Chemistry 1152 Final Exam Study Guide

Transcription

1 Chemistry 52 Final Exam Study Guide Exam will have 200 points. You will have from :30 pm 4:00 pm on Saturday, May 3 rd to complete the exam. The exam will be a mix of multiple-choice and show your work problems, following the format of the quizzes and mid-term exam. Review the quizzes, mid-term exam, homework problems, class notes and textbook. Chapter 3. Describe chemical equilibrium using the terms forward and reverse reactions and dynamic process. 2. Write the equilibrium constant in terms of the equilibrium concentration of products and reactants and their respective stoichiometric coefficients for both homogenous and heterogeneous equilibria. 3. Be able to convert between K p and K c. 4. Determine equilibrium constant given equilibrium concentration data. 5. Show that if a reaction can be expressed as the sum of two or more reactions, the equilibrium constant for the overall reaction is given by the product of the equilibrium constants of the individual reactions. 6. Relate equilibrium constant to rate constants from chemical kinetics. 7. Describe the relationship between reaction quotient and equilibrium constant and predict the direction a reaction will proceed to reach equilibrium. 8. Use the concepts of equilibrium to determine concentration of all species in a solution. 9. Use Le Châtelier s Principle to describe how changing concentration, volume, pressure, or temperature will shift the reaction so that a equilibrium will be maintained. 0. Describe the effect of a catalyst has on equilibrium concentrations. Chapter 4. Describe the properties of elemental hydrogen. Include its appearance, molecular structure, occurrence in nature, laboratory synthesis, industrial synthesis, and industrial use. 3. Classify binary hydrides as ionic, covalent, or metallic. 5. Classify oxides as basic, acidic, or amphoteric. 6. Identify compounds as oxides, peroxides, and superoxides.

2 Chapter 5. Compare and contrast Arrhenius, and Brønsted acids and bases. 2. Describe what is meant by conjugate acid-base pairs and give several examples. 3. Use K w to determine [H + ] and [OH - ] of solutions. 4. Discuss the ph scale and calculate ph and poh given either + ][H or [OH - ]. 5. Define strong and weak acids and bases and give several examples of each. 6. Relate properties of conjugate acid-base pairs. 7. Determine K a from experimental data. 8. Calculate ph, [H + ], weak acid concentration, and conjugate base concentration given K a and the initial concentration of the weak acid using the quadratic equation or using an appropriate approximation, as needed. 9. Calculate percent ionization for a weak acid. 0. Calculate ph, [OH - ], weak base concentration, and conjugate acid concentration given K b and the initial concentration of the weak base using the quadratic equation or using an appropriate approximation, as needed.. Show the relationship between a, K, K b, and K w. 2. Calculate concentrations of all species present at equilibrium for diprotic and polyprotic acids. 3. Relate molecular structure and the strength of acids. 4. Predict the relative strengths of oxoacids. 5. Describe salt hydrolysis and explain how some salts produce neutral solutions, some acidic solutions and others basic solutions. 6. Calculate the ph of salt solutions and determine the percent hydrolysis. 7. Describe acid-base properties of oxides and hydroxides. 8. Give several examples of Lewis acid-base reactions. Chapter 6. Describe the common ion effect as a special case of Le Châtelier s principle. 2. Use the Henderson-Hasselbalch equation to determine the ph of a solution containing a weak acid (weak base) and its conjugate base (conjugate acid). 3. Describe what a buffer solution is and its importance in chemical and biological systems. 4. Calculate the ph of a buffer solution using the Henderson-Hasselbalch equation. 5. Calculate the ph of a buffer solution after the addition + of or HOH_. 6. Describe how to prepare a buffer of a desired ph. 7. Predict the ph profile of a strong acid-strong base titration and calculate the ph at any stage of the titration. 8. Predict the ph profile of a strong acid-weak base (or strong base-weak acid) titration and calculate the ph at any stage.

3 9. Describe common acid-base indicators and suggest the correct method of selection for a specific titration. 0. Use the concepts of equilibrium to relate ion product, Q, with sp to Kpredict if a solution is unsaturated, saturated or supersaturated.. Distinguish between solubility product, molar solubility and solubility. 2. Calculate the solubility of an insoluble ion when a common ion is present. 3. Describe how changing ph can effect solubility. 4. Use the concept of equilibrium, formation constant f) and (K complex ion formation to predict ion concentration in solutions. 5. Describe how solubility product principle is used in qualitative analysis. Chapter 7 2. Provide several examples of spontaneous processes. 4. Define the standard conditions under which thermodynamic properties are measured (25 C, Molar and atm). 5. Define entropy using the terms disorder or randomness. 6. Justify why entropy of one mole of steam is greater then the entropy of one mole of water. 7. Predict the sign on the change in entropy S) ( for common processes and chemical reactions... Predict if a reaction is spontaneous, spontaneous in the reverse direction or at equilibrium from the sign on G. 2. Use thermodynamic tables to determine G of a reaction, given G f, H f, and S s Hess Law. 3. Rationalize the direction of a spontaneous reaction given H, S and T. 5. Be able to calculate quantities using the relationship between G and G, the gas constant, temperature and the reaction quotient G ( = G 0 + RT lnq). 6. Relate G and K, equilibrium constant G( 0 = RT lnk). 7. Calculate K, equilibrium constant, using data from thermodynamic tables.

4 Chapter 8. Describe the concept of redox reactions using such terms as reduction, oxidation, reducing agents, oxidizing agents and oxidation numbers. 2. Be able to calculate oxidation numbers and identify the oxidizing agent and reducing agent for redox reaction. 3. Describe an electrochemical cell using such terms as oxidation, reduction, galvanic cell, electrolytic cell, anode, cathode, half-cell reactions, salt bridge, cell voltage and emf. 4. Use standard cell diagrams to describe an electrochemical cell. 5. Use standard reduction potentials to predict the emf of a cell. 6. Define what SHE (standard hydrogen electrode) means and relate its significance to the standard reduction potential table. 7. Predict the outcome of reactions based on standard reduction potentials and overall. E 8. Mathematically relate G, K and E cell to each other and to reactions that are spontaneous, at equilibrium or are non-spontaneous. Potentially Useful Information: Happy studying!! Good Luck!! atm = 760 torr = 760 mm Hg N o = x 0 23 R = L.atm / (mole. K) = 8.34 J /(mole. K) = kj/(mol.k) Charge on one electron =.60 x 0-9 C D = x 0-30 Cm P l 2 = P Hvap R T T 2 year = 3,557,600 s year = days day = 24 hours hour = 60 min min = 60 s /2 /2 /2

5 Information that will given on the final exam: atm = 760 torr = 760 mm Hg N A = 6.022x0 23. R = L atm mol K = 8.34 [ ln A ] 0 [ A] = kt J mol K ln P 2 P [ A ] [ A ] = kt 0 ( ) = H vap R ( T ) T 2 ln( 2) = kt 2 [ A ] = kt 2 0 ln k ( k 2 ) = E a R ( T T 2 ) K w = ph = pk a + log base or acid HA ph = log K a A x = b ± b2 4ac 2a G = H T S G = Go +RT lnq G = RT ln K G = nfe