Chemistry Exam Notes PDF

General Equilibrium Reversible Reactions: ​ Many chemical reactions are reversible and can proceed in both the forward and reverse directions. ​ Denoted using a double headed arrow (⇌) ​ Ex.N2O4(g)⇌ 2 NO2(g) ​ The forward and reverse reactions are constantly going but there is a p...

General Equilibrium Reversible Reactions: ​ Many chemical reactions are reversible and can proceed in both the forward and reverse directions. ​ Denoted using a double headed arrow (⇌) ​ Ex.N2O4(g)⇌ 2 NO2(g) ​ The forward and reverse reactions are constantly going but there is a point when there is no net change in the amounts of reactants and products ​ When a chemical system reaches a point where there is no net observable change in the reactants and products, the system is said to be at EQUILIBRIUM. Chemical Equilibrium: ​ The forward and reverse reactions do not stop, but the rate of the forward reaction equals the rate of the reverse reaction ​ Since the 2 processes are still occurring, chemical equilibrium is considered to be DYNAMIC. ​ At equilibrium, the rates of the 2 processes are the same but the amounts of reactants and products are not necessarily equal ​ Equilibrium can only be reached by CLOSED systems. ​ Equilibrium can be approached from either the reactant or product side ​ In most reactions, either the forward or reverse process is favoured ​ At a given temperature, the point at which a system reaches equilibrium is constant. Representing Equilibrium Graphically: ​ It is possible to determine the point at which equilibrium is reached by constructing a concentration vs time graph for a reversible reaction ​ Since equilibrium is reached when there is no net change in the concentrations of reactants and products, the slopes of all lines on the graph will be zero at equilibrium. Quantitative Look at Equilibrium: ​ For a given temperature, the concentrations of products and reactants at equilibrium are represented by the equilibrium constant (Keq) ​ Keq is determined by comparing the concentrations of the products to the concentrations of the reactants at equilibrium in the equilibrium expression ​ The equilibrium expression is written as the concentration of the products raised to their molar coefficients divided by the concentration of reactants raised to their molar coefficients. ​ Solids and liquids are never included in the equilibrium expression. (only gases and aqueous solutions!) ​ Once the rates of the forward and reverse reactions have stabilized, a system can continue at equilibrium indefinitely ​ That is, until something comes along to disturb it Le Châtelier’s Principle: ​ When a stress is applied to a system at equilibrium, the equilibrium will shift(increase in speed) in the direction necessary to relieve that stress. ​ A stress can be a change in concentration, volume, pressure, temperature or addition of an inert species. Question: At 100°C the reaction below has an equilibrium constant, Keq, value of 2.2x10-10. If 1.00 mol of phosgene, COCl2, is placed in a 10.0 L flask, calculate the concentration of carbon monoxide at equilibrium. The Approximation Rule for ICE Problems: ​ In some cases, the change in concentration is so small it is considered negligible. ​ To determine this, divide the smallest non-zero initial concentration by the Keq value. ​ If this is greater than 1000, any “x” being added to or subtracted from a value can be ignored ​ If not, keep all “x’s” and use the quadratic formula to solve for “x” (only 1 “x” value will be valid).  Solubility Equilibrium ​ Solution – homogeneous mixture of solvent and one or more solutes ​ Solvent – substance that has other substances dissolved in it ​ Solute – substance that is dissolved in solution ​ Solutions can have variable composition – different ratios of solvent to solute are possible ​ Ex. Weak (dilute) or strong (concentrated) solutions depend on how much solute is present ​ Aqueous solution – water is the solvent TYPES OF SOLUTES: ​ Ionic compound – compound formed by ionic bond (b/w + & - ions) ​ Polar molecule – molecule that has an uneven distribution of charge; one end has a partial -/ + charge ​ Non-polar molecule – covalently bonded molecule that does not possess a dipole moment; because of the arrangement of its bonds between atoms ​ Solubility – amount of solute that dissolves in a given quantity of solvent at a specific temperature. Ex. The solubility of NaCl in water at 20oC is 36 g/ 100 mL ​ Saturated solution – solution in which no more of a particular solute can be dissolved at a specific temperature ​ Unsaturated solution – solution in which more of a particular solute can be dissolved at a specific temperature ​ Soluble – solubility is greater than 1g / 100 mL of solvent ​ Insoluble – solubility is less than 0.1 g / 100 mL solvent ​ Slightly soluble – solubility range between 0.1 g – 1 g Remeber: LIKE DISSOLVES LIKE ​ Rate of dissolving – speed at which solute dissolves in solvent ​ Solubility – the maximum amount of solute that can dissolve in an amount of solvent FACTORS THAT AFFECT THE RATE OF DISSOLVING (TIME IT TAKES FOR SOMETHING TO DISSOLVE: 1.​ Temperature: rate of movement (energy), therefore # of collisions 2.​ Agitation:contact between solute and solvent 3.​ Particle size: size will change surface area FACTORS THAT AFFECT SOLUBILITY (AMOUNT OF SOMETHING THAT CAN DISSOLVE: 1.​ Molecule size: bigger size, lower solubility2 2.​.Temperature -​ Solid – temp, energy for breaking bond, therefore higher solubility -​ Gas – temp, energy of gas molecules, therefore more collisions and higher solubility 3.​ Pressure: pressure on gas, higher solubility  ​ DISSOCIATION: Separation of ions that are already present in a compound ​ IONIZATION: Formation of positively or negatively charged ions from molecules that do not initially contain ions EQUILIBRIUM AND SOLUBILITY: ​ All of these equilibria equations are dissolving equations. i.e. 1 solid splitting into its ions ​ Keq changes to Ksp, where “sp” stands for solubility product ​ Ksp indirectly represents the amount of solute that can dissolve ​ Ksp also represents the maximum concentration of the ions at equilibrium (i.e. saturation) RELATING KSP TO SOLUBILITY OF A SALT: ​ Ksp is dimensionless (ie we do not give it a unit) but is related mol/L ​ Solubility is usually measured in g/100 mL or g/L ​ The “x” (ie the change in concentration) value from the Ksp expression represents the solubility of the salt in mol/L (i.e. molar solubility). PREDICTING PRECIPITATION: ​ In order to determine if a precipitate will form or not (or if the total amount of a solute will dissolve) you must determine the TRIAL ION PRODUCT (Q) ​ Q is found by substituting the “ACTUAL ION CONCENTRATIONS” into the Ksp expression and comparing that value with the Ksp ​ If Q is > Ksp a precipitate will form (or not all will dissolve) ​ If Q is ≤ Ksp a precipitate will not form (or all will dissolve). ​ Selective precipitation is a situation where you have more than one type of cation or anion in solution and you want to precipitate one out and leave the other dissolved, or you want to make precipitations that contain only one of the ions at a time ​ Hint: to answer these questions, consider the relative solubility of the salts you could make – you want to add ions that will be very soluble with one of the ions but only slightly soluble with the other cation. If you are given the Ksp, the one that gives you the lowest concentration will precipitate first. Acid Base Equilibrium STRONG VS WEAK ACID: ​ Strong Acid: ionizes 100 % in water to form hydronium ions and the anion of the acid. ​ Examples are hydrochloric, nitric, sulfuric ​ Weak Acid: ionizes less than 50% in water to form hydronium ions and the anion of the acids. ​ Examples are acetic, formic, nitrous ​ Weak acids reach equilibrium with their anion and hydronium and therefore do not ionize any more. ​ The equilibrium constant for weak acids is called the Ka STRONG VS WEAK BASE: ​ Strong Base: metal hydroxides (NaOH, Ca(OH)2) that dissociate completely in water to form hydroxide and metal ions ​ Examples are sodium hydroxide, calcium hydroxide ​ Weak Base: substances that react with water to form hydroxide ions and a cation. ​ Examples are ammonia, carbonates Weak bases reach equilibrium with their cation and hydroxide and therefore do not ionize any more. ​ The equilibrium constant for a weak base is called the Kb PERCENT IONIZATION: ​ The percent ionization indicates the degree to which a weak acid or base ionizes in water ​ It is determined using the following formula: -​ % ionization = [H3O+]eq/ [HA]initial x 100%  ​ For any acid-base equilibrium, the Ka x Kb for any conjugate acid-base pair must equal Kw. AMPHOTERIC (AMPHIPROTIC) SUBSTANCES: ​ Substances that can act as an acid in one reaction but as a base in another ​.Eg. HCO3- + H2O H⇌2CO3 + OH- Or HCO3- + H2O CO⇌32- + H3O+ ​ Compare the Ks in order to see which one is most likely to occur POLYPROTIC ACIDS: ​ Polyprotic acids have more than 1 acidic hydrogen. ​ Each ionization will have its own Ka value and in general, Ka1> Ka2 > Ka3 etc. ​ Since the first ionization is the largest, only the first ionization is used to determine the pH. ACID-BASE PROPERTIES OF SALT SOLUTIONS: ​ Salts are solids at room temperature composed of cations and anions arranged in a crystal lattice (i.e. ionic compounds) ​ When salts dissolve in water, they dissociate into aqueous solutions of ions that may or may not affect the pH of a solution. ​ Salts that Form Neutral Solutions: -​ Salts with the cation of a strong base and the anion of a strong acid. (eg NaCl, K2SO4) ​ Salts that form Acidic Solutions: -​ Salts where the cation is the conjugate acid of a weak base and the anion of a strong acid. (eg NH4Cl) -​ Salts where the cation is a highly charged metal ion and the anion is from a strong acid. (eg AlCl3) -​ Other compounds that make Acidic Solutions: Non-metal oxides (eg CO2) ​ Salts that Form Basic Solutions: -​ Salts with the cation of strong base and the anion of a weak acid. (eg NaC2H3O2) -​ Other compounds that make Basic Solutions: Metal oxides (eg CuO) ACID-BASE TITRATION: ​ An indicator can be used to determine the end point. At the end point, the indicator will change colour. ​ The equivalence point or stoichiometric point occurs when equal moles of H3O+ and OH- ionshave reacted to produce water. ​ At the equivalence point, neutralization is complete and you are left with salt and water. Since salts can be acidic, basic or neutral, it does not always occur at a pH of 7. ​ When choosing an indicator for a titration, one tries to choose an indicator that will change when the equivalence point is reached. Strong Acid and Base Titration: ​ These will produce a neutral salt.Equivalence point will occur at a pH of 7. ​ Calculation problems are done as a simple limiting factor problem since there is complete ionization/dissociation. Weak Acid and Strong Base Titration: ​ These produce a basic salt ​ Equivalence point will occur at a pH above 7 as the conjugate base of the weak acid determines the pH of the solution ​ These problems are done using an ICE table. Strong Acid and Weak Base Titration: ​ These produce an acidic salt ​ Equivalence point will occur at a pH below 7 as the conjugate acid from the weak base determines pH of the solution ​ These problems are done using an ICE table.