CHM104 Dr Friesen Exam 1 Study Guide Lectures 10 18 Lecture 10 September 22 Textbook section 14 4 14 6 Kc vs Kp o Equilibrium equation can be written using molarity or partial pressure gases This can change k o Pressure of gas is proportional to molarity PV nRT P n V RT o Heterogeneous Equilibria Pure solids and liquids are replaced by 1 s Their concentrations DO NOT depend on their container volume just the density of the substance Example CaCO3 s CaO s CO2 g K 1 CO2 1 K CO2 EQ Lecture 11 September 24 Textbook Section 14 7 14 8 c aA bB cC dD d C D K a b A B Reaction quotient Qc as the ratio at any point in the reaction of the concentrations of the products raised to their stoichiometric coefficients divided by the concentrations of the reactants raised to their stoichiometric coefficents o Uses the partial pressure atm in place of concentration and is called Q p Difference between reaction quotient and equilibrium constant o At given temp the equilibrium constant has only one value and it specifies the relative amounts of reactants and products at equilibrium Summarizing direction of change predictions o Reaction quotient Q relative to the equilibrium constant k is a measure of progress of the reaction toward equilibrium Q K reaction goes to the right toward product Q K reaction goes to the left toward reactants Q K reaction is at equilibrium Lecture 12 September 29 Textbook Section 14 9 Heterogeneous Equilibrium Calculations o Guidelines Disturbance Response EQ Shifts Add Product Use Product Remove Product Generate More Exothermic Reaction reaction shifts to the left Endothermic Reaction reaction shifts to the right Effect of increasing the pressure of a system of gasses at equilibrium o Decrease volume o The system will try to lower the pressure o Pressure is proportional to the moles of gas Le Chatelier Principle o If an outside influence disrupts a chemical reaction that is at equilibrium the reaction will respond by shifting in a direction to partially counteract the disturbance and restore equilibrium Lecture 13 October 1 Textbook Section 15 1 15 4 Acids o o o o Tastes sour Can dissolve metals Neutralized by bases Common Acids HCl Hydrochloric Acid HF Hydrofluoric Acid HNO3 Nitric Acid H2SO4 Sulfuric Acid HC2H3O2 Acetic Acid Carboxylic Acid Group Citric Acid H3C6H5O3 Malic Acid CH2C4H6O5 Bases o Tastes bitter o Feels slippery o Common Bases NaOH Sodium Hydroxide NaHCO3 Sodium Bicarbonate Na2CO3 Sodium Carbonate NH3 Ammonia Amine CH3NH2 methylamine Definitions of Acids and bases o Arrhenius Acid substance that produces the H ions in aqueous solutions Base substance that produces OH ions in aqueous solutions Acids and bases react to form water o Bronstad Acid Substance that DONATES H Base Substance that ACCEPTS H H2O can either accept H act as base OR donates H acts as acid Amphoteric can act as an acid or a base o Lewis Acid ACCEPTS electron pair Base DONATES electron pair Lecture 14 October 3 Section 15 4 15 6 Strong Acid o Completely ionizes in reaction with H2O o Table of strong acids given on exam Weak Acid o Only partially ionized in reaction with H2O o Strength depends on how strong the H is bound to the acid Stronger the bond weaker the acid o Common weak acids HF HC2H3O2 H2SO3 H2CO3 H3PO4 The Ka Equilibria A aq HA aq H 2 O l H 3 O aq HA EQ O H3 EQ A K a EQ A base H O EQ H A K b EQ EQ Typical Ka s for weak acids range from 10 2 to 10 10 Auto ionization of water o Water can act as an acid OR base EQ 1 0 10 14 O H EQ H 3 O K w In pure water H 3 O O H K w K a K b pH Scale o We use the pH scale to quantify acidity H3 O pH log o pH is always positive due to H3O less than 1 therefore negative log will be positive Lecture 15 October 6 Textbook Sections 15 6 15 9 When calculating pH o Ka or Kb 1 0 x 10 4 for x assumed to be small rule HA EQ value less than 5 A 100 be valid ionization EQ Lectures 16 18 Calculation worksheets from class and discussion