Mixtures

A nonuniform mixture that has regions of different composition.
Ex. Sand and water

A uniform mixture that has the same composition throughout.
Ex. Salt and water

A homogeneous mixture of particles of molecular size; particles can be separated by physical means.

** Homogeneous
** Particles will not come out of solution
** Single phase (salt was a solid but is now dissolved
** Can pass through a filter w/o separating
** Clear – particles can not be seen.

The part of a solution in which a substance is dissolved.

The dissolved substance in a solution.

A solution where water is the solvent.

Alcohol is the solvent.

- Two driving forces behind formation of solution
1. Disordering/Randomizing*
2. Intermolecular Forces
** Whether or not a solution forms depends on both opposing forces
In the thermodynamics chapter this is called entropy and is defined more precisely

- Two gases mix spontaneously
** Due to random motions
** Mix without outside work
** Never separate spontaneously
- Tendency of system left to itself, to 
become increasingly disordered

- When mixed, for a solution to form,
** Solvent-to-solute attractions must be similar to attractions between solute alone and solvent alone
- Initially solute and solvent separate to make room for each other
** Solute molecules held together by intermolecular forces: energy added
** Solvent molecules held together by intermolecular forces: energy added

Solute molecules are attracted to solvent molecules: energy released
Solvent molecules are attracted to solute molecules: energy released
** If there is a net release of energy, the solution will form.
(shows that formation of solutions depend on INT Forces)

- Two liquids that dissolve in one another in all proportions
- Form solution
- Strengths of intermolecular attractions are similar in solute and solvent
- Similar polarity
e.g., Ethanol and water
** mixing alcohol (is polar) and water (is polar) = completely mix **like dissolves like; oil (nonpolar) and water = immiscible = London Dispersion Forces

can participate in hydrogen bonding = could act as a double agent — dissolving polar and non polar; different ends attracted to different polarity

first we break apart solute particles, then solvent particles: releases more energy then absorbed — energy is going out to the surrounding is an exothermic process (not the only possibility) ** however, solubility always requires energy. comparing the start of energy to the end of energy

- Two insoluble liquids
- Do not mix
- Get two separate phases (regions)
- Strengths of intermolecular forces are different in solute and solvent
- Different polarity
e.g., Oil and water
- - you could get liquids of different densities to mix, if something does not mix it is because they do not have the same intermolecular forces; However, density WOULD explain why something is at the top and another at the bottom

- Like dissolves like
** Use polar solvent for polar solute
** Use nonpolar solvent for nonpolar solute
- When strengths of intermolecular attractions are similar in solute and solvent, solutions form because net energy exchange is about the same
** soluble = similar intermolecular forces

- Basic principles remain the same when solutes are solids
- Sodium chloride (NaCl)
** Ionic bonding ( is stronger then all the INT F)
** Strong intermolecular forces
** Ions dissolve in water because ion-dipole forces of water with ions strong enough to overcome ion-ion attractions

ex: Na+Cl- (all sodium salts dissolve in water); However, with a salt (carbonates = however, there are exceptions) will not dissolve bc the intermolecular force is not strong enough to break those bonds — within the textbook there is a chart of solubility — what dissolves and what does not

- Ions at edges, fewer ion-ion attractions
** Multiple ion-dipole attractions formed with water
** These overcome stronger ion-ion electrostatic attractions
- New ion at surface
** Process continues until all ions in solution
- Hydration & Solvation

Ions surrounded by water molecules

General term for surrounding solute particle by solvent molecules

** H2O reorients so
- Positive H atoms are near negative ends of solute
- Negative O atoms are near positive ends of solute
** slightly positive hydrogen atoms — this is dipole dipole forces (polor - polar)

- Wax is a mixture of long chain alkanes and is nonpolar
- Benzene is a nonpolar solvent
- Benzene can dissolve waxes
** Weak London dispersion forces in both solute and solvent. Weak London forces will form between solute and solvent. “like dissolves like”
- Wax molecules
- WAX: C-C-C-C-C-C (surrounded by H) - nonpolar

** Easily slip from solid
** Slide between benzene molecules
** Form new London forces between solvent and solute

- Energy change associated with formation of solution
- Molar Enthalpy of Solution (ΔHsoln)
ex. If you mix stuff in water sometimes gets hot and sometimes cold; sodium hydroxide in water (becomes hot) is exothermic releasing that much energy — must put water in a ice bath

- Difference in intermolecular forces between isolated solute and solvent and mixture
- Cost of mixing
- Energy exchanged between system and surroundings

Amount of enthalpy exchanged when one mole of solute dissolves in a solvent at constant pressure to make a solution
** amount of heat absorbed or released be the solution

ΔHsoln > 0 (positive)
ΔHsoln < 0 (negative)
- Occurs depends on your system
ex. positive = forces energy into the solution (endothermic)
negative = releases energy to surrounding (ex. to your hand) - decreases potential energy of the system (exothermic)

- Costs energy to make solution
- Endothermic
- Increase in potential energy of system

- Energy given off when solution is made
- Exothermic
- Decrease in potential energy of system

Formation of solution from solid and liquid can be modeled as two-step process

- Separate solute and solvent molecules
** Break intermolecular forces
** Endothermic, increase in potential energy of system, “cost”

- Mix solute and solvent
** Come together
** Form new intermolecular forces
** Exothermic, decrease in potential energy of system, “profit”

- Application of Hess’ Law
** Way to take two things we can measure and use to calculate something we can’t directly measure
** ΔHsoln is path independent
** Method works because enthalpy is state function
** ΔHsoln = Hsoln + Hsolute + Hsolvent
- Overall, steps take us from 
solid solute + liquid solvent → final solution
- These steps are not the way solution 
would actually be made in lab
ex. positive, positive + negative = sum of all the energy chains (solute solvent solution)

** ΔHsolution = ΔHlatt + ΔHhydr
** ΔHsoln = 632 kJ mol–1
– 619 kJ mol–1
ΔHsoln = +13 kJ mol–1
Formation of KI(aq) is endothermic

ΔHlattice (NaBr) = 728 kJ mol–1
ΔHhydration (NaBr) = –741 kJ mol–1
ΔHsolution = ΔHlatt + ΔHhydr
ΔHsoln = 728 kJ mol–1
– 741 kJ mol–1
ΔHsoln = – 13 kJ mol–1
Formation of NaBr(aq) is exothermic
** put in some to get more back

The one that form the strongest intermolecular forces between IT and water: (higher the charge the stronger the intermolecular force between IT and the ion)

- One in which inter-
molecular attractive 
forces are identical
ΔHsoln = 0
- e.g., Benzene in CCl4
- All London forces
- ΔHsoln ~ 0
** Step 1 + Step 2 = –Step 3
** or
ΔHsolute + ΔHsolvent = –ΔHsolvation
** get back the same amount of energy that you put into the solution

- Only very weak attractions exist between gas molecules
** There are no intermolecular attractions in ideal gases
- When making solution with gas solute
** Energy required to “expand solute” is negligible
- Heat absorbed or released when gas dissolves in liquid has two contributions:

1. Expansion of solvent ΔHsolvent
2. Solvation of gas ΔHsolvation
- WHY: gases are already separated
- gas is already separated from its neighbors
- expansion of solvent (breaking the intermolecular forces)

** Generally endothermic
** ΔHsolvation > 0

** Generally exothermic (** Talking about sign numbers not absolute values — the lattice is ALWAYS POSITIVE )
** ΔHsolvation < 0
- ammonia is a gas that is extremely polar, so it could be dissolved in water

** A measurement of the mass of solute that forms saturated solution with given mass of solvent at specified temperature
- How much stuff could you dissolve in a specified about of substance in a specific temperature (is changing)

One that has dissolved as much solute as it can under existing conditions
- Existing Conditions: the temperature and pressure of the room (specific to our classroom)
ex. Access solute the solution above is saturated, because the solution above would be dissolved if that was not the case

- NaCl (aq) <—> NaCl (s) (mean it is saturated) — solid is at the bottom of the beaker (for one of the solid crystals to dissolve one of the NaCl (aq) must precipitate - one comes out for one to come in — solution equilibrium)
- How do we know this happens? allow to sit for long enough will grow to one whole solid

Dynamic equilibrium exists between solute in two phases (solid & solution)
Molecules moving in & out of solution

** Most substances become more soluble as T increases
** Amount solubility increases
- Varies considerably
- Depends on substance

- Solubility of gases usually decreases as T increases
- Solubilities of Common Gases in Water

- Pressure-Solubility Law
- Concentration of gas in liquid at any given temperature is directly proportional to partial pressure of gas over solution
Cgas = kHPgas
ex. partial pressure of the gas in the soda bottle example - CO2 goes into the head space if we have a MIXTURE of solution — 80%s atm = 4 atm (at the top - helping to redissolve CO2 — keeping the soda fizzy) in a solution

CO2 will will react with water to create carbonic acid (seltzer water) — carbon dioxide itself is tasteless

Solubility increases as P increases
ex. If your concentration of the gas increases so does P (keeping the ratio)
If you double one, double the other

A. At some P, equilibrium exists between vapor phase and solution
ratein = rateout
B. Increase in P puts stress on equilibrium
Increase in frequency of collisions so ratein > rateout
C. More gas dissolved
rateout will increase until rateout = ratein again

- Gas molecules with polar bonds are much more soluble in water than nonpolar molecules like oxygen and nitrogen
CO2, SO2, NH3 >> O2, N2, Ar
- Form H-bonds with H2O
- Some gases have increased solubility because they react with H2O to some extent

“like dissolves like”
- CO2 is non polar as a molecule, however it does have partially polar bonds — so some of it could be dissolved, just not much
- However, SO2 is actually polar with partially + and -
more CO2 reacts with water in order to replace — maintaining the positive concentration

- how to you keep CO2 molecules dissolved? high pressure
- gases are more soluble at low temp and high pressure

M = mol of solute/L of solution
- M varies with temperature
- Volume varies with temperature
- Solutions expand and contract when heated and cooled
- If temperature independent concentration is needed, must use other units
- When you change the temp — you change the solubility — meaning, the molarity of your saturated solution would also change

Temperature Independent Concentration
Percent Concentrations: Also called percent by mass or percent by weight (in respect to the solution)
Percent by mass = mass of solute/ mass of solution x 100
- This is sometimes indicated %(w/w) where “w ” stands for weight or mass
- The “(w/w)” is often omitted

Temperature Independent Concentration
- Number of moles of solute per kilogram solvent
molality = m = mol of solute/kg of solvent
- Also molal concentration
- m vs. M
Similar when d = 1.00 g/mL (DENS FOR WATER)
Different when d >> 1.00 g/mL or 
 d << 1.00 g/mL

Temperture Idependent Concentration
XA = # mol A/ Total moles of all component
Mole %:
mole%A = XA x 100%

- Physical properties of solutions
- Depend mostly on relative populations of particles in mixtures
- Don’t depend on their chemical identities
Effects of solute on vapor pressure of solvents:
- Solutes that can’t evaporate from solution are called nonvolatile solutes
Fact: All solutions of nonvolatile solutes have lower vapor pressures than their pure solvents

- not pure water: effecting vapor pressure (usually lowers VP)
- less vapor = less vapor pressure
- lower vapor pressure will effect BP and FP of a solvent
- FP = DEPRESSED = DEPRESSION (how much will it be depressed? depends on how much you dissolve in water)
- boil and a higher temp — with more
- salt water changing the specific heat changing the temp of gas
- Nonvolatile = does not evaporate usually referring to solid and ionic compounds

- Vapor pressure of solution, Psolution, equals product of mole fraction of solvent, Xsolvent and its vapor pressure when pure, Psolvent
Applies for dilute solutions: P (solution) = X (solvent) P (Solvent)
ex. sugar and water = 90% is water 10% is sugar — mole fraction of sugar is .10 mole fraction of water is .90
VP will be .90 bc the rest will NOT EVAPORATE
ex. vapor pressure is 30 torr 30 x .90 = 27 (new VP)
0 means under standard conditions of the solvent

- Plot of Psoln vs. Xsolvent 
should be linear
** Slope = X (solvent) P (Solvent)
Intercept = 0
- Change in vapor 
pressure can be 
expressed as: delta P = P (solvent) - P (solution)
- Usually more interested in how solute’s mole fraction changes the vapor pressure of solvent
** change in pressure = mole fraction of solute x normal VP of water

Freezing point of solution always lower than pure solvent = freezing point depression
Boiling point of solution always higher than pure solvent = boiling point elevation

Consider the phase diagram of H2O
Solid, liquid, gas phases in equilibrium
Blue lines
P vs. T

The salt lowers the freezing point to change the freezing point & the boiling point is increased (cars have an engine that produces tremendous amount of heat — too much heat — potential energy converted into heat energy — need to cool so energy does not overheat and break — there is a cooling system that goes around the engine within the heat is water which absorbs heat from the energy bc it has a high specific heat capacity — but this is why water freezes so engine then freezes); On cold days your water does not expand — elevates the boiling point — antifreeze is stuff dissolved in water.