes to a solution formation
1. If the sum of the endothermic terms is approximately equal in magnitude to the exothermic term then deltaHsoln = 0
* solution formed
2. Endothermic < exothermic magnitude then deltaHsoln = -
* solution formed
3. Endothermic>exothermic magnitude then deltaHsoln = +
* solution may or may not form
39
Hess law and *the enthalpy of solution* (deltaHsoln):
3 outcomes to a solution formation:
o If the sum of the endothermic terms is approximately equal in magnitude to the exothermic term then deltaHsoln = 0
Entropy drives the mixing of a solution however overall energy remains nearly constand
40
Hess law and *the enthalpy of solution* (deltaHsoln):
3 outcomes to a solution formation:
o Endothermic < exothermic magnitude then deltaHsoln = -
Tendency towards lower energy and greater entropy drive formation of a solution
41
Hess law and *the enthalpy of solution* (deltaHsoln):
3 outcomes to a solution formation:
o Endothermic>exothermic magnitude then deltaHsoln = +
As long as deltaHsoln is not too large then the tendency to greater entropy drives the formation of a solution. deltaHsoln too large then no solution forms
42
Heat of Hydration(deltaHhydration)
in an aqueous solution, the deltaHsolvent and deltaHmix can be combined into a single term.
- the enthalpy change that occurs when 1 mol of the gaseous solute ions is dissolved in water
43
Heat of Hydration(deltaHhydration) is ________.
always largely negative(exothermic) for ionic compounds because the ion-dipole interactions between a dissolved ion and the surrounding water molecules are much stronger than the hbonds of water.
44
with Heat of Hydration
deltaHsoln =
deltaHsolute(+) + (deltaHsolvent + deltaHmix[combined term = deltaHhydration exothermic and -])
45
For ionic compounds the deltaHsolute is the __________.
negative of the solutes lattice energy
deltaHsolute = -deltaHlattice
46
\deltaHsolute\ < \DeltaHhydration\
o The amount of energy required to separate the solute into its constituent ions is less than the energy given off from the hydration of ions
o deltaHsoln = exothermic and negative
feels warm to the touch
47
\deltaHsolute\ > \DeltaHhydration\
o The amount of energy required to separate the solute into its constituent ions is greater than the energy given off from the hydration of ions
o If a solution forms at all then deltaHsoln = endothermic and +
feels cool to the touch
48
\deltaHsolute\ = \DeltaHhydration\
o The amount of energy required to separate the solute into its constituent ions roughly equal to the energy given off from the hydration of ions
o deltaHsoln = 0 or close to it
neither appreciably exo or endothermic
no noticeable change in temperature
49
the dissolution of a solute in a solvent is ________.
an equilibrium process similar to the equilibrium process associated with phase changes.
50
What happens to the rates of dissolution and recrystallization as equilibrium is reached?
initially the rate of dissolution exceeds the rate of recrystallization but as the concentration solute dissolves the rate of recrystallization increases eventually reaching equillibrium.
51
Saturated Solution
a solution in which the dissolved solute is in dynamic equilibrium with the solid(undissolved) solute.
Adding solute will not dissolve
52
Unsaturated Solution
A solution containing less than the equilibrium amount of solute.
Adding solute will drissolve
53
Supersaturated Solution
A solution containing more than the equilibrium amount of solute.
Unstable and the excess normally precipitates out.
in some cases, if left undisturbed, a supersaturated solution can exist for an extended period of time.
54
How does the temperature affect the solubility of solids?
The solubility of most solids in water increases with increasing temperature.
55
Will more or less sugar dissolve as water temperature increases?
More sugar will dissolve as temperature increases.
56
Recrystallization
a common technique to purify a solid.
Enough solid is added to water(or another solvent) to create a saturated solution at an elevated temperature, then as the solution cools and become supersaturated the excess solid precipitates out of the solution.
If the solution cools slowly then a solid is formed.
57
During recrystallization the crystalline structure tends to _______.
reject impurities leaving a pure solid substance.
Rock candy is recrystallized sugar.
58
2 major factors affecting the solubility of gases in water.
Temperature and pressure.
59
How does temperature affect the solubility of gases in liquids?
The solubility of gases in liquids decreases with increasing temperature.
DIFFERENT for solids who tend to become more soluble with increasing temperature.
60
What is a major difference of temperature on solids and gases solubility?
Solids become more soluble with increasing temperature.
Gases become less soluble with increasing temperature
61
How does pressure affect the solubility of gases in liquids?
the higher the pressure of a gas above a liquid the more soluble the gas is in the liquid.
CO2 in a closed soda can is soluble but open the can and CO2 fizzes out as gas as the pressure decreases.
62
Henrys Law
Sgas = (Kh)(Pgas)
Sgas = solubility in M
Kh = constant of proportionality(henrys law constant) depending on the specific solute, solvent, and temperature
Pgas = partial pressure of the gas(ATM)
63
Henrys law shows that _______.
the solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid.
64
Constants(Kh in M/atm) of Several Gasses in Water at 25 degrees C:
O2
1.3 * 10^-3 M/atm
65
Constants(Kh in M/atm) of Several Gasses in Water at 25 degrees C:
N2
6.1 * 10^-4 M/atm
66
Constants(Kh in M/atm) of Several Gasses in Water at 25 degrees C:
CO2
3.4 * 10^-2 M/atm
67
Constants(Kh in M/atm) of Several Gasses in Water at 25 degrees C:
NH3
5.8 * 10^1 M/atm
68
Constants(Kh in M/atm) of Several Gasses in Water at 25 degrees C:
He
3.7 * 10^-4 M/atm
69
Dilute Solution
Contains small quantities of solute relative to the amount of solvent.
NaCl in water will not dehydrate you here
70
Concentrated Solution
Contains large quantities of solute relative to the amount of solvent.
NaCl in water will dehydrate you here.
71
Molarity(M)
the amount of solute(in moles) divided by the volume of solution(in L)
Moles of solute per liter of SOLUTION not per liter of solvent.
72
How do you make a solution of specified molarity?
Put the solute into a flask and then add water(or other solvent) to the desired volume of solution.
73
What is Molarity a convenient unit for?
making, diluting, and transferring solutions because it specifies the amount of solute per unit of solution.
74
What does Molarity depend on?
Volume
Volume depends on temperature so molarity varies with temperature!
1M aq soln at room temp will be slightly less than 1M at an elevated temperature because the volume of the solution will be greater.
75
Molality(m)
a concentration unit independent of temperature.
The amount of solute(moles) divided by the mass of the solvent(kg)
per SOLVENT not solution!
76
What is Molality(m) useful for?
To compare concentrations over a range of different temperatures.
does not fluctuate like Molarity since it is kg based not volume.
77
Parts by Mass:
the ratio of the mass of the solute to the mass of the solution multiplied by a multiplication factor.
(mass solute/mass solution)(mult factor)
78
Parts by Mass:
Percent by mass
(mass solute)/(mass solution)(100%)
14g of solute per 100g of solution
79
Parts by Mass:
Parts per million(PPM)
(mass solute)/(mass solution)(10^6)
15g of a solute per 10^6g solution
80
Parts by Mass:
Parts per billion(ppb)
(mass solute)/(mass solution)(10^9)
15g of a solute per 10^9g solution
81
Parts by Volume
typically used when both the solute and solvent are liquids.
usually a ratio of the volume of the solute to the volume of the solution multiplied by a multiplication factor
(volume solute/volume solution)(mult factor, %,ppm,ppb)
82
Using parts by mass(or parts by volume) in calculations
can use as a conversion factor between mass(or volume) of the solute and mass(or volume) of the solution.
a solution containing a 3.5% sodium chloride by mass:
- (3.5g NaCl)/(100g solution) converts g solution to g NaCl
- (100g solution)/(3.5g NaCl) converts g NaCl to g solution
83
Mole Fraction(Xsolute or Xsolvent
More useful in times where the ratio of solute to solvent can vary widely
- Xsolute = (amount solute in mol)/(total amount of solute and solvent in mol)
- (nsolute)/(nsolute + nsolvent)
- Xsolvent = (nsolvent)/(nsolute/nsolvent)
84
mole percent(mol %)
the mole fraction * 100%
- mol % = (Xsolute)(100%)
85
Colligative Property
a property that depends on the number of particles dissolved in solution NOT the type of particle.
86
Four Colligative Properties
Vapor Pressure Lowering
freezing point depression
boiling point elevation
osmotic pressure
87
What is the effect of a nonvolatile nonelectrolyte solute on the vapor pressure of the liquid into which is dissolves?
- the vapor pressure of the solution is lower than the vapor pressure of the pure solvent
- the nonvolatile solute particles interfere with the ability of the solvent particles to vaporize
* the rate of vaporization is thus diminished compared to that of the pure solvent
- subsequently the rates change allowing the rate of condensation to become greater than the rate of vaporization
- as molecules condense the vaporization pressure decreases until equilibrium is reached again
- natures tendency to mix(entropy)
- even if a pure solvent and a concentrated solution are combined in a seal container(even in separate beakers) the two will mix so the concentrated solution becomes less concentrated
- molecules evaporate from the pure solvent, then before dynamic equilibrium is reached the pressure increases to a point where the molecules condense in the solution
- this leads to the solvents molecules constantly evaporating as vapor pressure is never reached
88
Vapor Pressure Lowering:
Raoult's law
Psolution = XsolventPsolvent
- Psolution = vapor pressure of solution
- Xsolvent = mole fraction of the solvent
- Psolvent = vapor pressure of the pure solvent at the same temperature
89
25degree C water contains 0.90 mol of water and 0.10mol of a nonvolatile solute such as sucrose. The pure water has a vapor pressure of 23.8torr. Fine the vapor pressure of the solution:
Psolution = (XH2O)(PH2O)
= (0.90)(23.8toss)
Psolution = 21.4 torr
- the vapor pressure of the solution is directly proportional to the amount of solvent in the solution
90
vapor pressure lowering(deltaP) – deltaP = Psolvent – Psolution
the difference in vapor pressure between the pure solvent and the solution
- for a two component solution Xsolvent = 1 – Xsolute in raoults law
P solution = (Xsolvent)(Psolvent)
P solution = (1 - Xsolute)(Psolvent)
Psolvent – Psolution = (Psolvent)(Xsolute)
deltaP = (Xsolute)(Psolvent)
- indicates the lowering of vapor pressure is directly proportional to the mole fraction of the solute
91
Vapor Pressures of Solutes Containing a Volatile(nonelectrolyte) Solute
when there is a volatile solvent AND volatile solute BOTH solvent and solute contribute to the overall vapor pressure of the solution
92
idea solution
one which behavior follows Raoults law at all concentrations for both solvent and solute
- will follow Raoults Law exactly
93
nonideal solution
behavior of solvent and solute does not follow Raoults law
94
an ideal solution is similar to an ideal gas
- will follow Raoults Law exactly
- solute-solute, solvent-solvent, and solute-solvent interactions are similar in magnitude
- the solute simply dilutes the solvent and idea behavior is observed
- for a two-component solution containing liquids A and B:
- Pa = (Xa)(Pa)
- Pb = (Xb)(Pb)
- Ptot = Pa + Pb
95
in a nonideal solution the solute-solvent interactions are either stronger or weaker than the solvent-solvent interactions
- if solute-solvent interactions are stronger then the solute tends to prevent the solvent from vaporizing as readily as it would otherwise
- if the solution is sufficiently dilute then the effect will be small and Raoults law works as an approximation
- however, if the solution is not dilute the effect will be significant and the vapor pressure of the solution will be LESS than that predicted by Raoults law
- if solute-solvent interactions are weaker than solvent-solvent interactions then the solute tends to allow more vaporization than would occur with just the solvent
- if the solution is not dilute the effect will be significant and the vapor pressure of the solution will be GREATER than predicted by Raoults law.
96
vapor pressure lowering occurs at all temperatures
the net effect is that the solution has a lower melting point and a higher boiling point than the pure solvent
97
freezing point depression
solutions have lower freezing points than pure solvents
98
boiling point elevation
solutions have higher boiling points than pure solvents
99
the more concentrated the solution the _______.
more drastic the change to freezing point and boiling point
100
Freezing point depression:
deltaTf = (m)(Kf)
- deltaTf = change in temperature of the freezing point in Celsius(relative to the freezing point of the pure solvent)(usually reported as a + number)
- for water Kf = 1.86 C/m
- when an aqueous solution containing a dissolved solid solute freezes slowly, the ice that forms does not normally contain much of the solute
101
Boiling Point Elevation:
deltaTb = (m)(Kb)
- deltaTb = change in temperature of the boiling point in Celsius(relative to the boiling point of the pure solvent)
- for water Kb = 0.512 C/m
102
Osmosis
the flow of solvent from a solution of lower solute concentration to one of higher solute concentration
103
concentrated solutions draw solvent from more dilute solutions because of _______.
natures tendency to mix
104
Semipermeable Membrane
a membrane that selectively allows some substances to pass through but not others
- separates two halves of the cell
105
Osmotic Pressure
the pressure required to stop the osmotic flow
106
osmotic pressure =
MRT
- M = molarity of the solution
- T is the temperature in Kelvin
- R is the idea gas constant((0.08206L)(atm)/(mol)(K))
107
Van't Hoff Factor(i)
i = (moles of particles in solution)/(moles of formula units dissolved)
- the Van’t Hoff Factor approaches the expected value at infinite dilution(as approaching 0)
- the reason it is not the exact predicted amount(like 2 for 2mols) is because some of the ions will pair together(cation and anion) so the dissociation of the solute in the solution will never truly be 100% complete.
- this slightly reduces the number of particles in the solution
108
using the Van't Hoff factor with ionic solutions
to calculate freezing point depression, boiling point elevation, and osmotic pressure of ionic solutions plug Van’t Hoffs factor in to the equation:
- deltaTf = (im)(Kf) <- osmotic pressure
109
Strong Electrolytes and Vapor Pressure with electrolytle solutions
the vapor pressure lowering is greater than in nonelectrolyte solutions
110
Colligative Properties and Medical Solutions
Hyperosmotic, hyposmotic, and intravenous solutions
111
Colligative Properties and Medical Solutions:
Hyperosmotic
a solution with an osmotic pressure greater than those of body fluids
- takes water out of the cells and tissues
- cells tend to shrivel as this happens
112
Colligative Properties and Medical Solutions:
Hyposmotic
solution with an osmotic pressure less than those of body fluids
- pushes water into the cell, sometimes causing it to burst
113
Colligative Properties and Medical Solutions:
Intravenous solutions
those administered directly into a patients veins
114
Colligative Properties and Medical Solutions:
Intravenous solutions:
isosmotic(isotonic)
osmotic pressure = to those of body fluids
- usually an isosmotic saline solution
- percent mass to volume solution
- 0.9g NaCl per 100mL solution
- 0.9% mass/volume
115
colloidal dispersion
a mixture in which a dispersed substance(which is solute like) is finely divided in a dispersing medium(which is solvent like)
- a colloid
- fog, smoke, whipped cream, milk, opal
116
How is a colloid determined?
- if small(like individual molecules) then it is a solution
- if large, greater than 1um diameter then it can be colloidal which is a heterogeneous mixture
- dust
- the particles must be between 1nm and 1000nm to be a colloid
117
Size range of a colloid
1nm and 1000nm
118
Brownian motion
jittery particle motion observed under a microscope due to the particles colliding with molecules in the liquid
- was a decisive factor in confirming the molecular and atomic nature of matter
119
micelles
nonpolar hydrocarbon tails crowd into the center of a sphere to maximize their interactions with one another while ionic heads orient toward the surface of the sphere where they can interact with polar water molecules
- soap does this and creates balls inside water
- responsible for the hazy seen in soapy water because even though they are too small to be seen by the naked eye they still scatter light.
120
Tyndall Effect
the scattering of light by a colloidal dispersion
- can be used to test if a mixture is colloidal
- can see light beams then colloidal otherwise no
121
How are colloidal suspensions of micelles kept stable?
electrostatic repulsions that occur at their surfaces
122
How can you disrupt colloidal suspensions?
- adding heat(increasing collisions) or an electrolyte can disrupt the electrostatic repulsions and destroy the colloid
- soap does not work well in salt water but great in fresh water
123
Colloids can be:
- clusters of molecules or macromolecules
- protein solutions
- blood contains hemoglobin which is large enough to scatter light
- blood is colloid
