Learning Objectives: Chem 115 at West Virginia University Dr. Ratcliff (Spring 2018)

Transcription

1 All course materials, including lectures, class notes, quizzes, exams, handouts, presentations, and other materials provided to students for this course are protected intellectual property. As such, the unauthorized purchase or sale of these materials may result in disciplinary sanctions under the Campus Student Code. HOUR EXAM II perform Learning Objectives related exam 1 Learning Objectives related to basic stoichiometry use the coefficients of a balanced equation to convert between chemicals (moles moles, moles grams, grams grams). (For example: In the combustion of acetylene (C2H2), what mass of carbon dioxide can be produced from the reaction of 50. g of acetylene with an excess of oxygen?) identify a limiting reagent problem (amounts for two different reagents given) explain why an excess of one reagent would be used identify the limiting reagent and the excess reagent calculate the amount of product formed during a limiting reagent problem (For example: In the combustion of butane (C4H10), what mass of water can be produced from the reaction of 4.6 g butane with 15 g oxygen?) calculate the amount of excess reagent left over after the reaction is complete (For example: When bromine is reacted with aluminum, aluminum bromide forms. If 0.12 mol bromine is reacted with 0.10 mol aluminum, how many moles of which reagent are left?) determine how the following will affect the percent yield of a reaction (i.e., Will make percent yield less than 100% or greater than 100%): a competing reaction uses up some reactant or product to produce a second undesired product, the reaction proceeds less than 100% to completion, some of the product is stuck to the filter paper and cannot be removed, some of the product evaporates, the product is wet with solvent when it is weighed calculate the percent yield for a reaction (For example: An 11.2 g amount of ammonia is produced by reacting 10. g of nitrogen with 5.0 g of hydrogen. What is the percent yield for this reaction? Learning Objectives related to solutions and reactions in solution define soluble versus insoluble apply the solubility rules to predict whether an ionic compound is soluble in water (Remember, the solubility rules only apply to ionic compounds in water. Do not use them to predict whether acids (which in most cases are molecular compounds) are soluble in water. Most of acids that we deal with in Chemistry 115 are soluble in water.) understand and envision how soluble ionic compounds dissolve in water (e.g., Soluble ionic compounds are strong electrolytes and dissociate (split up into ions) 100% in water.) predict products for precipitation reactions identify spectator ions explain why most reactions take place in solution identify and define unsaturated, saturated and supersaturated solutions identify and define the solute, solvent, and resulting solution calculate the molarity of a solute in a solution (For example: A solution is made by dissolving We are what we repeatedly do. Excellence, then, is not an act, but a habit. Aristotle Page 1 of 5

2 18.6 g Ca(NO3)2 in enough water to prepare 525 ml of solution. What is the concentration of this solution?) use the molarity to calculate the moles of solute or mass of solute present (For example: You want to prepare 75.0 ml of a solution that is M in H3PO4. What mass of H3PO4 should be used?) explain dilution and how a more concentrated solution is diluted explain how addition of pure solvent affects the solution volume, moles of solute, and solution concentration calculate the concentration of a solution after dilution or volume of concentrated solution needed to perform a dilution, etc. (For example: What volume of 12 M H2SO4 solution is needed to prepare ml of 0.30 M H2SO4 solution? To prepare this solution, should the water be added to the 12 M acid or should the 12 M acid be added to the water?) calculate the stoichiometry of reactions occurring in solution (For example: What mass of lead(ii) nitrate is needed to react with 50.0 ml of 0.12 M sodium bromide solution? Hint: First write a balanced equation.) carry out calculations involving molarity of ions in solution (especially important when strong electrolytes are present) (For example: A 300. ml solution that is prepared from lead(ii) nitrate is found to be 0.33 M in nitrate ions. What mass of lead(ii) nitrate was used to prepare this solution? OR What is the concentration of chloride in a solution that is 0.50 M in zinc(ii) chloride? OR What is the concentration of sodium ions in a solution that is prepared by dissolving 2.0 g sodium carbonate in enough water to prepare 100. ml of solution?) Learning Objectives related to electrolytes, acids, bases, and acid-base titrations recite the definitions and characteristics of strong, weak and non-electrolytes identify substances that act as strong, weak, and non-electrolytes when dissolved in water write balanced molecular, full ionic, and net ionic equations (For example: Predict products and write the net ionic equation obtained when aqueous solutions of cobalt(iii) nitrate and hydrosulfuric acid are mixed.) identify spectator ions recite the definitions of acids (Definitions: Arrhenius acid produces H + /H3O + when dissolved in water; Bronsted Lowry acid H + donor and loses the H + ) recite the definitions of bases (Definitions: Arrhenius base produces OH - when dissolved in water; Bronsted base H + acceptor and gains the H + ) recite the characteristics of strong & weak acids and bases identify strong & weak acids and bases and envision how they dissolve in solution identify nonmetal and metal oxides as acid or base anhydrides and be able to predict the acid or base obtained when the oxide is dissolved in water predict products for acid-base reactions define the equivalence point, end point, titration, indicator and generally understand how a titration is carried out calculate the concentration of a solution given titration data (For example: During a titration, the phosphoric acid in ml of solution is completely neutralized by addition of ml of M KOH solution. What is the concentration of the phosphoric acid solution?) classify chemical reactions as precipitation, acid-base, or redox reactions and explain the driving force for each We are what we repeatedly do. Excellence, then, is not an act, but a habit. Aristotle Page 2 of 5

3 understand that H2CO3, H2SO3, or NH4OH are unstable or somewhat unstable and if predicted as products of a precipitation or acid-base reaction decompose to other products (H2CO3 gives CO2(g) + H2O(l); H2SO3 gives SO2(g) + H2O(l); and NH4OH gives NH3(g or aq) + H2O(l)) Learning Objectives related to oxidation-reduction reactions use the rules for assigning oxidation numbers to assign oxidation numbers identify atoms being oxidized and reduced identify oxidizing agents and reducing agents determine the number of electrons transferred during a redox reaction balance simple (no polyatomic ions present as reactants or products) redox reactions by either the oxidation number method or the half-reaction method balance more complicated redox reactions (those involving polyatomic ions) by either the oxidation number method or the half-reaction method. use the metal activity series to determine if a reaction between a metal and a metal cation will occur (For example: given Na is a stronger reducing agent than Al, predict if Na + will react spontaneously with Al.) Be able predict products of the spontaneous reaction and balance the resulting redox reaction. given a series of reactions that occur or do not occur (no reaction), predict the metal activity series and give a list of metals in order of decreasing ability to act as reducing agents. Recall the non-metal activity series (F2 > O2 > Cl2 > Br2 > I2 > S) and that this lists the nonmetals according to decreasing ability to act as oxidizing agents use the non-metal activity series to predict whether reactions will occur spontaneously between two non-metals (For example: Determine which of the following reactions will occur spontaneously? O2 + NaCl ; KBr + I2 ; Cl2 + Na2S ) Learning Objectives related to heat and thermochemistry produce the following equations from memory: E = ½mv 2 ; c = υ λ; E = hυ; E or E = hc/λ; λ = h/mv; q = mc T; q = K T; q = nc T; heat gained = - heat lost; E = q + w and w = -P V; Hrxn = Hproducts - Hreactants. explain, define and use proper terminology related to thermochemistry (e.g., temperature, heat, kinetic energy, potential energy, energy units (J, kj, cal, kcal, Cal), system, surroundings, internal energy, first law of thermodynamics, work of expansion, isolated system, insulated system, internal energy, enthalpy, exothermic, endothermic, calorimetry, standard conditions, etc.) convert between the various energy units (the conversion of 1 cal = J will be given but a conversion like 1 kj = 1000 J will not be given) distinquish between the concepts of kinetic energy and potential energy explain that kinetic energy and potential energy can be inter-converted. explain how chemical reactions can do work of expansion be able to explain and identify the directionality of energy (in the form of work or heat) with reference to the system (For example: If the system of a chemical reaction is expanding, energy in the form of work is leaving the system and so w = negative.) use the change in temperature of the surroundings to determine whether a reaction is We are what we repeatedly do. Excellence, then, is not an act, but a habit. Aristotle Page 3 of 5

4 exothermic perform calculations relating heat, temperature change and specific heat capacity. (For example: When 8.66 kj of heat is added to 665 g of a substance, its temperature changes from 10.9 C to 87.3 C. What is the specific heat of this substance (in J/g- C)? or If 4.5 kj of heat is added to 425 g of a substance (c = 0.34 J/g- C) at 22 C, what is the final temperature of the substance?) Relate the First Law of thermodynamics to solve heat gained/heat lost problems (For example: In an insulated system, 25 g of a substance (c = 0.45 J/g- C) at 210. C was added to a mass of liquid water (c = J/g- C) at 12 C. If the temperature of the resulting mixture was 17 C, what mass of liquid water was present?) solve calorimetry problems use the thermochemical equation to calculate/relate the heat absorbed or released to amount of chemical reacted or produced (For example: Given, 2 HgO(s) 2 Hg(l) + O2(g) H = kj, how much heat is involved (absorbed or released) when 15 kg of HgO is decomposed?) use Hess s law to calculate enthalpies of reactions given 2 or more thermochemical equations. write the balanced reaction that goes along with the standard enthalpy of formation for any chemical compound calculate standard enthalpy change (or standard heat of reaction) for any reaction given tabulated standard heat of formation ( Hf ) values recall the standard enthalpies of formation for elements in their standard states under standard conditions ( Hf = 0) recall the standard state of common elements (For example: standard state of fluorine is F2(g), for carbon is C(s,graphite), and for nickel is Ni(s)) Learning Objectives related to light and atomic spectra, and wave-particle duality define terms dealing with light energy and electronic structure: wavelength, frequency, amplitude, speed of light, Hertz, electromagnetic radiation, photon, quanta/quantum, white light, continuous spectrum, discrete/line spectrum, ground state, excited state, Bohr model, wave-particle duality, de Broglie wavelength, electromagnetic spectrum understand the relationship between frequency, wavelength and energy (directly related or inversely related?) recall the order of light in the electromagnetic spectrum in terms of energy and therefore frequency and wavelength (For instance: x-rays are more energetic and have higher frequency but lower wavelength than visible radiation) understand the dual nature of light (electromagnetic radiation), properties associated with its wave nature (frequency and wavelength, that it can undergo refraction, diffraction, and constructive and destructive interference) and properties associated with its particle nature (that it comes in discrete packets/particles or quanta of light energy called photons) calculate properties of light given other properties (given the frequency of light, calculate the energy or wavelength or given the wavelength of light calculate the frequency or energy) explain what atomic line/discrete spectra tells you about the energy levels with the atom (a continuous set of energy levels or discrete set of energy levels?) explain whether when an electron moves from one energy level to another within the atom is a photon of light is absorbed or emitted Use the Rydberg equation to calculate the wavelength of light emitted by excited state We are what we repeatedly do. Excellence, then, is not an act, but a habit. Aristotle Page 4 of 5

5 hydrogen and/or the difference in energy between the two energy levels (indexed by the n- value) or given the wavelength of light emitted be able to calculate one of the energy levels (nvalue) Explain the relationship between the wavenature (specified by wavelength) and the particle nature (specified by momentum or mass x velocity) as shown by the de Broglie wavelength equation. Be able to calculate the de Broglie wavelength and understand that as the mass increases, the de Broglie wavelength decreases and the wavenature becomes less important the more massive the particle explain the relationship between the wavenature (specified by wavelength) and the particle nature (specified by momentum or mass x velocity) as shown by the de Broglie wavelength equation. Be able to calculate the de Broglie wavelength and understand that as the mass increases, the de Broglie wavelength decreases and the wavenature becomes less important the more massive the particle Learning Objectives related to quantum atomic concepts recognize and draw the shapes of the s, p (px, py, and pz) and d (dxy, dxz, dyz, dx2-y2, and dz2) orbitals, explain, define and recall the values for the four different quantum numbers, determine the angular momentum quantum number for s, p, and d, etc. subshells and the values for the magnetic quantum numbers for orbitals within a given subshell (e.g., f subshells have l=3 and possible ml values of -3, -2, -1, 0, +1, +2, +3 and thus there are 7f orbitals within any f subshell), qualitatively discuss the quantum mechanical theory developed by Schrodinger (i.e., how it treats the motion of electron about the nucleus as a wave using mathematical theory of wave mechanics) and the results that come from quantum mechanics (i.e., wavefunction, energy, quantum numbers, and square of wavefunction), discuss the meaning that quantum mechanics gives to orbital shapes and that orbital shapes essentially give information on the probability of finding the electron in the region of space subtended by the orbital boundaries, understand that the orbital shapes do not give information on the path the electron takes about the nucleus, predict allowed sets of quantum numbers for an electron in an atom (i.e., Can an electron in an atom have the following set of quantum numbers? n = 5, l = 3, ml = -2, ms = -1/2. Why or why not? How about the following sets? n = 6, l = 2, ml = -2, ms = -1 or n = 4, l = 1, ml = -2, ms = ½ or n = 5, l = -3, ml = 0, ms = -1/2), predict correct orbital designations (e.g. An electron is in an orbital with n = 3, l = 1, and ml = 0. What is the orbital designation? Which of the following orbital designations are not allowed? 5p, 4g, 3s, 1p, 2d, 7f. Explain.), determine the number of electrons that can reside in any orbital, subshell, or shell (e.g., How many electrons can reside in a 5px orbital? in a 7f subshell? in the l = 2 subshell? in the fourth shell?), explain the Pauli Exclusion Principle, the Aufbau Principle, and Hund s Rule and be able to determine whether the electrons filled into an atom obey or disobey these rules, determine the electronic configurations (using both the diagonal filling pattern and the periodic table) for neutral atoms, cations and anions We are what we repeatedly do. Excellence, then, is not an act, but a habit. Aristotle Page 5 of 5