Chem 1062 Final Study Guide PDF

Chapter 5: The temperature at which a phase change occurs depends upon the molecular structure of the compound. For molecular substances, as a substance changes, phase intermolecular forces are overcome (not chemical bonds). The direction of change is d...

Chapter 5: The temperature at which a phase change occurs depends upon the molecular structure of the compound. For molecular substances, as a substance changes, phase intermolecular forces are overcome (not chemical bonds). The direction of change is determined by an increase in the total entropy change (ΔStotal) or the Gibbs energy change (ΔG). Measuring temperature changes can be related to molecular level changes in interaction strength by the thermodynamic function ΔH and bond energies. Explain the difference/relationship between temperature, thermal energy, and kinetic energy. Explain why particles in gases move with a range of different velocities at a given temperature. Identify the Boltzmann distributions of particles at different temperatures or for particles of different molecular weights at a specified temperature. Explain how temperature and kinetic energy are related, including the energy associated with vibration, rotation, and translation in different phases. Explain the causes of water’s anomalous properties (high melting point, boiling point, lower density of ice relative to liquid water, specific heat). Explain why the heat capacity of a substance is affected by the molecular-level structure. Draw heating or cooling curves showing how the temperature changes when thermal energy is added to a substance (including a phase change). Explain why the temperature changes except during the phase change. Define and give examples of open, closed, and isolated systems. Explain the difference between state and path functions, and give examples. For exothermic and endothermic processes, identify the direction of thermal energy change and the sign of q or ΔH. For phase changes, identify the direction of thermal energy change and the sign of q or ΔH. Explain the role of probability in entropy changes. Predict the sign of the entropy change for simple systems. Explain the second law of thermodynamics in terms of the system and surroundings Explain why we usually use ΔG instead of the total entropy change to predict whether a process is thermodynamically favorable. Classify a chemical reaction as endothermic or exothermic Interpret an enthalpy diagram Sketch heat flow Understand the meaning of specific heat capacity  Calculate q Manipulate qlost + qgained = 0 to solve for m, c, T, and/or q Distinguish between qrxn and H Compare calculated H values to tabulated values Calculate Ccal Calculate the final T of a system Determine H from calorimeter data Apply stoichiometry to thermochemical reactions Write and apply the energy-to-mole ratios in a thermochemical equation Write the heat of formation reaction for a given compound Diagram a heat of formation reaction Diagram a heat of reaction for a thermochemical equation Interpret a heat of reaction enthalpy diagram Calculate Hrxn using heats of formation Describe standard state conditions State the first law of thermodynamics State the second law of thermodynamics Describe microstates, energy dispersal, and entropy Define standard molar entropy and the conditions that apply Predict relative entropy changes of a system Compare relative entropies of two or more substances Predict the sign of Sºrxn Calculate Sºrxn from Sº Compare S of system and surroundings to determine spontaneity Recognize Sºrxn is zero for a system at equilibrium and explain why Calculate the temperature at which a process becomes spontaneous Define spontaneous change Describe how heat leaving a system impacts the entropy of the surroundings Define standard free energy and its relationship to ΔH and ΔS Calculate ΔG°rxn from standard free energy values (ΔG°f)  Calculate ΔG°rxn from standard enthalpy and entropy values (ΔH°rxn and ΔS°rxn) Understand how T impacts the spontaneity of a reaction Draw and interpret heating and/or cooling curves Explain why gas molecules do not stick together at room T Explain why not all gas particles in a given sample move at the same speed at the same T Describe how intermolecular forces impact the specific heat of a substance Interpret a KMT diagram Compare average molecular speed of gas samples Illustrate gas particle behavior Chapter 6 A solution is a stable heterogeneous molecular mixture; it will not become “unmixed” over time. The factors that affect solubility are the enthalpy change (ΔH) and the entropy change (ΔS). A rule of thumb for predicting solubility is “like dissolves like”, but from a molecular perspective, this is not the reason why a solute dissolves in a solute. Also, it does not explain quantitative differences in solubility (for example, unlike dissolves at low levels into like). Molecules that have both polar and nonpolar parts will often assemble to form larger structures, through a process driven by a decrease in entropy (involving the water molecules, primarily). Understand and explain the terms: solute, solvent, solution, solvation, molarity, concentration, dilute, concentrated, micelle, colloid, emulsion Perform calculations using molarity (M) and volume of solution. Be able to calculate mass of solute needed, and determine volume of a dilute or concentrated solution. Interconvert between solution units such as % by mass, ppm and ppb. Relate temperature changes to the interactions that are broken and formed during the solution process. Describe the solution process in terms of entropy changes for both solute and solvent. Explain why free energy change is the thermodynamic function that allows us to predict solubility. (Why can’t we just use entropy or energy changes?) Predict which materials might be soluble in water given the interactions that hold their particles together. (“Like dissolves like” – but remember this is not the explanation for why something dissolves.)  Draw molecular level diagrams showing how solute and solvent interact, including how water molecules solvate ionic compounds. Predict which molecular compounds can form hydrogen bonds with water. Understand that solubility is not an all or nothing property and be able to estimate the relative solubilities of a range of given compounds. Explain why nonpolar molecules are not soluble in water, being sure to include the role of entropy in your discussion. Discuss how large molecules with both polar and nonpolar “parts” behave, including the role of the hydrophobic effect on how biomolecules fold, and how micelles and bilayers form. Understand and explain the terms colloids and emulsions. Explain how they differ from solutions. Explain the role of temperature in solubility. Explain how different inclusions in metals can alter the properties of the metal. (Consider steel, brass, bronze) Predict and explain the effect of temperature on the solubility of gases and solids. Identify intermolecular forces involved in the formation of a solution Describe the impact of relative size of polar and nonpolar portions of a molecule on their solubility in a nonpolar solvent or a polar solvent Describe the “like dissolves like” rule and explain why it works Define heat of solution (ΔHsoln) Predict which materials might be soluble in water given the interactions that hold their particles together (“like dissolves like”, but remember this is not the explanation for why something dissolves) Relate temperature changes to interactions that are broken and formed during the solution process Define heat of hydration (ΔHhyd) Describe the steps involved in forming a solution and whether the process is exo- or endothermic Draw molecular level diagrams showing how solute and solvent interact, including how water molecules solvate ionic compounds Predict which molecular compounds can form hydrogen bonds with water Describe the role of entropy (S) in the solution process Define saturated, unsaturated, and supersaturated solutions Describe the role and effect of temperature on the solubility of solids in water Describe the role and effect of temperature on the solubility of gases in water Define and calculate molality Use molality as a conversion factor Define and calculate mass percent Use mass percent as a conversion factor  Define and calculate ppm and ppb Interconvert between any two concentration units Define and identify substances as electrolyte or nonelectrolyte Define and calculate boiling point elevation Define and calculate freezing point depression Define and calculate molarity, as well as use it as a conversion factor Determine the molarity of ions in a given solution Chapter 7 Chemical reactions involve rearrangements of atoms to produce new chemical species with new bonds. Chemical reactions produce new species with properties than cannot be predicted. Chemical reactions involve only changes in arrangements of atomic cores and valence electrons. Chemical reactions can be classified by how the valence electrons behave during the course of the reaction (for example acid-base reactions or redox reactions). Energy changes in reactions arise from the changes in bond energies as bonds in reactants are broken (which requires an energy input) and new bonds are formed (which releases energy to the surroundings) Chemical reactions involve rearrangements of atoms to produce new chemical species with new bonds. Chemical reactions produce new species with properties than cannot be predicted. Chemical reactions involve only changes in arrangements of atomic cores and valence electrons. Chemical reactions can be classified by how the valence electrons behave during the course of the reaction (for example acid-base reactions or redox reactions). Energy changes in reactions arise from the changes in bond energies as bonds in reactants are broken (which requires an energy input) and new bonds are formed (which releases energy to the surroundings) Define Arrhenius acids and bases Define and recognize a neutralization reaction Define and identify common strong acids (HCl, HBr, HI, HNO3, H2SO4) and bases (NaOH, KOH) Define and identify weak acids and bases Compare and contrast weak and strong acids and bases Define Kw and the autoionization of water  Describe the relationship between [–OH] and [H3O+] in aqueous solution Calculate pH from [H3O+] and vice versa Calculate pH from [–OH] and vice versa Define a solution as acidic, basic, or neutral using the pH scale Calculate pH from [strong acid] Calculate pH from [strong base] Define Bronsted-Lowry acids and bases Identify conjugate acid-base pairs Use molecular scenes to predict net direction of acid-base reactions Define Lewis acids and bases Recognize a Lewis acid-base reaction Assign oxidation states to elements in atomic states, within compounds, and within polyatomic ions Identify substances undergoing oxidation and reduction Identify oxidizing and reducing agents in a given reaction Chapter 8: Rates of reactions depend on the probability that molecules will collide with enough energy to surmount the activation energy barrier. The extent of a reaction (position of equilibrium) is related to the free energy change from reactants to products and the temperature The position of equilibrium (but not the equilibrium constant) can be changed by changing concentrations of reactants or products. Explain what rate means when discussing a chemical reaction. Discuss the factors that affect rates of reactions, and explain how and why each factor affects the rate. Graphically determine the rate and the order of reaction from concentration vs. time data. Draw graphs that show how concentration changes with time for all reactants. Define and use half-life for a first order reaction. Draw and label reaction energy profiles when given a reaction mechanism, and show the effect of a catalyst. Use rate equation to determine which mechanism is more likely. Calculate an equilibrium constant when given concentrations at equilibrium. Discuss the concept of dynamic equilibrium and explain how it differs from a steady state. Predict which side the equilibrium lies from data such as the equilibrium constant or the Gibbs energy.  Use acid dissociation constants (Ka) to calculate pH of weak acid. Use the pH of a weak acid to calculate the Ka. Predict and explain how a position of equilibrium will shift when conditions are changed. Use the relationship between Q and K to predict how a reaction will shift. Define and differentiate between average, instantaneous, and initial rates of reaction Understand the impact of concentration and temperature on reaction rate Express reaction rate in terms of reactant and product concentration change Sketch a graph of concentration (both reactants and products) as a function of stoichiometry, as the reaction proceeds Define the rate law of a given reaction Understand that a rate law must be determined by experiment Define reaction order Determine reaction order from a given rate law Determine units on rate constant, given a rate law Determine reaction order, given a series of initial concentrations and initial reaction rates Determine the rate constant, given experimental data Graph experimental data using integrated rate laws to identify the reaction order as zero, first, or second order Determine the half-life of a first order reaction, given experimental data Determine amount of substance remaining, given an amount of time elapsed for a first-order reaction Interpret a graph of concentration vs. time, for a first order reaction, to answer questions about half-life Describe how changes in the number of molecular collisions impacts rate of reaction Calculate activation energy from experimental rate data Define transition state Recognize a transition state species and where it belongs on a reaction energy diagram Describe the effect of temperature on rate constant and rate Interpret a reaction energy diagram Determine if a proposed mechanism matches the experimentally observed rate law Describe how a catalyst impacts the activation energy, rate constant, and rate of a reaction Describe the relationship between forward rate and reverse rate of a reaction at equilibrium  Write the relationship between equilibrium constant and the rate constants for the forward and reverse reaction Write the equilibrium expression for a given reaction Explain why solids and liquids are not included as concentrations in the equilibrium expression Understand what a large, intermediate, or small equilibrium constant indicates about extent of reaction Define and write an expression for the reaction quotient of a given reaction (Q) Understand the relationship between Q & K and when they are equivalent Sketch a graph of concentration vs. time during a given reaction, of both reactants & products Describe the relationship between K for the forward reaction and K for the reverse reaction Define the equilibrium constant or reaction quotient of a reaction that is multiplied by a factor or reversed Express the equilibrium constant in terms of pressure Compare Q and K to determine reaction direction Use molecular scenes to determine reaction direction Solve for K using equilibrium quantities of reactants and products Solve for equilibrium quantities of reactants and/or products, given K and starting quantities of reactants and products Set up an ICE table to solve equilibrium problems Identify when able to use approximations in solving equilibrium problems Identify which direction a reaction will shift when disturbed by a sudden change in concentration, temperature, volume, or addition of a catalyst Chapter 9: Buffers can resist changes in pH Acidic and basic groups will protonate and deprotonate depending on the pH Reactions can be coupled by common intermediates so that unfavorable reactions can be driven by coupling to a favorable reaction Explain how and why a buffer can resist changes in pH, and relate this to the buffer capacity and range Calculate pH of buffer solutions, and the change in pH when a strong acid or base is added Predict the approximate pH of a salt, and explain what processes cause a pH change when it dissolves in water. Explain how and why reactions are coupled by a common intermediate, and give example Define an acid-base buffer Identify solutions that behave as buffers  Describe the common-ion effect Write reactions to show how a buffer absorbs a strong acid or strong base Calculate the pH of a buffer, given the concentrations of weak acid/conjugate base or weak base/conjugate acid Calculate the pH of a buffer after adding a strong acid or strong base Define buffer capacity Estimate the capacity of a given buffer Define buffer range Estimate the effective range of a given buffer Select a weak acid/conjugate base buffer or a weak base/conjugate acid buffer to maintain a given pH value Recognize the equivalence point(s) on an acid-base titration curve Differentiate between strong acid/strong base, weak acid/strong base, weak base/strong acid titration curves Identify the major species in solution at each point along any titration curve Calculate the pH at any point along any titration curve Define and identify the buffer region on any titration curve Use acid dissociation constants (Ka) to calculate pH of weak acid. Use the pH of a weak acid to calculate the Ka Predict and explain how a position of equilibrium will shift when conditions are changed Explain how and why a buffer can resist changes in pH, and relate this to the buffer capacity and range Predict the approximate pH of a salt, and explain what processes cause a pH change when it dissolves in water Explain how and why reactions are coupled by a common intermediate, and give examples