exam 1

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Solutions: homogeneous mixtures – the solvent dissolves the solute... since the solvent is the component present in greatest amount, it determines the phase of the solution...

in relative terms: a concentrated solution contains a lot of solute, dilute solution contains little solute

Enthalpy of solution (Hsoln) and intermolecular forces in solution determine whether a solution can form:

When all forces are similar, Hsoln = 0: ideal solution (also: Vsoln = 0)

Solute–solvent interactions dominate: Hsoln < 0: solution is non-ideal, and forms exothermically

Solute–solute and solvent–solvent interactions dominate: Hsoln > 0: solution is non-ideal, and forms

endothermically (if it forms)

Solute–solute and solvent–solvent interactions are >> solute–solvent interactions: Hsoln >> 0 and

solution can’t form

Rule of thumb for predicting solubility: “like dissolves like” (not always true, but a good guide)

For ionic compounds in H2O: ion–dipole forces (+ end of H2O dipole attracts anions; – end of H2O dipole attracts cations) disrupt the solid  hydrated ions in solution

Dissolving too much solute  saturated solution – a dynamic equilibrium is established: continuous dissolving and crystallizing, but apparently nothing happening (concentration of dissolved solid remains constant; quantity of undissolved solid remains constant)

Solubility of ionic compound = how much is dissolved at saturation... for most ionic compounds: solubility in H2O increases with temperature...

Solubility of gases (in H2O) decreases with increasing temperature, but (in all solvents) increases with increasing pressure... Henry’s Law: solubility is proportional to pressure, Cgas = kPgas (know this equation)

CONCENTRATION

− always defined as “amount of solute”/“amount of solution” (or “amount of solvent” for molality)

(you need to know all of the following equations/definitions:) molarity = moles solute (mol) / volume of solution (L)

for % compositions: use “(part/whole)  100” with appropriate units (mass, volume, etc...). ppm, ppb, ppt are similar to %: defined as (part/whole)  106 or 109 or 1012, respectively

molality = moles solute (mol) / mass of solvent (kg) ; molality is temperature-independent, so is useful when temperature varies

mole fraction = (moles of component)/(total moles present); mole % = mole fraction  100

Conversions between different units of concentration are best achieved by first deciding upon a fixed amount of solution. Suggestions:

if original concentration is: mass %

volume % molarity (M) molality (m) mole fraction ()

start with:

100 g of solution

100 mL of solution

1.00 L (i.e., 1000 mL) of solution 1.00 kg (i.e., 1000 g) of solvent 1 mol of (solute + solvent)

Colligative properties – depend on quantity (concentration) of solute particles, not on their identity

 Raoult’s Law: Psolution = solventP°solvent (know this equation) – adding a solute lowers vapor pressure compared to the vapor pressure of the pure solvent (remember to use solvent!)

 Freezing point depression: Tf = mKf (f.p. is lower than pure solvent) (know this equation)

 Boiling point elevation: Tb = mKb (b.p. is higher than pure solvent) (know this equation)

 Osmosis: concentrated solution/dilute solution separated by semi-permeable membrane; solvent passes from

dilute  concentrated; flow continues until osmotic pressure, , is reached.  = MRT or (know these equations)... two solutions with same value of  are isotonic (isosmotic)

Solutions of electrolytes

Electrolytes produce ions in solution, so # particles in solution > # formula units dissolved... Equations for colligative properties must be modified by van’t Hoff factor (i):

ΔTf =imKf ΔTb =imKb =iMRT

V = nRT

(for Raoult’s Law, the equation is the same, but x must be calculated based on numbers of moles of all the particles in solution – remember that there are ions present)

(know how to modify and use the equations) i (at infinite dilution) = number of ions produced in solution per formula unit of solute

For more concentrated solutions, ion pair formation reduces the effective value of i

Colloids are mixtures with particle size between that of solutions and suspensions – they won’t settle out... they can be distinguished from true solutions by the Tyndall effect (light passing through colloid is scattered in all directions)