CHM 11600 Exam 3 Study Guide

What is a buffer solution?

• 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 the First Law of Thermodynamics.

• 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” If you want to learn more check out psyc1001 - psychology 1001 notes

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.

The total enthalpy change of a reaction is the sum of all of the steps/changes of the reaction.

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. We also discuss several other topics like acct3031

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. We also discuss several other topics like acc 203

• 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 If you want to learn more check out phys 284 concordia

• 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 If you want to learn more check out ams 5 ucsc

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: We also discuss several other topics like assimilation vs accommodation psych

▪ 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 Ebonds broken - Σ 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.