Final - Solutions

solution

homogeneous mixture of two or more substances or components

solvent

the majority component in a solution (usually H2O)

solute

the minority component in a solution

solubility

the maximum amount of a substance that will dissolve in a given amount of solvent at a specified temperature

* typically given as a concentration

aqueous solutions

water as the solvent

* can have any other soluble solid, liquid, or gas as a solute

gas as solutions

solute - gas

solvent - gas

ex - air (O2 and N2)

liquid - g and l

solute - gas

solvent - water

ex - club soda (CO2 and H2O)

liquid - l and l

solute - liquid

solvent - liquid

ex - liquor (ethanol and water)

liquid - s and l

solute - solid

solvent - liquid

ex - sea water (salt and H2O)

solid as solutions

solute - solid

solvent - solid

ex - brass (copper and zinc)

* called alloys

entropy

measure of energy randomization or energy dispersal in a system

* disorder of a system
* mixtures will mix randomly if other forces do not prevent (prefers disorder)
* driving force behind creation of solutions

what forces keep mixtures from randomly mixing?

intermolecular forces stop formation of solutions

how IMFs stop formation of solutions

if IMFs of individual species are much stronger than that of interspecies mixing, then a solution can’t occur

* like dissolves like

ion-dipole forces

allows the solution to be formed

* polar species can also often dissolve ionic compiunds
* aqueous solutions

determining solubility

both need to be polar or non-polar to be soluble

enthalpy of solution

∆H(sol) = ∆H(solute) + ∆H(solute-solvent) + ∆H(solvent)

∆H(sol) = (+) + (-) + (+)

enthalpy of solutions - solids and liquids

solute - overcomes IMFs with ∆H

* in solution, there is a lower state of energy than the solute and solvent separately

∆H(sol) - endothermic

|∆H(solute) + ∆H(solvent)| > |∆H(solute- solvent)|

* + > -
* feels cold
* + ∆H(sol)

∆H(sol) - exothermic

|∆H(solute) + ∆H(solvent)| < |∆H(solute- solvent)|

* feels hot
* - ∆H(sol)

insoluble

if the solution cannot dissolve most of the solute

soluble compounds have ____ solubility

high

insoluble compounds have very very ___ solubility

low

saturated

a solution that has this dynamic equilibrium present

* combo of dissolved and crystallized
* unsaturated = less than this

supersaturated

a solution with more than the saturated levels, which cna only occur under particular conditions such as by heating

* if cold, the extra will crystallize

equilibrium

when a solid substance is in a saturated solution

* rate of dissolution = rate of recrystallization

temperature and solubility of solids in liquids

increasing temp = increases solubility of a solid in a liquid

* IMFs more easily overcome

temperature and solubility of gas in liquid

higher temperature = less solubility

* IMFs between the fas and liquid are mroe easily overcome as the temperature increases

pressure and solubility of gas in a liquid

henry’s law

* increasing pressure = increases solubility
* S(gas) = kP
* solubility of gas = constant (Pressure)

molarity (M)

moles solute/ L solution

molality (m)

mols solute/kg solvent

mole fraction (x)

mols solute/total mols

* units = none

mole percent

mols solute/total mols x 100

parts by mass

mass solute/ mass soluteion x mult factor

percent by mass

mass solute/ mass solution x 100

* units = %

parts per million by mass (ppm)

mass solute / mass solution x 10^6

* units = ppm

parts per billion by mass

mass solute / mass solution x 10^9

* units = ppb

parts by volume (%, ppm, ppb)

volume solute / volume of solution x mult factor

colligative property

a property that depends on the number of particles dissolved in a solution rather than the type of particle

* vapor pressure lowering, freezing point depression, boiling point elevation, and osmotic pressure

electrolytes

dissociated in water to form ions

* ionic compounds = electrolytes
* ex - NaCl (s) → Na+ (aq) + Cl- (aq)

non-electrolytes

covalent = non-electrolyters

C6H12O6 (s) → C6H12O6 (aq)

van’t hoff factor (i)

ratio of particles in solution to formula unit

* NaCl = 2 (Na + Cl)
* MgCl2 = 3 (Mg + 2Cl)
* C6H12O6 = 1

raoult’s law

P(solution) = X(solvent) P°(solvent)

* P(solution) = vapor pressure of the solution
* X(solvent) = mole fraction of the solvent
* P °(solvent) = vapor pressure of the pure solvent

calculate the vapor pressure of each mixture and total vapor pressure

1. find moles of each mixture
2. calculate the mole fraction of each
3. plug into P(a) = X(a) P°(a) for each
4. add together for total

phase diagram

x-axis = temp (C)

y-axis = pressure (atm)

clockwise from top left = solid, liquid gas

phase diagram - freezing temperatures lower

△T(f) = T(solvent) - T(solution)

△T(f) = i \* m \* K(f)

* i = van’t hoff factor
* m = molality (mol A/ kg solvent)

phase diagram - boiling points raise

△T(b) = i \* m \* K(b)

△T(b) = T(solution) - T(solvent)

* opposite △T(f)

osmosis

flow of solvent from area of low solute concentration to area to high solute concentration

* moving from low to high concentration across a semi-permeable membrane

osmotic pressure

the amount of pressure that would be required to push on a pure solvent to prevent it from passing thru the membrane into a solution

* pressure that prevent them from moving

osmotic pressure - equation

⫪ = MRT

* M = molarity (mol A/L sol)
* R = ideal gas constant
* T = temp (K)

isosmotic/ isotonic

osmotic pressure equal to body fluids (a)

ex - normal red blood cells

hyperosmotic/ hypertonic

osmotic pressure greater than body fluids

* higher solute concentration
* ex - shriveled red blood cells

hyposmotic/ hypotonic

osmotic pressure less than that of body fluids

* lower solute concentration
* ex - swollen red blood cells