Unit 5: Solutions, Kinetics and Equilibrium

Solutions Parts of a Solution -- any homogeneous mixture Not to be confused with: a. suspensions: small particles of one substance suspended in another b. colloids: a suspension that doesn’t settle Parts of a Solution 1. Solute : the material that gets dissolved makes up less than 50% 2. Solvent : material that dissolves the solute makes up more than 50% Often the solvent is a liquid --but not always

Solution Concentration Qualitative: lots of solute per unit solvent = concentrated little solute per unit solvent = dilute (how much Mio you add to your water- a little or a lot) Quantitative: This is an actual measured quantity Molarity (M) moles per liter solute solution moles M = Liters (Table T)

Example: 100 g of MgO are dissolved in enough water to make 5.0 L of solution. What is the molarity? 100 Solution: first, find the moles of solute moles = = 2.5 moles 40 2.5 next, use molarity formula M = = 0.5M 5.0 Example: How many grams of BaCl2 are needed to make 300.0 mL of a 0.70M solution? moles Solution: find the moles of solute required = 0.21 moles 0.7 = 0.3 grams now, change moles to grams = 0.21 = 44 g 207

) 2. Percent (%) (Table T) % = ( amount of solute . x 100 % = ( amount of solute . ) x 100 amount of total solution amounts can be volumes or masses, but units must agree -- except it’s OK to mix g and mL because 1mL has a mass of 1 g Example: 100 mL alcohol is diluted with enough water to form a final volume of 250 mL. % alcohol = ? 100 Solution: x 100 = 40% 250 Example: 100 mL alcohol is mixed with 250 mL water. % alcohol=? 100 29% Solution: x 100 = 350

( ) ( ) ( ) 3. Parts per Million (ppm) (Table T) amount of solute . x 1,000,000 amount of total solution Example: a 50.0 gram sample of apple contains 3.2 x 10-4 g of pesticide. ppm = ? ( 3.2 x 10-4 ) ( ) Solution: 1,000,000 = 6.4 ppm 50.0 When solutes dissolve in liquids, they change some physical properties of the liquid, such as: 1. The freezing point is lowered -- salt used to melt ice on roads -- salt used to freeze home-made ice cream -- antifreeze used to prevent car freeze-ups

2. The boiling point is raised --antifreeze also prevents boiling These changes depend on the concentration of solute particles, not what kind Some substances cause larger-than-expected changes -- ionic compounds (salts) -- acids (H+) and bases (OH-) These substances dissociate, or ionize, when dissolved in water (separate into + and – ions) Example: NaCl (s)  Na+(aq) + Cl-(aq) 1 mole 1 mole 1 mole 2 moles So, 1M NaCl affects the fp and bp as if it were a 2M solution Examples: CaCl2(s)  Ca+2(aq) + 2Cl-(aq) Na2CO3(s)  2Na+(aq) + CO3-2(aq)

These ions behave independently in solution Example: Na2CO3 + CaCl2  2NaCl + CaCO3 2Na+(aq) + CO3-2(aq) + Ca+2(aq) + 2Cl-(aq)  2Na+(aq) + 2Cl-(aq) + CaCO3(s) This is an Ionic Equation The solid that forms from a solution (CaCO3) is a precipitate The Na+ and Cl- do not participate in the reaction -- they are spectator ions The Net Ionic Equation is: Ca+2(aq) + CO3-2(aq)  CaCO3(s)

Reactions between ions in solution will occur only if one of the products is insoluble (gas or solid), or water Example: NaCl + KNO3  ? Solution: the possible products, NaNO3 and KCl, are both soluble so, no reaction Use reference table F to predict solubilities Pb+2 + 2I- = PbI2

Kinetics All substances contain chemical potential energy A  B high PE low PE Energy released -- Exothermic low PE high PE Energy absorbed -- Endothermic Most chemical reactions require some energy to get started -- activation energy

Potential Energy Diagrams A – PE of reactants 50 40 30 20 10 PE B – Activation energy B kJ C – PE of products D – Heat of reaction; ΔH D A C Progress of Reaction In this diagram, PE decreased energy released; exothermic ΔH is negative

Exothermic reactions have a negative ΔH Endothermic reactions have a positive ΔH See reference Table I Notice that ALL values of ΔH have a sign (+/-)!

How do we tell if a reaction will occur spontaneously? Consider 2 factors: 1. Heat of Reaction (ΔH) --low energy conditions are more stable -- exothermic reactions (- ΔH ) are favored 2. Change in entropy (ΔS) --entropy is a measure of disorder, or randomness gas liquid solid disordered neat, orderly high entropy low entropy --increasing entropy (+ ΔS ) is favored

Favorable conditions are: Exothermic -ΔH Increasing entropy +ΔS Any reaction with -ΔH, +ΔS is always spontaneous +ΔH, -ΔS is never spontaneous Reactions where only one factor is favorable are spontaneous under some conditions – depends on temperature. Example: 2C(s) + 3H2(g) → C2H6(g) ΔH = -84.0 kJ Exothermic, favorable (Table I) ΔS solid to gas, entropy increases, favorable So this reaction is spontaneous

Example : H2O (solid)  H2O (liquid) Types of Reactions heat of reaction(ΔH) [favorable?] entropy (ΔS) [favorable?] 1. decrease ( - ) [yes] increase ( + ) [yes] YES [no] decrease ( - ) 2. increase ( + ) [no] NO 3. decrease ( - ) decrease ( - ) Only at low temp [yes] [no] 4. increase ( + ) [no] increase ( + ) [yes] Only at high temp The result in cases 3 and 4 depends on temperature Example : H2O (solid)  H2O (liquid) -- solid to liquid, energy increases, energy absorbed ΔH positive, unfavorable -- solid to liquid, entropy increases ΔS positive, favorable So, ice melts spontaneously only at high temperatures

Reaction Rates Collision Theory --reactions occur when molecules physically collide -- the collision must be effective – it must have 1. enough energy 2. proper orientation -- any change that increases the number of effective collisions will increase the reaction rate

Factors affecting reaction rate 1. Temperature -- faster motion of molecules means more frequent and more effective collisions 2. Surface Area -- more contact between reactants means more frequent collisions (increase surface area by dividing a solid into many small pieces) Concentration of Reactants – higher concentration = more frequent collisions in solutions – raise concentration by dissolving more reactant, or by reducing water in gases – raise concentration by raising pressure

4. Nature of Reactants -- --if bonds must be broken first, reaction is slow --if not, reaction is fast ex., reactions between solutions Catalysts– a catalyst is a substance that increases the rate of a chemical reaction catalysts are not used up during the reaction catalysts do not affect the ΔH or the spontaneity of the reaction catalysts work by lowering the activation energy

ineffective no catalyst low energy collision catalyst effective

Equilibrium Kinds of Equilibrium -- a state of balance between opposing forces -- no overall changes occur at equilibrium -- in a dynamic equilibrium, opposing changes occur at equal rates and cancel each other Kinds of Equilibrium 1. Phase equilibrium: ex.) ice and water at 0ºC ex.) water and water vapor in a closed jar 2. Solution equilibrium 3. Chemical equilibrium

Solution Equilibrium When solute is added to solvent, it dissolves until an equilibrium is reached: solid dissolved The rate of dissolving = the rate of precipitation when the solution is saturated. A saturated solution contains the maximum amount of solute that can be dissolved under specific conditions, like temperature. By comparison, an unsaturated solution contains less, and a supersaturated solution contains more. (Table G) The solubility of a substance is its concentration at equilibrium (at saturation with given conditions) at STP

Factors affecting solution equilibrium: A. solid/liquid solutions -- Temperature almost all solids are more soluble at high temps

B. gas/liquid solutions -- Temperature gases are less soluble at high temperatures See reference Table G The lines that go DOWN are gases -- Pressure gases are more soluble at high pressures. Think of soda. What happens when you open it?

Chemical Equilibrium A  B Many chemical reactions are reversible forward reaction A  B reverse reaction B  A Both reactions use the same PE diagram forward: ΔH +; endo act. E act. E B reverse: ΔH ; exo ΔH ΔH ΔH’s are equal but opposite A

A  B Equilibrium occurs when: 1. the rate of the forward and reverse reactions are equal 2. the concentration of the product and reactant remain constant 3. no visible changes are occurring

Equilibrium Shifts LeChatelier’s principle: When a system at equilibrium is subject to a change, the equilibrium will shift to undo the change Three factors can affect chemical equilibrium: 1. Concentration of product or reactant Example: A + B  C + D is at eq.; then more D is added Result: the equil. will shift to reduce the amount of D more A and B will form, and C will be used up to combine with the excess D. the equilibrium shifts left Example: AgCl(s)  Ag+(aq) + Cl-(aq) at eq.; then NaCl is added Result: adding NaCl would add more Cl-; therefore, the equil. shifts left, and AgCl(s) will precipitate

2. Pressure: affects equil. involving gases If pressure increases, the equil. shifts to the side with fewer moles of gas Example: N2(g) + 3H2(g)  2NH3(g) at eq., then pressure 1 mole + 3 moles  2 moles Result: to cause the pressure to go back down, the eq. shifts right Example: CaCO3(s) + HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g) What happens to the amount of CaCl2 if pressure decreases? Result: 0 moles  1 mole Eq. shifts right to restore pressure, producing more CaCl2 If gases are equal on both sides, equil. does not shift

3. Temperature: consider heat as a product or reactant Example: exothermic reaction NaOH + HCl  NaCl + H2O + heat at eq., then temp Result: Eq. shifts left to get rid of excess heat Example: endothermic reaction N2 + O2 + heat  2NO at eq., then temp Result: Eq. shifts right Catalysts do not shift the equilibrium -- they speed up the forward and reverse reactions equally