Concentrations of Solutions, Types of Solutions, Solubility, Reaction Kinetics
Concentrations of Solutions
- Solutions can be described as concentrated or dilute.
- Concentration: Measure of the amount of solute in a given amount of solvent or solution.
- Expressed as molarity (M) or molality (m)
- Molarity (M):
- M = \frac{moles \ of \ solute}{liters \ of \ solution}
- Molality (m):
- m = \frac{moles \ of \ solute}{kilograms \ of \ solvent}
Practice Time!
Find the molarity of a solution of 58.0 \ g of AgNO_3 dissolved in 1.00 \ L of solution.
- Molar mass of AgNO_3 = 169.87 \ g/mol
- Moles of AgNO_3 = \frac{58.0 \ g}{169.87 \ g/mol} = 0.341 \ mol
- Molarity = \frac{0.341 \ mol}{1.00 \ L} = 0.341 \ M
How many grams of KCl should be used to prepare 4.00 \ L of 0.500 \ M solution?
- Moles of KCl = 0.500 \ M \times 4.00 \ L = 2.00 \ mol
- Molar mass of KCl = 74.55 \ g/mol
- Grams of KCl = 2.00 \ mol \times 74.55 \ g/mol = 149.1 \ g
To what volume should 10.0 \ g of NH_4Cl be diluted to prepare a 0.35 \ M solution?
- Moles of NH_4Cl = \frac{10.0 \ g}{53.49 \ g/mol} = 0.187 \ mol
- Volume = \frac{0.187 \ mol}{0.35 \ M} = 0.534 \ L
Dilutions
Dilution: the process of reducing the concentration of a solute in a solution, usually simply by mixing with more solvent.
- M1V1 = M2V2
- M_1 = initial molarity.
- V_1 = initial volume.
- M_2 = final molarity.
- V_2 = final volume.
- M1V1 = M2V2
Practice Time!
How much 0.05 \ M \ HCl solution can be made by diluting 1 \ L of 2.00 \ M \ HCl?
- M1V1 = M2V2
- (2.00 \ M)(1 \ L) = (0.05 \ M)(V_2)
- V_2 = \frac{(2.00 \ M)(1 \ L)}{0.05 \ M} = 40 \ L
How much concentrated 12 \ M nitric acid is needed to prepare 250 \ mL of a 6.0 \ M solution?
- M1V1 = M2V2
- (12 \ M)(V_1) = (6.0 \ M)(250 \ mL)
- V_1 = \frac{(6.0 \ M)(250 \ mL)}{12 \ M} = 125 \ mL
How much water would need to be added to 500 \ mL of a 2.5 \ M \ NaCl solution to make a 1.0 \ M solution?
- M1V1 = M2V2
- (2.5 \ M)(500 \ mL) = (1.0 \ M)(V_2)
- V_2 = \frac{(2.5 \ M)(500 \ mL)}{1.0 \ M} = 1250 \ mL
- Water added = 1250 \ mL - 500 \ mL = 750 \ mL
Practice Time!
You are given 15 \ g of NH_4Cl. Your boss asks you for 0.5 \ L of a 0.10 \ M solution. How would you make this solution?
- Molar mass of NH_4Cl = 53.49 \ g/mol
- Moles needed = 0.10 \ M \times 0.5 \ L = 0.05 \ mol
- Grams needed = 0.05 \ mol \times 53.49 \ g/mol = 2.67 \ g
- Dissolve 2.67 \ g of NH_4Cl in enough water to make 0.5 \ L of solution.
Types of Solutions
HOMOGENEOUS MIXTURE
- UNIFORM DISTRIBUTION OF PARTICLES
HETEROGENEOUS MIXTURE
- NON-UNIFORM DISTRIBUTION OF PARTICLES
Heterogeneous: uneven distribution throughout
- Suspensions: particles are so large that they will settle out unless the mixture is constantly stirred (Ex. Muddy water or oil in water)
Colloids: heterogeneous mixtures with particles that are intermediate in size between solutions and suspensions
- Tyndall Effect: colloids scatter light, making the mixture appear turbid.
Solubility
Solubility: a measure of how well one substance dissolves in another.
- The amount of a substance required to form a saturated solution with a specific amount of solvent at a specific temperature.
- "Like often dissolves like" is useful for predicting solubility.
- Polar molecules dissolve polar and nonpolar molecules dissolve nonpolar.
- Soluble: capable of being dissolved
General Solubility Guidelines:
CO3^{2-} and PO4^{3-}
- Soluble only when combined with Group 1 ions or NH_4^{1+}
OH^{1-} and S^{2-}
- Soluble only when combined with Group 1 ions, Ca^{2+}, Ba^{2+}, Sr^{2+}, or NH_4^{1+}
Water: The Universal Solvent
Water's polarity and other properties of water make it an incredibly versatile solvent.
Liquid at room temperature
High melting and boiling points compared to its size
- Due to the strength of its hydrogen bonds
High specific heat
- It has to absorb a LOT of energy, in the form of heat, to increase its temperature.
Dissolving
Dissolution: the process of dissolving in forming a solution.
- The solute separates into ions or molecules, and each one is surrounded by molecules of solvent.
Dissociation: the separation of ions that occurs when an ionic compound dissolves
Solvation: the interactions between the solute and solvent particles when the solute particle is surrounded by molecules of a solvent.
Dissolving Writing Net Ionic Equations
A net ionic equation is often written for reactions of ions in aqueous solutions instead of a chemical formula equation.
- It shows only the compounds and ions that undergo a chemical change in the reaction.
- All other ions (spectator ions) are canceled out on both sides.
Electrolyte is a substance whose aqueous solution conducts an electric current.
Nonelectrolyte is a substance whose aqueous solution does not conduct an electric current.
Factors that Affect Solubility
Surface area of solute
- The smaller the solute, the greater its surface area, and thus the more places for contact.
Agitating the solution
- The more the mixture is stirred, the more the fresh solvent is brought into contact with undissolved solute.
Heating a solvent
- Raises the average kinetic energy of the solvent molecules
- Solvent molecules collide more often with solute
- Spreads the molecules apart, allowing more solute to enter between them.
Pressure changes (when gases dissolve in liquids)
- The higher the pressure, the more soluble the gas.
- Henry's Law: The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of the liquid.
Types of Solutions Based on saturation
- Saturated: a solution containing the max amount of dissolved solute at a given temperature.
- Supersaturated: a solution containing more dissolved solute than a saturated solution in the same conditions; above its max capacity.
- Unsaturated: a solution containing less dissolved solute than a saturated solution in the same conditions; below its max capacity.
Properties of Solutions
The presence of solutes affects the properties of solutions.
Colligative properties: properties dependent on the concentration of solute particles but not on their identity.
Boiling and freezing points of solutions are different from pure solvents.
Ex. The boiling point and freezing point of salt water is different from pure water.
A nonvolatile solute added to a solvent will lower the freezing point and raise the boiling point.
Nonvolatile solute: a substance that has little tendency to become a gas under its existing conditions.
- Equilibrium vapor pressure
Vapor pressure: the pressure caused by molecules in the gas phase that are in equilibrium with the liquid phase.
- As concentration increases, the number of solutes increases in a given volume, and the proportion of solvent (like water) decreases.
- Therefore, less water molecules can escape the liquid phase less vapor decreased vapor pressure.
Freezing-Point Depression
The difference between the freezing points of the pure solvent and a solution from that solvent.
It is directly proportional to the molal concentration of the solution.
- The greater the solution's molality, the greater the difference.
A solution always has a lower freezing point than its pure solvent.
- Ex. If you add salt to water, it will lower the freezing point.
Boiling-Point Elevation
The difference between the boiling points of the pure solvent and a solution from that solvent.
It is directly proportional to the molal concentration of the solution.
- The greater the solution's molality, the greater the difference.
A solution always has a higher boiling point than its pure solvent.
- Ex. If you add salt to water, it will raise the boiling point.
Osmotic Pressure
Osmosis: Solvent will move through a semipermeable membrane from the side of higher solvent and lower solute (and thus lower concentration) to the side of lower solvent and higher solute (and thus higher concentration).
Osmotic pressure: the external pressure that must be applied to stop osmosis.
- The greater the concentration, the greater the osmotic pressure.
Reaction Kinetics
Reaction rate: measured by the change in conc. of reactants or products per unit of time.
Factors that affect reaction rate:
Concentration
- The higher the conc., the more particles in a given space, the more likely they are to collide and thus react, increasing the rate.
Surface area
- The greater the surface area, the more space for particles to come into contact with each other and potentially react, increasing the rate.
Temperature
The higher the temperature, the faster the molecules are moving, and the more likely they are to collide and thus react, increasing the rate. Adding a catalyst also increases the reaction rate
- Catalyst: lowers the amount of activation energy needed, it speeds up the reaction w/out being permanently changed itself.
Chemical Equilibrium
Reactions can be reversible.
- 2SO2(g) + O2(g) \rightleftharpoons 2SO_3(g)
When the forward and reverse processes are happening at the same rate, we say that the system is in equilibrium.
Indicated by a double arrow.
Chemical Equilibrium: A dynamic process when there is no net change occurring in the amount of reactants and products in the system, thus no visible change.
When products and reactants are forming at the same rate.
When reversible processes are occurring at the same rates.
- Consider a rxn at equilibrium: A + B \rightleftharpoons AB
- AB is happening at the same rate as
- A + B
- AB \rightleftharpoons A + B
Le Chatelier's Principle
Le Chatelier's Principle: If a stress is applied to a system in dynamic equilibrium, the system changes in a way that relieves the stress.
- Stress: any kind of change to the system at equilibrium.
Ex. Blood pH and exercise
- Stress = Exercise adds CO_2
- H2CO3 (aq) \rightleftharpoons CO2 (aq) + H2O (l)
- Shift = the rxn shifts to the reactants
- H2CO3 starts to build up in the bloodstream, increasing the acidity of the blood.
- In response, as you exercise, the body starts to breath rapidly, removing CO_2 from the bloodstream.
- Stress = Breathing faster removes CO_2
- H2CO3 (aq) \rightleftharpoons CO2 (aq) + H2O (l)
- Shift = the rxn shifts to the products.
Factors that Affect Equilibrium
Temperature change
Exothermic = rxns that, overall, release heat
- Example:
- N2(g) + 3H2(g) \rightleftharpoons 2NH_3(g) + 91 \ kj
- Adding heat would favor the reverse rxn
Endothermic = rxns that, overall, absorb heat
- Example:
- N2O4(g) + heat \rightleftharpoons 2NO_2(g)
- Adding heat would favor the forward rxn
Pressure change (in gases)
- If pressure is increased, the rxn will favor whichever direction produces fewer gas molecules
- If pressure is decreased, the rxn will favor the direction that produces more gas molecules
Examples:
- CaCO3(s) \rightleftharpoons CaO(s) + CO2(g)
- Unaffected, because there aren't gases on both sides.
- N2(g) + 3H2(g) \rightleftharpoons 2NH_3(g)
- Increase in pressure favors forward, decrease in pressure favors reverse
- H2(g) + Cl2(g) \rightleftharpoons 2HCl(g)
- Unaffected, because there are equal moles* of gas on each side
- CaCO3(s) \rightleftharpoons CaO(s) + CO2(g)