Chemistry 1B Notes
Chemistry 1B Notes Chemistry 1B
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Popular in Chemistry
This 7 page Bundle was uploaded by Akash Patel on Wednesday June 15, 2016. The Bundle belongs to Chemistry 1B at University of California Riverside taught by Glen millhauser in Winter 2016. Since its upload, it has received 11 views. For similar materials see General Chemistry in Chemistry at University of California Riverside.
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Date Created: 06/15/16
THERMODYNAMICS Spontaneous Processes: ● Spontaneous Process: a process that proceeds without outside intervention ● Nonspontaneous Process: process that cannot happen on its own ● Experimental conditions like temperature and pressure affect spontaneity ○ nonspontaneous process can become spontaneous under conditions like temperature and pressure ● Many exothermic reactions including combustion reactions and metal corrosion are spontaneous ○ some exothermic reactions aren’t spontaneous and some endothermic reactions are spontaneous ○ depends on reaction conditions ● Endothermic reactions are spontaneous due to their particles being more spread out and/or having more freedom of motion than the particles in their reactants Entropy: ● When an isolated thermodynamic system undergoes a change where the particles in the system disperse into a larger volume and/or gain freedom of motion, the process is always spontaneous ○ System experiences an increase in entropy ● Second law of thermodynamics: entropy of an isolated thermodynamic system always increases during a spontaneous process ● Entropy (S): a measure of how dispersed the energy of a system is ● Energy Dispersion ○ In the lab experiment, energy initially confined to the metal flows spontaneously from the hot metal into the cool water and is dispersed throughout the metal/water system ○ Can happen when all the components of a system, or system and its surroundings are at the same temperature ● The kinetic molecular theory can help predict that gases mix spontaneously but it can be predicted with the statistics of energy dispersion and entropy ○ KMT: random motion of gas particles/ diffusion causes expansion of gas particles ○ Statistics of energy dispersion and entropy: most probably state is the observed state ● Statistics of energy dispersion and entropy comes to the conclusions of: 1. Substances (like gases) that are miscible in each other mix with each other spontaneously because there is a higher probability of a mixed distribution 2. The most probable mixing pattern produces a uniform distribution of particles throughout the volume they occupy ● Ludwig Boltzmann proposed that the entropy of a system was related to the number of different ways the particles in it could be arranged that was consistent with the system’s total internal energy ○ He linked the entropy (S) of the system to the number of probable arrangements (W) of its particles and particle energies ○ S= kBnW ■ kB= Boltzmann constant (equal to the gas constant divided by avogadro’s number: 23 23 ■ (8.314 J/mol x K) / (6.0221 x 10/mol) = 1.381 x 10 J/K ■ ΔS= Sfinalinitial ■ ΔS= kBnWfinalBnWinitiaBn (Wfinalinitial ● When W finalr than W , initialitive and ΔS is positive ○ Accessible microstates: probable number of the particles in a system at a given temperature ○ Macrostate: the observable system as a whole ○ An increase in the number of accessible microstates leads to an increase in entropy and according to the second law, is spontaneous ○ Boltzmann’s constant, W, incorporates particle motion and position ■ W is the number of energyequivalent microstates ■ We must consider the kinds of motional energy particles have ● Transitional energy: movement of the whole molecule through space ● vibrational energy: periodic motion of atoms within a molecule ● rotational energy: rotation of a molecule about an axis ■ In a solid, crystalline structure, motion is restricted to vibrations only (smallest number of microstates) ■ In a liquid, increased freedom, has vibration motions and rotation motions, more microstates than solids ■ In a gas, molecules are spread out and independent of each other, translational, vibrational, and rotational energy, largest number of microstates ■ To calculate the number of accessible microstates, rearrange the terms in the entropy equation to solve for W and then insert the entropy of 1 mole of the desired molecule ● illustrates the huge number of ways that the energy of a single gas phase atom/molecule can be dispersed ● Statistical definition of entropy based on accessible microstates is conceptually compatible with the thermodynamic, macroscopic view based on energy dispersion ● Third law of thermodynamics: a perfect crystalline solid has zero entropy at absolute zero ○ if the particles of a crystalline solid are perfectly aligned and in their lower possible energy states, the crystal only has one microstate and its entropy is 0. Absolute Entropy and Molecular Structure ● Standard molar entropy (S°): entropy of one mole of a substance in its standard state] ● The change in entropy with temperature is not linear because heating a substance at a higher temperature produces a smaller entropy increase than adding the same quantity of heat to the same substance at a lower temperature ● Factors that affect entropy change: ○ Entropy increases when temperature increases ○ Entropy increases when volume increases ○ Entropy increases when the number of independent particles increases ■ Entropy increases because each change increases the dispersion of the energy of a system’s particles ■ Particles have access to more microstates ■ Can make qualitative predictions of entropy using these three factors ● Standard molar entropies of molecular substances are linked to the masses of their molecules and their molecular structures ● The quantities of energy involved in rotational motion depend on how massive the molecules are and on how their masses are distributed ○ The more massive the molecule the more microstates there are: higher entropy ○ The more spread out, the spacing between its rotational energy states is smaller and more accessible: higher entropy ● Rigidity influences entropy ○ Softer molecules have more access to more microstates: higher entropy Applications of the Second Law ● Universe = system + surroundings ● ΔSuniv sys surr ● When a spontaneous process occurs in an isolated system, ΔSsys greater than 0 ○ process has no impact on its surroundings ○ ΔSunivreater than 0 because ΔSsysgreater than 0 ● A spontaneous process produces an increase in the entropy of the universe ○ positive value of univ ● Second law assumes that a physical or chemical change in a closed or open thermodynamic system can alter the entropies of both the system and its surroundings ● Second law states the process is spontaneous when ΔSuniv eater than 0 ○ provides a thermodynamic requirement for nonspontaneity: a process that produces a decrease in the entropy of the universe will not occur on its own ○ The reverse of any spontaneous process is nonspontaneous ● Exothermic reactions have the capacity to raise the entropy of their surroundings ○ the more energy that flows into the surroundings, or into any collection of particles, the greater energy dispersion among the particles and the greater increase in entropy ○ heating particles that are already hot produces a smaller gain in entropy than heating the same particles at a lower temperature ○ ΔS = (qrev T ■ inverse relationship between entropy gain and temperature ■ q : the reversible energy flow caused by the difference in temperature rev between the system and its surroundings ● Reversible process: process that happens so slowly that an incremental change can be reversed by another tiny change, restoring the original state of the system with no net flow of energy between the system and its surroundings ● Irreversible processes cannot be undone by exactly reversing the change to the system Calculating Entropy Changes ● The entropy of a system is a state function ○ the change in entropy that accompanies a process depends only on the initial and final states of the system, not on the pathway of the process ○ the change in entropy experienced by the system is the difference between its initial and final absolute entropy levels ● ΔS = S S sys finalinitial ● ΔS°rxnn products productsreactants reactants ● ΔS° like ΔH°, is an extensive thermodynamic property that depends on the amounts of rxn substances consumed or produced in a reaction Free Energy ● The entropy change experienced by the surroundings of any chemical thermodynamic system depends on whether the process occurring in the system is exothermic or endothermic ● When energy flows from an exothermic process occurring at constant pressure into a system’s surroundings, the quantity of energy is equal in magnitude by opposite in sign to the enthalpy change of the system ○ q = ΔH surr sys ● When a system undergoes an endothermic process the direction of energy flow is reversed ○ ΔSsurr)surr ○ equal to (ΔHsys T ○ Substitute into ΔSunivation: ■ ΔSuniv sys) /sys ■ Allows us to predict whether or not a process is spontaneous at a particular temperature once we calculate the enthalpy and entropy change of the process ■ downside that spontaneity relies on the value of a parameter (ΔSunivat is impossible to determine directly and that has little physical meaning ● Free energy: measure of the maximum amount of work a thermodynamic system can perform ● Gibbs free energy (G): energy available to do useful work in processes happening at constant temperature and pressure or once the temperatures and pressures of reaction mixtures have returned to their initial values ● Absolute free energy values of substances are often of less interest than the changes in free energy that accompany chemical reactions and to processes ● Gibbs proposed that the change in free energy (ΔG) sys process occurring at constant temperature and pressure is linked directly to that temperature and ΔS univ ○ ΔG = T ΔS univ ○ If ΔG < 0, then ΔS > 0; reaction is spontaneous sys univ ○ if ΔGsys0, then ΔS<univeaction is nonspontaneous; reaction is running in reverse of the process is spontaneous ○ If ΔGsys0, then ΔS univreaction has reached chemical equilibrium ○ T ΔSunivΔS + sys sys ○ ΔG sys sys sys ■ ΔG = ΔH TΔS ● can calculate the change in free energy of a process ● Thermodynamic factors that contribute to a decrease in free energy and making the process spontaneous: ○ The system experiences an increase in entropy (ΔS > 0) ○ The process is exothermic (ΔH < 0) ● How the signs of ΔH and ΔS impact the sign of ΔG and how they define the conditions under which a process is spontaneous ● ΔH°=rxn n ΔH° Σ n ΔH° productsf,products reactsntsf,reactants ● ΔS°rxn n products productsreactants reactants ● Standard free energy of formation (ΔG° ) : change in free energy associated with the f formation of 1 mole of a compound in its standard state from its elements in their standard states ○ ΔG°rxn products f,productsreactsntsf,reactants CHEMICAL KINETICS Cars, Trucks, and Air Quality: ● Photochemical smog: mixture of gases formed in the lower atmosphere when sunlight interacts with compounds produced in internal combustion engines and other pollutants ○ rates of photochemical smog being produced/consumed influence when smog events happen, how intense they are, and how long they persist ○ depends on chemical kinetics ● Chemical kinetics: study of the rates at which reactant and product concentrations change during a chemical reaction ○ provides understanding required to develop catalytic converters and other devices aimed at improving air quality Reaction Rates: ● Reaction rate: how rapidly a reaction occurs ● Related to rates of change in the concentrations of reactants and products over time ● Rate of change in the concentration of a reactant (Δ[reactant] / Δt) has a negative value because reactant concentrations decrease as a reaction proceeds ● The measured rate of any reaction is defined as a positive quantity ○ describes the rate at which reactants form products ○ negative sign is used with Δ[reactant] / Δt values to get an overall positive value for reaction rate ■ Reaction rate = ()rate of2onsumption = (Δ[N2/Δt ■ reaction rates are expressed in units of concentration per time (M/s) ● The coefficients in an equation expressing the relative rates of change of reactants and products are the reciprocals of their coefficients in a balanced chemical equation describing the reaction ● to give all of the terms positive values, the rates representing decreasing reactant concentrations receive minus signs ○ N2g) + 3H2) → 2NH 3g) ○ relative rates: (Δ)/Δt = ⅓([ΔH])/Δt = ½(Δ[NH])/Δt 2 2 3 ● Rate of formation: product is always positive ○ (Δ[NO])/Δt = ([NO] [NO] ) / (Δt Δt final initial finalintial) ● Rate of consumption: reactant is always negative ○ (Δ[N])/Δt = ( [N] ) / (Δt Δt 2 2final2initialfinal intial) ■ final concentration is smaller than initial concentration ● Example: over a 10second time period, the concentration of ammonia increases from 0.133 M to 0.605 M ○ Average rate in ammonia concentration: ■ Δ[NH3/ Δt = (0.6050.133)M / 10 s = 0.0472 M/s ○ Rate of change in concentration of N 2 ■ Δ[N2 / Δt = 1/2Δ[N3 / Δt = ½(0.0472M/s) = 0.0236 M/s ○ rate of reaction is usually the one with the coefficient of 1 ■ Rate = Δ[N2 / Δt = 0.0236 M/s ● to find average rate of formation or rate of change in concentration of reactants find the slope [concentrationfinalentrationinitialfinalinitial ● use the rate where the molecule has the coefficient of 1 ● Difference between the average and instantaneous reaction rates ○ Average is overall ○ instantaneous is at a particular time period ● With a graph, draw a tangent line on the desired instantaneous point to find the instantaneous reaction rate ○ calculate the slope of the line Effect of Concentration on Reaction Rate ● When a reaction is over, no more change occurs in the concentration of any remaining reactant or in the concentration of any product ● The most rapid changes in reactant and product concentrations take place early in most reactions ● We assume that reactions take place as a result of collisions between reactant molecules, then the more reactant molecules there are in a given space, the more collisions per unit time and more opportunities for reactants to turn into products ○ increasing concentrations increases collisions between reactants, leading to the formation of products ● Reaction order: parameter derived from experiments that tells us how reaction rate depends on reactant concentrations ○ Knowing order of reaction gives insight into how the reaction takes place ● Rate law: equation that defines the relation between reactant concentrations and reaction rate ● Rate constant: proportionality constant that relates the rate of a reaction to the concentrations of reactants (expressed as k) ● A + B → C m n ○ Rate= k[A][B] ○ m is reaction order of A, n is reaction order of B ○ determine reaction order by comparing differences in reaction rates to differences in reactant concentrations ● Overall reaction order: sum of the exponents of the concentration terms in the rate law ● Reaction rate depends on the concentration of the reactants, but the rate constant does not change ○ rate constant changes with changing temperature or in the presence of a catalyst ● k = Rate / concentration of reactants multiplied together ● Integrated rate law: ln [X] = kt 0
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