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REVIEW SHEET SEMESTER TWO PREAP CHEMISTRY Bonding 1. 2. 3. 4. .In ionic bonding electrons are ___________________. In covalent bonding electrons are ________________. Why do atoms bond?________________________________________________ Using electronegativities determine what type of bonding, molecular shape and molecular polarity for the following molecules. Compound Type of bond Molecular shape Molecular polarity NaCl CH4 H2O NH3 AlCl3 Kinetic Theory and States of Matter 1. Name the four states of matter. 2. Name the three intermolecular forces. Which is the strongest ?_________________ Which is the weakest?________________ 3. What are the four elements necessary for H-bonding? 4. Arrange the following molecules in order of decreasing boiling point. CH4 H2S N2 NH3 5. Why does ice float? 6. At what temperature is water most dense? Why? 7. What is rate of flow? 8. What factors affect rate of flow? 9. What substance will decrease the surface tension of water?______________________ 10. A highly volatile substance will have _______ intermolecular forces. 11. What type of intermolecular force will be the strongest in polar molecules?________________Nonpolar molecules?_______________ 12.Adding heat to a solid will ____________ increase/decrease the kinetic energy and cause the intermolecular forces to ________________. 13. Of the three phases of matter, ______________ has the strongest attractive forces, and _________________ has the weakest. 14. Where will water boil at a lower temperature 250 feet or 10,000 feet above sea level? _____________________________________________ Why?_____________________ 15. Explain on the molecular level how solids, liquids and gases compare. Which has the most kinetic energy? 16. Label the heating curve of water wit liquid, solid, gas, boiling point, freezing point, melting point, condensation point. Show in which direction energy increases. 17. How much energy would it take to melt 10 moles of ice if the molar heat of fusion for water is 6.009 KJ/mole? 18. q=m(ΔT) Cp How much heat does a 100g cube of ice need to go from -200Cto 00CC? The specific heat of ice is 2.06 J/g 0C. 19. Draw a phase diagram for CO2 and label the normal melting point, normal boiling point, triple point, critical temp and critical pressure. The Gas Laws 1. If pressure decreases, temperature _____________, and volume ___________________. 2. The number of collisions of gas molecules is measured by ________________. List three ways to increase the number of collisions. ____________________________________ ___________________________________ ____________________________________ 3. Gases have _______________kinetic energy, __________________distances between molecules and therefore ____________________ compressed. 4. Standard temperature is ___________________0C and ____________________K 5. Convert. a. 160C = _____________________K b. 16K = _____________________0C 6. a. b. c. d. e. Convert 0.95 atm = ________________________mmHg 790 mmHg = ______________________kPa 125 kPa= _________________________mmHg 100 mmHg = ______________________ Torr 95 kPa = __________________________atm 7. If gas A travels slower than gas B at 350C, then what can be said about the mass of A and B? 8. LAW Charles’ CONSTANT FORMULA Bolye’s Gay-Lussac’s Ideal N/A Combined Gas N/A 9. At 1120C, the volume of a gas is 2.3L. Find the volume if the temperature is changed to 650K and the pressure remains the same. 10. A gas collected when the pressure is 80 kPa has a volume of 175 mL. What volume will the gas occupy at standard pressure? 11. At 327K, a gas exerts a pressure of 0.7 atm. If the volume remains constant, what will happen to the pressure if the temperature is raised to 876K. 12. What is the volume, in liters, of 0.250 mol of oxygen gas at 20.00C and 0.974 atm pressure? Solutions 1. What are the three factors affecting rate of solubility? 2. Define Saturated Unsaturated Supersaturated 3. What is meant by dissociation? 4. In a solution, the ___________________ is dissolved in the ____________________. 5. When two substances will not dissolve in one another the are said to be _____________. When they will the are _________________________. 6. Calculate the molality (m) prepared by dissolving 15 g of NaCl in 500 g of water. 7. A 250 mL solution has a molarity (M) of 2.5 M. How many moles of solute are contained in this solution? 8. A solution contains 10.0 g of CuSO4 in 100mL of solution. What is the percent (m/v) of the solution? 9. 10 mL of a 3M solution of KCl is diluted with 350mL of water. What is the molarity of the diluted solution? Reaction Rates and Equilibrium 1. List the parts to the collision theory. 2. In the following reaction mechanism, identify the intermediates, and give the net equation. Step 1 2NO(g) N2O2(g) Step 2 N2O2(g) + O2(g) 2NO2(g) 3. What is a catalyst?___________________________________________ 4. What kind of catalyst is in your car?_____________________________ 5. What is an inhibitor?_________________________________________ 6. List the five factors that influence reaction rates. ___________________________________ __________________________________ ____________________________________ ___________________________________ ___________________________________ 7. What is the sixth factor that only applies to gases?___________________________ 8. In the following one-step reaction give the rate law. 2NO(g) + O2(g) 2NO2(g) What order of reaction is it overall? In terms of NO? In terms of O2? 9. Given the following reaction mechanism, write the rate law. 2NO N2O2 FAST N2O2 + O2 2NO2 SLOW 2NO + O2 2NO2 10. What is the rate determining step for the above reaction? 11. Draw and label an exothermic and an endothermic energy diagram. Label them with Ez, Ez’, ΔE, reactants, products, and activated complex. Make sure and label whether ΔE is positive or negative. 12. Calculate the equilibrium constant Keq for the following reaction. 2A <=> 2C + D At equilibrium, the concentration of A is 6.59 X 10-4 M, the concentration of C is 4.06 X 10-3 M, and the concentration of D is 2.03 X 10-3 M. 13. Explain how a system at equilibrium responds to a stress and list factors that can be stresses on an equilibrium system. 14. Explain how decreasing volume of the reaction vessel affects each equilibrium. a. 2SO2(g) + O2(g) <-- > 2SO3(g) b. H2(g) + Cl2(g) < -- > 2HCl(g) 15. Decide whether higher or lower temperatures will produce more CH3CH0 in the following equilibrium. C2H2(g) + H2O(g) < --> CH3CHO(g) ∆H = -151 kJ 16. Make sure you can calculate rate law from data given! Thermochemistry 1. Calculate the Gibb’s Free Energy change for the following reaction at 25oC. Ca(s) + 2H2O(l) Ca(OH)2(s) + H2(g). Given Hrxn= -411.6 kJ/mol; Srxn=.0318 kJ/mol (K) 2. Give the formula for Hreaction and state Hess’s Law. 3. Use Hess’s law and the changes in enthalpy for the following two generic reactions to calculate ∆H for the reaction 2A + B2C3 2B + A2C3. 2A + 3/2C2 A2C3 ∆H = -1874 kJ 2B + 3/2 C2 B2C3 ∆H = =285 kJ 4. Use standard enthalpies of formation from the table to calculate ∆Hrxn for the following reaction. P4O6 (s) + 2O2(g) P4O10(s) 5. Is the following reaction endothermic or exothermic. How much energy is gained or lost if 2 moles of iron reacts completely with the oxygen? 4Fe(s) + 3O2(g) 2Fe2O3(s) + 1625kJ 5. How can we predict if a reaction is spontaneous? Complete the table. ∆H ∆S ∆G Negative (-) value Positive (+) value Negative (-) value Negative (-) value Positive (+) value Positive (+) value Positive (+) value Negative (-) value Acids and Bases 1. List 4 properties of acids and 4 properties of bases. 2. Define an Arrhenius Acid and an Arrhenius Base 3. Define a Bronsted-Lowry acid and base. 4. Define a Lewis acid and base. 5. The pH of an acid is ____7, a base is _____7 and a neutral solution is _______7. 6. Give the formula to find pH. 7. How do you find pH if you are given [OH-] ? 8. What are the products of a neutralization reaction? 9. Label the acid, base, conjugate acid, and conjugate base for the following equation. NH3 + HNO3 NH4+ + NO3- 10. Find the pH of the following and label as acid, base(alkaline), or neutral [H+]= 1.0 x 10-5 [H3O+] = 2.0 x 10-13 [OH-] = 3.0 x 10-4 11. Given pH = 4.3, find pOH, [OH], [H3O+1] 12. Write formulas for and give names for the following acids. H2SO4 HNO3 HF HI Phosphoric acid Sulfurous acid 13. Write formulas for and give names for the following bases. NaOH Ca(OH)2 NH3 Barium hydroxide Magnesium hydroxide 14. If 20.00 mL of a 0.01 M solution of HCl is titrated with NaOH, 15.00 mL of NaOH is used at the endpoint. What is the molarity of the base? 15. What is the Ka of an acid that has a [H+] of 2.5 x 10-3M and the concentration of athe acid is .2M? 16. If the concentration of [Ag+1] is 2.53 x 10-4 M, what is the Kspof Ag2CO3? REDOX 1. Give the oxidation numbers of all the elements in the following compounds. KCN MgSO4 PO4-3 CaCO2 2. Identify the following as a redox or nonredox H2 2H 2Na + Cl2 2 NaCl HCl + NaBr HBr + NaCl 3. The oxidation number of a free element is always ____________________. 4.The most active reducing agent among the elements is ____________________. 5. For each of the following equation s determine whether oxidation or reduction is happening. K K+ + eS + 2e- S-2 6. What is the difference between a Voltaic cell and a electrolytic cell.? 7. Calculate Eocell for the following. a. Na+/Na and Ni2+/Ni b. Cu2+/Cu and Ag+/Ag 8. Identify the compound being oxidized, reduced, the oxidizing agent, and the reducing agent. Write the half reactions and balance. Sb + H2SO4 Sb2(SO4)3 + SO2 + H2O 9. Draw and label a voltaic Cell with the following cell notation; Zn/Zn+2//Cu+2/Cu 10. Draw and label an electrolytic cell that shows silver electroplating of a nickel ring.