Chapter 17 Review

Equilibrium

applies to the extent of a reaction, the concentration of reactant and product present after an unlimited time, when no further change occurs. In equilibrium, no further net change occurs because forward and reverse reactions are in balance. Essentially, given sufficient time, the concentrations of the reactants and products attain certain values that no longer change. This is because, all reactions are reversible and reach a state of equilibrium. At equilibrium: rate(fwd)=rate(rev) at a particular temperature.

Kinetics

applies to the speed (rate) of a reaction, the concentration of reactant that disappears (or of product that appears) per unit of time.

Law of Chemical Equilibrium or law of mass action

at a given temperature, a chemical system reaches a state in which a particular ratio of reactant and product concentrations has a constant value. for a particular system and temperature, the same equilibrium state is attained regardless of starting concentrations.

Qc = Kc

equilibrium; no net change

Simplifying Assumption

if a reaction has a relatively small K and a relatively large initial reactant concentration, the concentration change (x) can often be neglected. Predict assumption IF: 1. [A]initial/Kc > 400 assumption is justified. 2. [A]initial/Kc < 400 assumption is NOT justified.

Qc < Kc

more product forms because shifting to right to reach equilibrium.

Qc > Kc

more reactant forms because shifting to left to reach equilibrium.

Form of Q for a reaction with coefficients multiplied by a common factor:

multiplying all the coefficients of the equation by some factor also changes the form of Q. The same is true for the equilibrium constant, raising the multiplied value to the power. If all the coefficients of the balanced equation are multiplied by some factor, that factor becomes the exponent for relating the reaction quotients and the equilibrium constants.

At equilibrium NOTE:

the exponent of the RT term equals the change in the amount (mol) of gas (delta mol gas), from the balanced equation. Kp = Kc(RT)^change in amount (mol) gas. (if there is no change in moles from the balanced equation, the exponent is 0, dropping the RT term, and Kp = Kc).

Form of Q for a forward and a reverse reaction

the form of a reaction quotient depends on the direction in which the balanced equation is written. A reaction quotient (or equilibrium constant) for a forward reaction is the reciprocal of the reaction quotient (or equilibrium constant) for the reverse reaction. A high Kc value means the reaction is going far to the right, meaning lots of products are formed at equilibrium. A low Kc value (which would be the reverse reaction of the same equation) does not form a lot of products.

Magnitude of K

the magnitude of K is an indication of how far a reaction proceeds toward product at a given temperature. 1. a SMALL K means a reaction yields little product before reaching equilibrium. This can be used to assume there is "no reaction". (ex. K = 1.0*10^-30) 2. a LARGE K means when a reaction reaches equilibrium, there is very little reactant remaining and a lot of product is formed. This can be used to assume a reaction "goes to completion". (ex. K = 2.2*10^22) 3. an INTERMERDIATE K means at equilibrium, significant amounts of product and reactant are present. (ex. K = 5)

equilibrium position

the specific equilibrium concentrations (or pressures).

Reaction Quotient (Q):

this is the particular ratio of concentration terms that we write for a given reaction. It is essentially the same as K, but Q is not at equilibrium, it is just product concentrations over reactant concentrations from the chemical formula. K is unitless. At equilibrium: Q = K The reaction quotient must be written directly from the balanced equation (UNLIKE THE RATE LAW FOR AN OVERALL REACTION).

Equilibrium problems

typically use quantities (concentrations or pressures) of reactants and products to find K, or we use K to find quantities.

Reaction Quotient when dealing with pure liquids and solids:

we eliminate the terms for pure liquids and pure solids in the expressions for Q and K because we are concerned only with concentrations that change as they approach equilibrium, and the concentrations, like the densities, of pure solids and pure liquids are constant.

Le Chatelier's Principle:

when a chemical system at equilibrium is disturbed, it retains equilibrium by undergoing a new reaction that reduces the effect of the disturbance.

If you want the reaction to proceed towards the side with fewer moles of gas:

you need to decrease the volume, increase the pressure.

If you want the reaction to proceed towards the side with more moles of gas:

you need to increase the volume, decrease the pressure.

The van't hoff equation: the effect of T on K for an EXOthermic reaction:

(a negative value, adding heat to the product side-as temperature increases, the value of Kc decreases). An increase in temperature for an exothermic reaction results in a shift to the left and a decrease in K.

The van't hoff equation: the effect of T on K for an ENDOthermic reaction:

(a positive value, adding heat to the reactant side- as temperature increases, the value of Kc increase). Everything to the left side of the equation is <0, and everything on the right side of the equation is >0. SO; and increase in temperature for an endothermic reaction results in a shift to the right and an increase in K.

If volume becomes larger:

(smaller pressure), the reaction shifts so that the total number of gas moles increases.

Three ways pressure change can occur:

1. Changing the concentration of a gaseous component. 2. Adding an inert gas (one that does not take part in the reaction). As long as the volume of the system is constant, adding an inert gas has no effect on the equilibrium position because all concentrations, and thus partial pressures, remain the same. 3. Changing the volume of the reaction vessel. This change can cause a large shift in equilibrium position only for reactions in which the number of moles of gas changes.

There are two types of Equilibrium problems:

1. Knowing the equilibrium quantities and solving for K. 2. Knowing K and the initial quantities and solving for the equilibrium quantities.

Three possibilities for the relative sizes of Q and K:

1. Q < K: If Q is smaller than K, there are more reactants. In order for Q = K, the reactants must decrease and the products must increase. The reaction at this point is not at equilibrium, and will progress to the right, toward products, until Q = K. 2. Q > K: If Q is larger than K, there are more products than reactants. In order for Q = K, the reactants must increase and the products must decrease. The reaction as this point is not at equilibrium, and will shift to the left, until Q = K. 3. Q = K: This situation occurs when the reactant and product terms equal their equilibrium values. No further net change takes place.

Preliminary steps in setting up to solve an equilibrium problem:

1. Write a balanaced equation. 2. Write the reaction quotient, Q. 3. Convert all amounts into the correct units (M or atm).

How is a system disturbed?

At equilibrium, Q = K, the system is disturbed when a change in conditions forces it temporarily out of equilibrium. There are three common disturbances: 1. Change in concentration. 2. Change in pressure (caused by a change in volume). 3. Change in temperature.

Reactions in which the number of moles of gas DOES NOT change (Le Chatelier's):

Because a change in volume has the same effect on the numerator (products) and denominator (reactants), there is NO effect on the equilibrium position if number of gas moles are not changing.

Reactions in which the number of moles of gas change (Le Chatelier's):

Boyle's Law: pressure and volume are inversely proportional (if the volume is halved, the pressure is doubled). To reduce this disturbance, the system responds by reducing the number of gas molecules in order to reduce the pressure. The only way of doing that is through a net reaction toward the side with fewer moles of gas. (RE: when volume is halved, concentration and pressure are doubled). 1. If the volume becomes smaller (higher pressure), the reaction shifts so that the total number of gas molecules decreases. 2. If the volume becomes larger (smaller pressure), the reaction shifts so that the total number of gas moles increases.

The addition of a catalyst (Le Chatelier):

No effect on the equilibrium POSITION, but it does shorted the amount of time it take to REACH equilibrium.

Removing a reactant, decreasing concentration of reactants (Le Chatelier's):

The system will replace the removed reactant by consuming more of the product and proceeding towards the reactants, or to the left. (in terms of Qc, when the amount of reactant decreases, and product increases, Qc increases). Overall: the equilibrium position shifts to the left when a component on the left is removed.

Reaction Quotient based on concentrations:

This is the most common form of Q, showing reactant and product terms as molar concentrations. It is denoted as Qc. (Equilibrium constant based on concentrations is denoted as Kc). Qc: a ratio of product concentrations multiplied together and divided by reactant concentrations multiplied together, with each term raised to the power of its balancing coefficient. This requires 2 steps. 1. Balance the overall equation. 2. Arrange the terms and exponents.

Expressing equilibrium and reaction quotient with pressure terms (Kp)

This is used when a reaction involves gas, because it is easier to measure gas pressure instead of gas concentration. The reaction quotient is also expressed in terms of partial pressures instead of concentrations for reactions involving gas. The ideal gas law allows us to relate pressure (P) to concentration (moles/V); PV = nRT. - at a constant temperature, pressure is directly proportional to molar concentration. Kp: the equilibrium constant based on pressures (most cases this value is different than Kc). However, if you know one, the change in amount (mol) of gas, form the balanced equation allows you to calculate the other (the change in amount (mol) of gas is found from final mol minus initial mol).

The effect of change in concentration (Le Chatelier's):

When a system at equilibrium is disturbed by a change in concentration of one of the components, it reacts in the direction that reduces the change. 1. If the concentration of reactants increases, the system reacts to consume some of it. 2. If the concentration of reactions decreases, the system reacts to produce some of it.

Exothermic Reaction:

When heat is added to the product side of a chemical reaction equation. It is when heat is lost from the system. The system feels hot. An exothermic reaction is always less than 0.

Shift to the left:

a net reaction that converts product to reactant until equilibrium is reached.

Shift to the right:

a net reaction that converts reactants to product until equilibrium is reached.

Heat:

a rise in temperature occurs when heat is added to the system. A drop in temperature occurs when heat is removed from the system. The system shifts to reduce the effect either change. A temperature increase (adding heat) favors the endothermic (heat-absorbing) direction, and a temperature decrease (removing heat) favors the exothermic (heat-releasing) direction. Always flows from hot to cold substances. Think ice cube in hand, your hand releases heat which melts the ice cube.

If volume becomes smaller:

(higher pressure), the reaction shifts so that the total number of gas molecules decreases.

Gas constant R:

8.314 J/mol

The effect of a change in pressure (Volume) (Le Chatelier's):

A change in pressure has its largest effect on equilibrium systems containing gas. A change in pressure has a negligible effect on liquids and solids because they are nearly incompressible.

A positive delta H* rxn:

Endothermic: a temperature increase will increase the value of Kc. Adding heat to the left of the equation, will shift the equation to the right.

A negative delta H* rxn:

Exothermic: a temperature increase will decrease the value of Kc. Adding heat to the right side of the equation, will shift the equation to the left.

Reaction Table

I.C.E. Table. This table shows the balanced equation and the; (I): initial quantities (concentrations and pressures) of reactants and products; (C): changes in these quantities during the reaction; (E): the equilibrium quantities. The final quantities are those at equilibrium.

Form of Q for an overall reaction:

If an overall reaction is the sum of two or more reactions, the overall reaction quotient (or equilibrium constant) is the product of the reaction quotients (or equilibrium constants) for the steps. 1. Q(overall) = Q1 * Q2 * Q3 * ... 2. K(overall) = K1 * K2 * K3 * ...

Adding or removing product (Le Chatelier's):

If you increase the amount of product, more reactants must be made to reach equilibrium, so the reaction shifts to the left until equilibrium is met (in terms of Qc: a higher product and lower reactant makes for a a higher Qc value. If Qc > Kc, the reaction shifts left). If you decrease the amount of product, more reactants are actually being made, therefore the reaction must shift to the right to make more products (in terms of Qc: a lower amount of product and higher amount of reactant makes for a smaller Qc value. If Qc < Kc, the reaction shifts right).

Equilibrium Constant (K):

K = k(fwd)/k(rev) K is a number equal to a particular ratio of equilibrium concentrations of products and reactants at a particular temperature. K is a special value of Q that occurs when the reactant and product concentrations have their equilibrium values. K is unitless.

Comparing Q and K to determine reaction direction

K is a value that remains constant at a particular temperature, so if you know K at a given temperature, and you determine Q, you are able to determine, at which point Q is monitored in the reaction, what direction the reaction is progressing. More product makes Q larger, and more reactant makes Q smaller. A larger Q would shift the reaction to the left, and a smaller Q would shift the reaction to the right (to form more product).

The effect of a change in temperature (Le Chatelier's):

Of the three types of disturbances (concentration, pressure, and temperature), only temperature changes can alter the value of K (this is seen by focusing on the sign of enthalpy change (H) of reaction).

Adding reactant, increasing concentration of reactants (Le Chatelier's):

The system will consume the added reactant, shifting toward the product, or the right. (in terms of Qc, when the amount of reactant increases, Qc decreases). Overall: the equilibrium position shifts to the right when a component on the left is added.

What is a net reaction?

This refers to a shift in the equilibrium position to the right or the left. Concentrations (or pressures) change in a way that reduces the effect of the change in conditions, and the system attains a new equilibrium position.

Endothermic Reaction:

When heat is added to the reactant side of a chemical reaction equation. It is when heat is gained by the system. This system feels cold. An endothermic reaction is always greater than 0. The addition of heat on the Q value: when adding heat to the reactant side of the equation, you are increasing the denominator of the Q value, so the new equilibrium (Kc) value is smaller than the initial Kc value. The release of heat on the Q value (decrease of temperature): this is the addition of heat on the product side of the equation, which increases the numerator of Q and decreases the denominator, thus increasing the value of Q, so the new equilibrium (Kc) value is greater than the initial Kc value.

Qc:

When the denominator is greater than the numerator: smaller K value. When the numerator is greater than the denominator: larger K value.

ICE tables starting with a mixture of products and reactants

You must first compare Q and K values to find the direction the reaction is proceeding to reach equilibrium.

In terms of K (Le Chatelier's pressure change with no mole change)

a change in volume is, in effect, a change in concentration. A change in pressure due to a chance in concentration (volume change) does NOT alter Kc (because there is no mole change).

Quantitative view of equilibrium

a constant ratio of constants. The ratio of constants creates the equilibrium constant.

Kinetics and equilibrium

are both distinct aspects of a reaction system, and the rate and extent of a reaction are not necessarily related.

van't Hoff equation

shows quantitatively how the equilibrium constant is affected by changes of temperature. equation: lnK2/K1 = -H*rxn/R (1/T2 - 1/T1). R= universal gas constant (8.314 J/mol*K) This is slightly similar to the Arrhenius equation. Do not get them mixed up.

Value of Kc with change in concentration:

the Kc value DOES NOT CHANGE with a change in concentration at a given temperature.