Chem Exam 2

Homogenous mixture that consists of one or more solutes uniformly dispersed at molecular or ionic level throughout a medium known as the solvent • Not always liquid Ex: Air: solution of nitrogen, oxygen, other gases Former dental fillings: silver amalgams (mix of silver and mercury)

Not possible to discern boundaries between components of the mixture. Ex: Normal saline (uniform throughout sample)

Separates regions of a mixture where the chemical or physical properties of the mixture change

Material that got dissolved; component kf solution present in smaller quantity. Ex: Sodium chloride

Material that does dissolving, usually water

Moles of solute per liter of solution M = mol of solute/L of solution (mol/L) • Molar concentrations are conversion factors btw moles of material and liters of solution • Depends on temperature of solution Molar concentration will decrease at temperature increases

Express concentration in terms of moles of solute per kilogram of solvent m=mol of solute/kg of solute • Used in physical chemistry -> quantities of solute and solvent are considered separately • Mass is not temperature dependent, molality is not temperature dependent • Less convenient in analysis bc quantities of a solution measured out by volume or mass in lab include both solute and solvent • When doing stoichiometry, molality requires additional calculation to take into

Molality and molarity are never equal, but difference is is smaller as solutions become more dilute Convert btw molarity and molality, need to know density of solution molality of solution = x mol of solute/y kg of solvent

The percent of concentration you encounter in a clinical setting when measuring out a volume of medicine in syringe Grams of solute per 100mL of solution %w/v = g of solute/100mL of solution OR %w/v= g of solute/mL of solution (100%) First equation is useful as a conversion factor btw grams of solute and milliliters of solution Second equation is useful tp calculate concentration of a solution • To relate to percent by weight to percent weight to volume, use the density of solution Percent by volume (%v/v): never used in analytical lab bc values are not additive

Analogous to a mole

Analogous to molarity

1 equivalent of a substance contains 1 mole of chemical reactivity N= equivalents of solute/L of solution Normality is equal to equivalents of solute per liter of solution Unless context of chemistry is specified, normality is ambiguous

Concentration of extremely dilute solutions is sometimes expressed as part per million A ppm concentration is analogous to a percent concentration, except you are comparing amount of solute to a million grams of solute instead of 100 gram. ppm = g of solute/1x10^6g of solution = g of solute/g of solution (1x10^6)ppm

Some solutes are more soluble in given solute than others The solubility of solute is amount of solute that will dissolve in a given amount of solvent at a given temp Saturated solution contains maximum amount of a solute, as defined by its solubility. No more solute will dissolve in a solution saturated with solute. if solution is not saturated, more solute will dissolve in solution.

Supersaturated solution contains more solute than allowed by solubility of solute. • NOT a stable system, bc there is more solute dissolved in sample than solvent can accommodate • Excess solute will come out of solution, crystallizing as a solid, separating as a liquid, or bubbling out as a gas Ex: when blood or urine in kidneys become supersaturated with calcium oxalate or calcium phos, kidney stone can form • If solute is a gas in liquid solvent, will see bubbles forming in solution (Fizz when opening bottle of beer or soda)

Two liquids are miscible if they are soluble I each other in all proportions Ex: Alcohol and water are miscible with each other. Oil and water are immiscible Like dissolves like • Polar solutes are more soluble in polar solvents • Nonpolar solutes are more soluble in non polar solvents; insoluble in water • Most organic molecules are nonpolar • Most ionic compounds are polar -> dissolve in water but not soluble in organic compounds (ether, hexane)

When a solute dissolves in a solvent, there is an associated energy change and temperature change to solution The energy change when using heat and cold packs is called heat of solution or the enthalpy of solution: Δhsoln Defined as the energy change that accompanies dissolving exactly 1 mole of solute in a given solvent Enthalpy H is equal to the heat Q as long as the pressure remains constant The energy change may be endothermic or exothermic Whether the heat of solution is endothermic or exothermic depends on the relative magnitudes of the lattice energy and the heat of solvation * If tearing the ions apart requires more energy than is released by solvation, then “Δ”Hsoln is going to be positive (endothermic) • If the energy released by solvation is greater than the energy required to tear the ions apart, “Δ”Hsoln is going to be exothermic

• If the solution process is exothermic, energy flows out of the system (solvent and solute) into the surroundings, resulting in a temperature increase in the solution • Energy is lost

• If the solution process is endothermic, energy flows from the surroundings into the system, resulting in a temperature decrease in the solution • Energy is gained

As pressure increases, the solubility of a gaseous solute in a liquid solvent increases * With less pressure, solubility of gaseous solute decreases (CO2 escaping champagne bottle) • Since solids and liquids are not very compressible, at least not compared to gases, pressure has very little effect on the solubility of solid and liquid solutes The quantitative relationship between pressure and solubility is given by Henry’s law: S = k(H)P(gas) S = solubility kH = Henry’s law constant (0.042g/L/atm Pgas = partial pressure of the gas Increased partial pressure of gas -> more gaseous molecules zipping around at near surface of the liquid

The solubility of solid and liquid solutes in liquid solvents generally increases with increasing temperature • As temperature increases, vapor pressure of gaseous solutes increases to point that they escape the solvent into gas phase

The vapor pressure of a solution Decreases with increasing solute concentration The vapor pressure of a liquid results from the most energetic molecules near the surface of the liquid escaping into the gas phase The most likely escape sites for the liquid molecules are at or near the surface of the liquid As we begin to introduce solute molecules, some of these escape sites are occupied by the solute molecules, so fewer solvent molecules can escape into the gas phase Therefore, the vapor pressure of the solution is less than the vapor pressure of the pure solvent Raoult’s law states the vapor pressure of a volatile component of a solution (P) is equal to the vapor pressure of the pure substance (Po) times the mole fraction (ᵡ) of that substance P = χPo

The boiling point of a solution INCREASES with increasing solute concentration Temperature at which the vapor pressure of the material is equal to the ambient pressure The boiling point of a solution increases as the concentration of solute(s) increases The change in boiling point is directly proportional to the molal concentration of the solute particles ΔTbp = Tbp,solution – Tbp,solvent = kbp · mtotal ΔTbp = the number of degrees by which the boiling point increases Tbp,solution = the boiling point of the solution, Tbp,solvent = the boiling point of the pure solvent kbp = a constant (called the ebullioscopic constant) that is characteristic of the solvent mtotal = the molal concentration of all solute particles Activity is the effective concentration of a solute Activity is always less than molality

The freezing point of a solution DECREASES with increasing solute concentration Temperature at which the liquid phase of the material is in equilibrium with the solid phase In order to enter into the solid state, the molecules (or ions or atoms) of the sample need to settle into an orderly, crystalline lattice structure The presence of solute particles interferes with this process by getting in the way It is necessary to cool the sample to lower temperatures, thereby lowering the kinetic energy of the molecules even further, before they will settle into the solid phase The relationship that quantifies the degrees of freezing point depression has an identical form to boiling point elevation ΔTfp = Tfp,solution − Tfp,solvent = kfp · mtotal ΔTfp = the number of degrees by which the freezing point decreases Tfp,solution = the freezing point of the solution, Tfp,solvent = the freezing point of the pure solvent kfp = a constant (called the cryoscopic constant) that is characteristic of the solvent mtotal = the molal concentration of all solute particles