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CHEM 14B Lecture Notes - Lecture 1: Ammonia, Equilibrium Constant, Partial Pressure

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turquoisemoose666

1 Feb 2021

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At molecular level: r forming p & reverse reaction p forming r (but at an equal rate) Under same conditions, reaction at equilibrium has same. Equilibrium constant (kc): ratio of the concentration of the product to the concentration of the reactant. Tells the relative stability of the products and reactants. Reaction at equilibrium aa + bb cc + dd. A, b, c, d: stoichiometric coefficients of blanched equation. Example: calculate kc for n2 mixed with h2 at 500 c to produce ammonia. At equilibrium: n2 = 0. 305m, h2 = 0. 324m, nh3 = Reverse rxn: 2nh3 n2 + 3h2. Kc = ([n2][h2]^3)/[nh3]^2 = 1/61. 0 = 0. 016. All r and p in same phases: homogeneous equilibrium. One of more r or p in different phase: heterogeneous equilibrium. Molar concentration of a pure substance (solid or liquid) does not change in a reaction. Solids and liquids not included in k expression. Example: ca(oh) ca+2 + 2 oh-

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Related Questions

Â± Relating Different Forms of the Equilibrium Constant

For chemical reactions where all reactants and products are in the gas phase the amount of each gas in the vessel can be expressed either as partial pressures or as concentrations. As such the equilibrium constant for a gas phase reaction can also be expressed in terms of concentrations or pressures. For the general reaction,

aA(g)+bB(g)âcC(g)+dD(g)

Kp=(PC)c(PD)d(PA)a(PB)b and Kc=[C]c[D]d[A]a[B]b.

It is possible to interconvert between Kp and Kc using

Kp=Kc(RT)În

where R=0.08314 L bar molâ1 Kâ1 and În is the difference in stoichiometric coefficients between gaseous products and gaseous reactants.

If we assume that the gaseous reactant and products are âidealâ and the partial pressures of the reactants and products were expressed in bar, then the thermodynamic equilibrium constant, K is the magnitude (or numerical value) of Kp.

Part A

For the reaction

2CH4(g)âC2H2(g)+3H2(g)

Kc=0.160 mol2 Lâ2 at 1621 âC. What is Kp for the reaction at this temperature?

Enter your answer numerically.

Part B

For the reaction

N2(g)+3H2(g)â2NH3(g)

Kp=3.80Ã10â3 barâ2 at 265 âC. What is Kc for the reaction at this temperature?

Enter your answer numerically.

ENT ives To determine equilibrium concentrations for species in solution . Examine temperature effects on reaction equilibrium Investigate the shift in equilibrium from different stresses on a system kills . Use colorimeter to determine unknown solution concentrations ced .Preparing temperature baths to change reaction conditions tionEquilibrium is defined as a state in which the concentrations of both the reactants and products have no tendency to change with time. We can restate this by saying that the rate of the forward reaction is equal to the rate of the reverse reaction. Recall that with rate we are simply defining how fast a reaction is proceeding, a speed at which the reaction proceeds. Take note that when writing an equilibrium expression, pure liquids (denoted by l or 2) and pure solids (denoted by s) are omitted. Typically we express the equilibrium by K or Kp where c and p denote either concentration or partial pressure respectively. We refer to this value of K, or Kp as an equilibrium constant For any reaction, Kc can be written as the product of the concentration of products, raised to their stoichiometric coefficients, divided by the product of the concentration of the reactants raised to their stoichiometric coefficients. See Equation (5-1) below for an example: aA(aq) + bB(aq-cC(aq) + dD(aq) (5-1) A] [B]b As mentioned before, K, or K, are equilibrium constants meaning that they will have the same value unless temperature is changed. This brings in two important ideas 1) We can perturb the equilibrium concentrations of a reaction, say [A] for instance which will cause the other concentrations to shift so that K for the reaction remains constant. 2) If we change temperature, we will either increase or decrease our value of K depending on the specific reaction we are examining. Determining K, and Equilibrium Concentrations During lecture we often discuss the idea of using ICE tables as a method for organizing and calculating concentrations of reactions as they proceed towards equilibrium. In a typical ICE table you need to know initial concentrations and the magnitude of K to determine the equilibrium concentrations. Alternatively, if you know the equilibrium concentrations for a reaction you can calculate K. For this experiment we will be examining a complex ion formation, when iron(III) ion