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Chem.200-13.Solution
General Chemistry Ch. 13 - The Properties of Mixtures: Solutions and Colloids
| Question | Answer |
|---|---|
| Solute | The substance that dissolves into the solvent |
| Solvent | The solution into which the solute dissolves |
| Solubility (S) | The max amount that dissolves in a fixed quantity of a particular solvent at a specified temperature, given that excess solute is present. |
| Is the solvent always the most abundant component of a given solution? | Usually, but not always. In some cases the substances are miscible. |
| Miscible | If substances are miscible it means they’re soluble in each other in any proportion. In such cases it may not be meaningful to call one the solute and the other the solvent. |
| Dilute/concentrated | Contains much less dissolved solute than the concentrated one. Dilute and concentrated are qualitative terms, unlike solubility which is quantitative. |
| Side note: butter doesn’t dissolve in water but it dissolves in… | … cooking oil |
| How do intermolecular forces relate to solubility? | Substances with similar types of intermolecular forces dissolve in each other. Rule of thumb: “like dissolves like” |
| Dissolve: ionic compounds in water | Water breaks apart the ionic compounds and form hydration shells around the ions |
| The size of the hydration shell (i.e. the number of water molecules surrounding the ion) depends on | The ions size |
| Dissolve: Hydrogen bonds in water | H-bonds are the primary factor in the solubility of many biological compounds in water including alcohols, sugars, amino acids, etc. |
| Dissolve: Dipole-dipole forces in non-aqueous solvents | In the absence of H-bonds dipole-dipole forces account for the solubility of polar organic molecules such as acetaldehyde (CH3CHO) in non polar aqueous solutions like chloroform (CHCl3) |
| Dissolved Ion-induced dipole forces example | Example: Fe^2+ (iron ion) in hemoglobin binding to O2 in the bloodstream. |
| How do dissolved ion-induced dipole forces affect solubility? | Because an ion increases the magnitude of any nearby dipole, ion-induced dipole forces also contribute to the solubility of salts in less polar solvents. |
| Dissolved dipole-induced dipole forces | Weaker than ion-induced. E.g. the solubility of O2, N2, and the noble gases in water is due to this. Paint thinners and grease solvents function this way as well. |
| Dissolved dispersion forces | Contribute to the solubility of all solutes in all solvents. |
| Oil doesn’t dissolve in water but does dissolve in hexane. Why? | In water: they’re not like substances, H-bonds can’t be substituted by the dispersion forces of the oil and vice versa. In hexane: they both have a lot of dispersion forces which can be substituted for one-another. |
| Alcohol’s behavior as a solute in water | Alcohol CH3(CH2)nOH = hydrocarbon with a hydroxyl group. The hydrocarbon portion is hydrophilic and not very soluble in water; the hydroxyl group however has H-bonds and is very soluble. n is proportional to its size |
| Alcohol’s solubility in water and hexane as related to the size of the molecule | Small: very soluble in water but not in hexane (the hydroxyl group has more influence). Large: not very soluble in water but very soluble in hexane (the hydrocarbon portion dominates and dispersion forces become more important) |
| Why is N2 not very soluble in water | It’s nonpolar. The solute-solvent forces are weak. |
| Why does O2 seem more soluble in blood than water? | Because it binds to the iron in hemoglobin |
| Some nonpolar substances appear to dissolve more readily in solvents due to… | The fact that they’re actually involved in reactions, not just dissolving. E.g. O2 in blood, or CO2 in water (because it’s actually reacting with H2O to form H+ and HCO3-) |
| Gas-Gas solutions | All gases are infinitely soluble in one another. E.g. air which is a solution containing about 18 different gases. |
| Gas-solid solutions | When a gas dissolves in a solid, it occupies the spaces between the closely packed particles. |
| Interesting fact: how to purify H2 gas? | By passing an impure sample through a solid metal such as palladium. Only the H2 molecules are small enough to enter the spaces between the Pd atoms. |
| Solid-Solid solutions | Because solids diffuse so little, their mixtures are usually heterogeneous; an example is gravel mixed with sand. Some solid-solid solutions can be formed by melting the solids and then mixing them, e.g. alloys. |
| Substitutional alloys | e.g. brass and sterling silver, atoms of another element substitute for some of the main element’s atoms. |
| Interstitial alloys | E.g. carbon steel, atoms of another element fill some of the spaces between atoms of the main element. |
| Proteins | Unbranched polymers formed from monomers called amino acids. They range from 50 amino acids to several thousand. Protein shapes are determined completely by the sequence of amino acids in the chain. |
| Amino acids | Organic compounds with amine (-NH2) and carboxyl (-COOH) groups in the same molecule. There are 4 groups attached to the alpha-carbon. The amine group, carboxyl group, H atom, and a side chain. |
| Side chain | A molecular structure ranging from H atom to C9H8N. |
| Peptide bond | The covalent linkage between amino acids in a protein and is formed between the carboxyl group of one amino acid and the amine group of the next. |
| In essence, the backbone of a protein consists of | An alpha-carbon bonded to a peptide bond which is bonded to the next alpha-carbon and so forth with various side chains dangling off them. |
| Disulfide bridge | The -SH ends of two cysteine side chains often form an -S-S- bond, a covalent “disulfide bridge” that brings together distant parts of the chain. |
| Salt link | Sometimes oppositely charged side chains lie near each other and the –COO- and –NH3+ groups form an electrostatic salt link (or ion pair), which secures the chain’s bends. |
| The amino acid sequence determines a protein’s shape, and the shape determines the… | Protein’s function |
| Soap | The salt of a strong base (metal hydroxide) and a fatty acid (a carboxylic acid with a long hydrocarbon). Soap has an even number of carbons: an odd number in the long chain, and a single one at the head (-COO-). |
| How does soap help clean grease on your hands? | The nonpolar tails of the soap molecules interact with the nonpolar grease molecules through dispersion forces, while the head interact with water. The tails of the soap pierce the greasy clumps while the heads poke out. |
| Detergent vs. soap | Detergents are very similar except the heads are composed of an –SO3- group rather than -COO- |
| Phospholipids | Molecules consisting of two long fatty acid chains and a charged organophosphate group linked to glycerol, a three-carbon trialcohol. In head, the head is polar, the tail is not. |
| In water, what do phospholipids do? | Self assemble into lipid bilayers, a sheetlike double layer of molecules where the polar heads cover the surfaces and the nonpolar tails face each other on the interior. |
| How do proteins orient themselves on lipid bilayers? | Their nonpolar regions orient themselves on the interior, the polar regions on the exterior. |
| Action of antibiotics | E.g. gramicidin and others: tubular structured molecules that are nonpolar on the outside, polar on the inside embed themselves perpendicularly through the lipid bilayers and allow ions to diffuse through, imbalancing and killing the cell. |
| Nucleic acids | Unbranched of polymers consisting of monomers called mononucleotides. |
| Mononucleotides | Consists of an N-containing base, a sugar, and a phosphate group. In DNA: sugar = 2-deoxyribose. |
| The repeating patter of a DNA chain | Sugar linked to phosphate linked to sugar linked to phosphate, etc. Attached to each sugar is one of four nitrogen-containing bases. The bases are ring-shaped and dangle off the sugar-phosphate chain similar to side-chains of proteins. |
| Structure of double helix | On the exterior are the sugar-phosphate chains forming ion-dipole and H-bonds with the aqueous surroundings. On the interior are the bases complimentarily H-bonding with each other. |
| Polysaccharides | Polymers containing monomers called monosaccharides (or simple sugars). |
| Three major polysaccharides: | Cellulose, starch, and glycogen all consist entirely of glucose monomers but differ in the way they’re linked and extent of crosslinking. |
| Cellulose | The only unbranched polysaccharide of the three and consists entirely of long chains of glucose monomers, i.e. parallel chains of linked glucose units lying next to each other in three-dimensions, their –OH groups everywhere. |
| The strength of wood is primarily derived from | The countless H-bonds among cellulose chains which promote H-bonding and dispersion forces. |
| For one substance to dissolve another, three events must occur: | (1) solute particles must separate from each other, (2) some solvent particles must separate to make room for the solute particles, and (3) solute and solvent particles must mix together. |
| No matter what the nature of the attractions within the solute and within the solvent… | Some energy must be absorbed for particles to separate, and some energy is release when they mix and attract each other. |
| Solution process, step 1 | Solute particles separate from each other. This step involves overcoming intermolecular attractions, so it is ENDOTHERMIC: Solute (aggregated) + heat -> solute (separated). deltaH_solute > 0 |
| Solvent particles, step 2 | Solvent particles separate from each other. This step also involves overcoming attractions, so it is ENDOTHERMIC: Solvent (aggregated) + heat -> solvent (separated) deltaH_solvent > 0 |
| Solvent particles, step 3 | Solute and solvent particles mix. The particles attract each other, so this step is EXOTHERMIC. Solute (separated) + solvent (separated) -> solution + heat. deltaH_mix < 0 |
| Heat of solution (deltaH_soln) | The total enthalpy change that occurs when a solution forms from solute and solvent |
| Thermochemical solution cycle formula | (Yet another application of Hess’ law): deltaH_soln = deltaH_solute + deltaH_solvent + deltaH_mix |
| The solution process is exothermic if… | (deltaH_solute + deltaH_solvent) < |deltaH_mix|. Thus, deltaH_soln would be negative |
| If deltaH_soln is higly positive then… | The solute may not dissolve to any significant extent in that solvent. |
| The deltaH_solvent and deltaH_mix components of heat of solution are difficult to measure individually. Combined these terms represent the enthalpy change during… | solvation |
| Solvation | The process of surrounding a solute particle with solvent particles. |
| Hydration | Solvation in water. |
| Heat of hydration (deltaH_hydr) | The enthalpy change for separating the water molecules (deltaH_solvent) and mixing the solute with them (deltaH_mix) are combined into heat of hydration (deltaH_hydr) |
| Using heat of hydration what does the thermochemical solution cycle formula become? | deltaH_soln = deltaH_solute + deltaH_hydr |
| Heat of hydration exhibits trends based on | Charge density |
| Charge density | The ion’s charge to its volume |
| Heat of hydration trend | The higher the charge density is, the more negative deltaH_hydr is |
| What attracts H2O molecules more strongly: a 1+ or 2+ molecules (if both are similar sized? | 2+ |
| What attracts H2O molecules more: a large 1+ molecule or small one? | Small |
| Periodic trend of charge density (and thus more negative deltaH_hydr) | Down a group: charge density decreases, across a period: charge density increases |
| The energy required to separate an ionic solute into gaseous ions is its … | Lattice energy, thus deltaH_soln = deltaH_lattice + deltaH_hydr |
| Relate lattice energy to the temperature change during the process of dissolving the solute | If the lattice energy is about the same as the heats of hydration, there will be no noticeable temperature change. If the lattice energy is much smaller, then very exothermic and hot. If the lattice energy is much larger, then endothermic/hot. |
| Entropy (S) | Directly related to the number of ways that a system can distribute its energy, which in turn is closely related to the freedom of motion of the particles and the number of ways they can be arranged |
| Compare entropy of liquid vs. solid | Solids have particles that are fixed in their positions, whereas in liquids they’re free to move around. This greater kinetic energy means they’re free to move around each other and distribute their kinetic energy in more ways. |
| Compare entropy of gas and liquid | S_gas > S_liquid |
| A solution usually has higher entropy than… | The pure solute and pure solvent, thus: S_soln > (S_solute + S_solvent), or deltaS_soln > 0 |
| He solution process involves the interplay of two factors: | The change in enthalpy and the change in entropy |
| Systems tend toward … | A state of lower enthalpy and higher entropy |
| At what point is a solution saturated? | When equilibrium has been reached between the undissolved solute and the dissolved solute. That is, the number of solute particles dissolving per unit time equals the number recrystalizing. |
| What happens if you add more solute to a saturated solution? | It will not dissolve. |
| Unsaturated | Solute will still dissolve until fully saturated |
| Supersaturated | A solution is supersaturated when it contains more than the equilibrium concentration of dissolved solute. At this point, if you add just a small amount of “seed” crystal of solute, or just tap the container, it will crystallize. |
| How can a supersaturated solution be prepared? | By dissolving solute in the material while it’s warmed up because it has higher solubility at higher temperature. Then cool the material down and it will be supersaturated because it has exceeded its solubility at the cool temperature. |
| Are all solids more soluble at higher temperatures? | No, MOST are. Seom compounds, e.g. Ce2(SO4)3 have decreased solubility at higher temperatures. Many sulfates behave similarly. |
| Gas solubility in water _______ with rising temperature | Decreases. Why? Gases have weak molecular forces so there are relatively weak interactions between gas and water. Thus, when temperature increases, gas is able to easily break free and return to gas phase. |
| Equation that shows gas solubility in water decreases with rising temperature | Solute(g) + water(l) -> saturated solution(aq) + heat |
| Thermal pollution | Industrial plants will heat up water during processes and dump it back into lake. The solubility of O2 will decrease and this will harm aquatic life in the area. |
| Pressure’s effect on liquid and solid solubility | Very little because liquids and solids are virtually incompressible. |
| Henry’s Law | Henry’s law expresses the quantitative relationship between gas pressure and solubility: the solubility of a gas (S_gas) is directly proportional to the partial pressure of the gas (P_gas) above the solution: S_gas = k_H * P_gas |
| k_H | Henry’s law constant. |
| Molarity | Amount (mol) of solute / volume (L) of solution. I.e. The number of moles of solute dissolved in 1 L of solution. |
| Why would expressing concentration in terms of molarity have drawbacks (1)? | Since volume is affected by temperature so is molarity. A solution expands when heated, so a unit volume of hot solution contains slightly less solute than a unit volume of cold solution. |
| Why would expressing concentration in terms of molarity have drawbacks (2)? | Because of solute-solvent interactions that are difficult to predict, solution volumes may not be additive; that is, adding 500mL of one solution to another 500mL solution may not give 1000mL |
| Molality | i.e. amount (mol) of solute / mass (kg) of solvent. I.e. the number of moles of solute dissolved in a 1000g (1kg) of solvent. |
| Note regarding the difference between the denominators of molarity and molality | Molarity is concerned with the volume of the SOLUTION whereas molality is concerned with the mass of the SOLVENT. |
| Why does molality use mass rather than volume? | Masses don’t change with temperature and they’re additive. Thus, molality is preferred if temperature (hence density) may change. |
| Special relationship between molarity and molality for water | For water 1L has a mass of 1kg, so molality and molarity are nearly the same for dilute aqueous solutions. |
| Parts by mass | Mass of solute / mass of solution |
| Mass percent | [ Mass of solute / (mass of solution) ] * 100 |
| Sometimes mass percent is symbolized … | As % (w/w), indicating that the percentage is a ratio of weights (more accurately masses). |
| Parts by volume | Volume of solute / volume of solution |
| Volume percent | [Volume of solute / volume of solution] * 100 |
| A common symbol for volume percent | % (v/v) |
| % (w/v) | A measure of concentration frequently used for aqueous solutions is % (w/v), a ratio of solute weight (mass) to the solution volume. |
| Mole fraction (X) | Amount (mol) of solute / (amount (mol) of solute + amount (mol) of solvent) |
| Mole percent | Mole fraction * 100 |
| Colligative properties | Properties of solutions that depend on the number of particles in a volume of solvent and not on the mass of the particles. |
| List colligative properties | Vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure |
| Nonvolatile electrolytes | Solutes that do not dissociate into ions and have negligible vapor pressure, even at the boiling point of the solvent. E.g. tables sugar. |
| Vapor pressure lowering (deltaP) | The vapor pressure of a solution of a nonvolatile nonelectrolyte is always lower than the vapor pressure of the pure solvent. |
| Raoult’s law | P_solvent = X_solvent * P °_solvent, where P_solvent = vapor pressure of solvent above the solution; X_solvent = mole fraction of solvent in the solution, and P °_solvent = vapor pressure of the pure solvent. |
| An ideal solution | One that follows Raoult’s law at any concentration. |
| In practice, Raoult’s law gives… | …a good approximation of the behavior of dilute solutions only, and it becomes exact at infinite dilution. |
| Boiling point elevation | A solution boils at a higher temperature than the pure solvent. |
| Why does boiling point elevation occur? Boiling point = the temp at which its vapor pressure equals the external pressure. And…. | Vapor pressure of a soln is lower than the external pressure at the solvent’s boiling point because the vapor pressure of a solution is lower than that of pure solvent at any temp. |
| Equation for boiling point lowering | deltaT_b = K_b*m, where m is the solution molality and K_b is the molal boiling point elevation constant. I.e. the magnitude of the boiling point elevation is proportional to the concentration of solute particles. |
| deltaT_b | The boiling point elevation; a positive value: deltaT_b = T_b(solution) – T_b(solvent) |
| Freezing point | That temperature at which the solution’s vapor pressure equals that of the pure solvent. At this temperature, the two phases, solid solvent and liquid solution, are in equilibrium. |
| Freezing point depression | deltaT_f. the magnitude is proportional to the molal concentration of solute: deltaT_f = K_f*m. Where K_f is the molal freezing point depression constant. |
| deltaT_f | Freezing point depression; a positive value: deltaT_f = T_f(solvent) – T_f(solution) |
| Semipermeable membrane | A membrane that allows solvent, but not solute |
| The phenomenon caused by a semipermeable membrane is called | Osmosis |
| Osmotic pressure | The applied pressure required to prevent the net movement of water from solvent to solution (symbol = the giant pi thing) |
| The osmotic pressure is proportional to… | The number of solute particles in a given volume of solution, that is, to the molarity. |
| Equation for osmotic pressure | Osmotic pressure = MRT |
| Review: what can solute not do? | They cannot enter the gas phase (leading to vapor pressure lowering), they cannot enter the solid phase (leading to freezing point depression), they cannot cross a semipermeable membrane (leading to osmotic pressure) |
| Review: what leads to the measured colligative property? | The presence of the solute decreases the mole fraction of the solvent, which lowers the number of the solvent particles leaving the solution per unit time; this lowering requires an adjustment to reach equilibrium again. |
| Of the four colligative properties, _____ creates the largest changes and therefore the most precise measurements | Osmotic pressure, note: thus it can be used to determine moles of solute in a solution |
| What is the effect on vapor pressure when the solute is volatile, that is, when the vapor consists of solute and solvent molecules? | P_total = (X_solvent * P°_solvent) + (X_solute * P°_solute). The presence of each volatile component lowers the vapor pressure of the other by making each mole fraction less than 1. |
| The vapor has a higher mole fraction of the… | More volatile solution component. |
| Van’t Hoff factor (i) | The ratio of the measured value of the colligative property in the electrolyte solution to the expected value for a nonelectrolyte solution |
| i formula | i = (measured value for electrolyte solution) / (expected value for nonelectrolyte solution) |
| To calculate the colligative properties of strong electrolyte solutions, we… | incorporate the van’t Hoff factor into the equation: vapor pressure lowering: deltaP = i(X_solute * P°_solvent); boiling point elevation: deltaT_b = i(K_b * m), osmotic pressure = iMRT; etc… |
| If strong electrolyte solutions behaved ideally, the factor i would be… | The amount (mol) of particles in solution divided by the amount (mol) of dissolved solute; that is, i would be 2 for NaCl, 3 for Mg(NO3)2, etc. |
| Are strong electrolyte solutions ideal? | Most strong electrolyte solutions are not ideal. |
| The measured value of the van’t Hoff factor is typically _____ than expected. Why? | Lower. This deviation implies that the ions aren’t behaving as independent particles. The explanation: clustered near a positive ion are, on average, more negative ions than positive ones, and vice versa. |
| Ionic atmosphere | The term describing the ions of opposite charge surrounding each other. I.e. Each ion is surrounded by an ionic atmosphere of net opposite charge. |
| Why is nonideal behavior of solutions more common than nonideal behavior of gases? | Because the particles in solutions are much closer to each other. |
| Suspension | Heterogeneous mixture containing particles large enough to be seen by the naked eye and clearly distinct from the surrounding fluid. |
| In contrast to sand, stirring sugar in water produces a …. | Solution: A homogenous mixture in which the particles are individual molecules distributed evenly throughout the surrounding fluid. |
| Colloidal dispersions, AKA colloids | The middle ground between suspensions and solutions, in which a dispersed (solute-like) substance is distributed throughout a dispersing (solvent-like) substance. E.g. proteins, synthetic polymers, etc. |
| The size of colloidal particles | Larger than simple molecules but small enough to remain distributed and not settle out. They range in diameter from 1 to 1000 nm |
| Surface area of colloids | They have an enormous surface area which allows many more interactions to exert a great total adhesive force, which attracts other particles and leads to practical uses of colloids. |
| Foam | Colloid: gas dispersed in a liquid |
| Tyndall effect | The light scattering phenomenon observed when light passes through colloids: the light is visibly broader than one passing through a solution, e.g. sunlight passing through dust. |
| Brownian motion | A characteristic movement in which the particles change speed and direction erratically. This motion results because the colloidal particles are being pushed this way and that by molecules of the dispersing medium. |
| Why is Brownian motion significant | These collision with the dispersing medium is primarily responsible for keeping colloidal particles from settling out. |
| When colloidal particles collide why don’t they aggregate into larger particles? | Many molecules arrange themselves so that the nonpolar portions are on the interior and the charged polar portions on the exterior |
| Strategies to eliminate colloids | Heating the air increases collisions. Adding electrolyte solutions nullifies the charged exteriors allowing them to coagulate and settle. |
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