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CHM 11600 Exam 3 Study Guide

by: Gayatri

CHM 11600 Exam 3 Study Guide CHM 116

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Includes learning objectives to know for Exam 3
CHM 116
Dr. Nash
Study Guide
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This 3 page Study Guide was uploaded by Gayatri on Friday April 3, 2015. The Study Guide belongs to CHM 116 at Purdue University taught by Dr. Nash in Winter2015. Since its upload, it has received 745 views. For similar materials see CHM 116 in Chemistry at Purdue University.


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Date Created: 04/03/15
CHM 11600 Exam 3 Study Guide • Determine whether an aqueous solution of a salt will be acidic, basic or neutral given values of Ka and Kb for conjugate acid-base pairs. o For any conjugate acid-base pair, the relationship between Ka and Kb is that (Ka)(Kb) = Kw = 1 x 10^-14 o To determine whether a solution is acidic, basic or neither, we need to find the pH. § First, set up the equation with the acid-base pair § Then, write equilibrium equation and set equal to Ka or Kb § Find unknown value, this will be H3O+ for Ka or OH- for Kb. § pH = -log(H3O), or 14 – (-log(OH-)) § pH<7 = acidic, pH 7 = neutral, pH>7 = basic. • Describe a "buffer solution”. o A buffer solution is a solution that decreases the impact on pH from the addition of an acid or a base, and it is made up of a conjugate acid-base pair (weak acid + conj. Base or weak base + conj. Acid) • Describe how either an acidic or basic buffer solution is prepared. o Step 1. Choose the conj.acid-base pair o Step 2. Calculate ratio of buffer component concentrations using the Henderson-Hasselbach equation o Find the buffer concentration, usually, one must be a solid and one must be a stock solution. • Describe "buffer capacity" o A measure of the strength of the buffer, its ability to maintain the pH following the addition of a strong acid or base. o If buffer concentration increases, capacity increases o If closeness of component concentration to buffer concentration increases, capacity increases • Describe the “common ion effect” o Remember the Le-Chatalier’s principle, which states that if equilibrium becomes unbalanced, the reaction will shift and adjust itself to restore the balance. This is essentially the idea of the common ion effect. o If a common ion is added to a weak acid or weak base equilibrium that already contains that ion, then the equilibrium position will shift towards the other side in order the consume the ion. • Describe how a buffer solution (either acidic or basic) is able resist large changes in pH when small amounts of either acid or base are added to the buffer solution. o A buffer is a special solution that can resist pH changes because it consists of conjugate acid and base pairs present in large amounts at equilibrium, so they are able to neutralize whatever small amounts of acids and bases are added to H3O+ and OH- o The Henderson-Hasselbach equation used to calculate change in pH: pH = pKa + log ([conj base]/[acid]) • Describe the thermochemical "universe" o The system is defined as the part of the universe that we are focusing on in a reaction, and the surroundings are everything else. The sum of the system and the surroundings make up the universe. • Describe “internal energy” o The internal energy is the sum of all the energies (potential and kinetic) in a particle in a system (E). o The internal energy changes when reactants in a chemical system change to products. • Describe two ways that a chemical system can change its internal energy. o Releasing energy to surroundings E(final) < E(initial), ΔE < 0 o Absorbing energy from the surroundings E(final) > E(initial), ΔE > 0 • Describe “heat” and “work”. o Heat (q): thermal energy, transferred as a result of difference in temperature between the system and the surroundings o Work (w): all other forms of energy, transferred when an object is moved by force. • Describe the two kinds of work normally associated with chemical reactions. o Work done by the system on the surroundings, w-, ΔE- o Work done on the system by the surroundings, w+, ΔE+ • Describe the sign conventions for changes in internal energy of the system when work is done by, or on, the system and/or heat is gained, or lost, by the system o When heat (q) is absorbed, +; when heat (q) is released, -. o When work (w) is done on system, +; when work (w) is done by system, -. o ΔE sign depends on work and heat • Describe the First Law of Thermodynamics o The First law of Thermodynamics states that the total energy of the universe is always constant. ΔE universe = ΔE system + ΔE surroundings = 0 • Describe a “state function” and list several examples of state functions o A state function is a property that is dependent only on the current state of the system, and not the path o Examples of state functions include enthalpy, pressure, volume, and temperature. • Describe enthalpy, and its relationship heat flow at constant pressure. o Enthalpy of a system is the internal energy plus product of pressure and volume. Since p is constant, enthalpy, which is ΔH = ΔE + PΔV, works out to be qp = ΔH, and stands for heat absorbed or released at constant pressure. • Distinguish “exothermic” and “endothermic” processes. o Exothermic reactions are those that release heat, -ΔE o Endothermic reactions are those that absorb heat, +ΔE • Describe each of the following types of enthalpy change: o heat of reaction (ΔHrxn) = energy absorbed/released through a chemical process o heat of combustion (ΔHcomb) = energy absorbed/released when a compound is broken into its separate components o heat of formation (ΔHf ) = energy absorbed/released when a compound is formed from its components o heat of fusion (ΔHfus) = heat absorbed/released when a substance is converted from solid to liquid o heat of vaporization (ΔHvap) = heat absorbed/released when a substance is converted from a liquid to a gas • Describe why the heat of vaporization for a substance is always larger in magnitude that the heat of fusion for the substance o ΔH vaporization is higher than ΔH fusion because in vaporization, all the bonds are broken since the substance is converted to gas, and all ideal gas molecules have no IMFs. As for fusion, only a few bonds are broken since the substance is converted from solid to liquid. • Describe how the value for ΔH is experimentally determined. o To find how much heat is absorbed/released in a reaction, we construct “surroundings” that retain the heat as reactants become products, and then note the temp change. In the lab, a calorimeter takes the place of the surroundings, which is device used to measure the heat released or absorbed by a chemical/physical process. There are two types of calorimeters that can be used: § Coffee cup calorimeter: for processes that occur at constant pressure, a “constant pressure calorimeter” is used, and –q(solid/sys) = -q(surr/H2O), where q = mass x specific heat capacity x ΔT • If a reaction takes place at constant pressure, the q(rxn) = ΔH § Bomb calorimeter: for processes where the volume is constant, a bomb calorimeter is a much more precise measurement, and helps us know the heat capacity of the entire calorimeter. The energy change measured is the heat released at constant volume, which equals ΔE = q + w • Describe heat capacity, specific heat capacity and molar heat capacity. o Heat capacity: quantity of heat required to change an object’s temp by 1K o Specific heat capacity (c): quantity of heat needed to change temp of 1 gram of object by 1 K. (J/g*K) o Molar heat capacity (C): quantity of heat needed to change temp of 1 mole of object by 1 K (J/mol*K) • Calculate ΔH by using: o bond energies § ΔH = Σ E bonds brokenΣ Ebonds formed o Hess’s law of heat summation § The total enthalpy change of a reaction is the sum of all of the steps/changes of the reaction o standard heats of formation § ΔHformation = Σ ΔHformation (products) - Σ ΔHformation (reactants) • Describe “homolytic” and “heterolytic” bond dissociation energies and the difference(s) between them. o Homolytic bond breaking is where each atom gets one electron each, even split, and this gives rise to free radicals. In calorimetry, homolytic BDEs are measured and recorded. o Heterolytic bond breaking is where one atom takes both electrons, giving rise to positive and negative ions. • Describe (list)the limitation(s), if any, of using bond dissociation energies, Hess’s Law and standard heats of formation for calculating ΔH. o A limitation of using bond energies (Hess’ Law) would be that they are always based on average bond enthalpies, so the actual one might be off, and they are only true for gaseous species, so for example the bond enthalpy of H2O(l) would be the same as that for H2O(g). • Describe “standard state” conditions and why they are needed in thermochemistry. o Standard states are a set of specific conditions that have been established to be able to study and compare reactions in thermochemistry. § For a gas, std. state is 1 atm and ideal behavior § For an aqueous solution, concentration =1 M § For a pure substance (element/compound), the most stable form at 1 atm and 25 degrees Celsius. • Describe why many compounds have negative standard heats of formation. o This is because most compounds have exothermic formation reactions. Under standard conditions, heat is released when most compounds form their elements.


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