Chemical Equilibrium

1. You will explain that there is a balance of opposing reactions in chemical equilibrium systems 2. You will determine quantitative relationships in simple equilibrium systems

You will define equilibrium and state the criteria that apply to a chemical system in equilibrium; i.e., closed system, constancy of properties, equal rates of forward and reverse reactions

You will identify, write and interpret chemical equations for systems at equilibrium

Collision-Reaction Theory: When a collision of reactants occurs, a reaction takes place.

Assumes:1. that all reactions are spontaneous2. all reactions happen very fast3. quantitative – certain mole ratios have to be met4. stoichiometric – happens to completion each time

Therefore, reactions must happen with greater than 99% efficiency to produce products.

-this is not always the case. Eg. A redox reaction is not always spontaneous. Therefore, even if collision occurs, a reaction may not take place.

However should this theory not also apply to products that come into contact? They should also be able to form reactants (reverse reaction)The truth is; there is often a presence of both reactants and products when no more observable change is taking place (reaction complete- has reached a state of equilibrium). This can be explained with the C-R Theory. If reactants can collide to produce products, products must also be able to collide to produce reactants (reverse reaction). This means there is competition between the forward and reverse reactions.

**In order for competition to take place, the system must be ‘closed’ or ‘isolated’(no reactants or products can escape therefore allowing them to still collide).

A closed system with no more observable change is at equilibrium.

A Dynamic Equilibrium is a balance between the forward and reverse processes. There are 3 types:

1. Phase equilibrium – balance between states of the substance. (s and l)2. Solubility equilibrium – balance between amount dissolved and the

amound undissolved. (s and aq)3. Chemical reaction equilibrium – balance between reactants and products

in a chemical reaction.Pg 680 equil arrows Table 3

The graph of the quantity of each substance against time shows that the rate ofreaction of the reactants decreases as the number of reactant moleculesdecreases, and the rate at which the product changes back to reactants increasesas the number of product molecules increases. These two rates must becomeequal at some point, after which the quantity of each substance present will notchange.

There are two ways to communicate equilibrium:1. % reaction: -only good for specific conditions of the reaction. Any change

in temp, conc, etc. changes the %.

%rxn or %yield = actual yield of product X 100 max. possible yield (stoich)

2. equilibrium constant: depends on the temperature. Works for moderate changes to concentration but not for large scale changes

-more accurate over a larger range than %rxn.-neither provides information on the rate of reaction.

EQUILIBRIUM CONSTANT Kc

You will define Kc to predict the extent of the reaction and write equilibrium-law expressions for given chemical equations, using lowest whole-number coefficients

You will write the equilibrium law expression for a given equation You will use experimental data to calculate equilibrium constants

The equilibrium constant provides a mathematical relationship for a chemical system over a range of concentrations.For an equilibrium constant to be accurate, equilibrium concentrations are to be used in the formula.

aA + bB cC +dD

Kc = [C] c [D] d [A]a [B]b

** pure liquids are not included in the expression (when they are the solvent) eg. WaterHowever, if the equilibrium involves the mixture of what were originally all pure liquids then all are included.

If the reaction is reversed, what happens to the Kc?

**To find the Kc of the reverse reaction, you would flip products for reactants. This is an inverse. Therefore 1/Kc for the forward reaction gives you the Kc for the reverse reaction.

ICE tables: A way to organize data for an equilibrium experiment and calculation.

I-initialC-changeE-equilibrium

Step 1:Calculate the initial concentrations if not already given.Step 2:Set up an ICE table. Record the Initial conc. of each substance in the

reaction. Let the Change in molar conc of one of the reactants be ‘x’ and fill in the corresponding change for the other substances paying special attention to their balancing #. Write and record the expressions for Equilibrium.

Step 3:Write the equilibrium law for the reaction. Substitute the expressions in for concentrations. Solve the expression for ‘x’.

Step 4:Calculate the equilibrium concentration of each gas using your ‘x’ value. Then, if necessary, use the volume to find the amount (moles) of each gas.

Eg. In a 500mL stainless steel reaction vessel at 900 degrees C, carbon monoxide and water vapour react to produce carbon dioxide and hydrogen. Evidence indicates that this reaction establishes an equilibrium with only partial conversion of reactants to products. Initially, 2.00 mol of each reactant is placed in the vessel. Kc for this reaction is 4.20 at this temp. What concentration of each substance will be present at equilibrium?

Reactions with a small Kc:If the constant is very small, then a very small amount of product will be present at

equilibrium. The amount of reactant at equilibrium will be almost the same as the initial amount. Therefore, you can make the approximation that the amount of reactant is the same before and after (when rounding to sig dig rules). Rule: if the equilibrium constant is of the order of 10-4 or smaller ( or 1000 times smaller than the conc of reactant) you can use the initial conc of a reactant in the Kc expression.

Le Chatelier’s Principle: You will predict, qualitatively, using Le Chatelier’s principle, shifts in equilibrium caused

by changes in temperature, pressure, volume, concentration or the addition of a catalyst and describe how these changes affect the equilibrium constant

You will predict variables that can cause a shift in equilibrium You will design an experiment to show equilibrium shifts in concentration You will perform an experiment to test, qualitatively, predictions of equilibrium

concentration shifts You will analyze, qualitatively, the changes in concentrations of reactants and products

after an equilibrium shift You will interpret data from a graph to determine when equilibrium is established and

to determine the cause of a stress on the system

When a chemical system at equilibrium is disturbed, the system appears to react in the direction that opposes the change until a new equilibrium is reached.

Eg. Remove product, shift to form more productThis example is used all the time in industry to produce more of a desired product making the process more efficient and economical.

Ways to shift the equilibrium:1. Concentration changes;

a. Adding more reactant, shifts toward products to reduce reactants Fig3 pg 691

b. Removing product (by physically removing it or by precipitation as a solid), shifts toward product to produce more product Fig 4 pg 692(reverse scenarios are also true but not as commonly used since the desired substance is generally a product)

Does this change the Kc??

No!

2. Temperature Changes;

Heat is treated as a reactant or product as per unit one. The same principles apply to the equilibrium shift as per the concentration changes.Eg. Endothermic: reactants + energy productsExothermic: reactants products + energyWhen cooled, energy is removed from the system and the equilibrium will shift to increase energy (to reactants for endo, products for exo)When heated, energy is added to the system and the equilibrium will shift to reduce energy (to products for endo, reactants for exo)

Does this change the Kc?

Yes!

3. Gas volume changes;

You must consider the total moles of reactant gases and product gases.Eg. 2SO2(g) + O2(g) 2 SO3(g)

If the volume is decreased, the overall pressure is increased, this causes a shift to reduce pressure by reducing the total number of moles of gases – shifting to the side with less moles of gas.If the volume is increased, the overall pressure is decreased - it shifts toward the side of the reaction with more moles of gas.**A system with equal number of moles of gases on both sides of the reaction will be unaffected by changes in volume/pressure. Systems with only solids and liquids are also unaffected by pressure changes. *Adding or removing gaseous substances not involved in the reaction will change the pressure but will not cause any shift to the equilibrium.

Does this change the Kc?

NO!

Catalyst: A catalyst decreases the time to reach equilibrium but does NOT affect the equilibrium position.

REVIEW WATER EQUILIBRIUM/PH AND POH: You will recall the concepts of pH and hydronium ion concentration and pOH and

hydroxide ion concentration, in relation to acids and bases

The equation for the ionization (dissociation) of water is:

H2O (l) H+(aq) + OH-

(aq) or 2H2O (l) H3O+(aq) + OH-

(aq) (Arrhenius definition)

There is very slight conductivity with water which implies that the equilibrium favors the left side.The equilibrium law for the above is:

Kc = [H + (aq)] [OH - (aq)] or (with 2nd eq.) [H2O (l)]

Given the above statement on conductivity, the Kc for water would be very large or small?

Since it is difficult to measure the concentration of a liquid solvent so we change the equilibrium law to an ionization constant for water Kw

Kw= [H+(aq)] [OH-

(aq)] or hydronium ion-[H3O+(aq)]for [H+

(aq)]

The Kw for any solution containing water is always the same, it equals 1.0 x 10-14 at SATP. If the solution is an acid the hydrogen ion concentration is larger than the hydroxide concentration but the product of the two is still Kw. In a neutral solution, the two ions have equal concentration. What would be there concentration?

Kw = x2 x=___________ (neutral solution)

Adding an acid to a neutral solution increases the [H+(aq)], but USING

LECHATELIER’S PRINCIPLE, the equilibrium of the solution will shift to reduce the two ion concentrations until they reach an equilibrium Kw of 1.0x10-14. (What happens when a base is added to a neutral solution?)

pH/pOH

negative logs of the Kw expression:

Kw= [H3O+(aq)] [OH-

(aq)]

for a neutral solution: 14 = 7 + 7 (add instead of multiply)

**The exponent on a number in scientific notation does not count as significant digits, therefore when you take the log, the numbers in front of the decimal (which represent the exponent) do not count as significant. Only the numbers behind the decimal are significant in a pH or pOH measurement. **The number of sign. digits in the concentration become the number of decimal places allowed in the pH; and vice versa-the number of decimal places on the pH become the number of sign. digits in the concentration.

Eg. [OH-(aq)] = 5.5 x 10-3 mol/L (2 digits) What is the pOH and pH

pOH = 2.26 (2 decimals)

STRENGHS OF ACIDS AND BASES

Strong and Weak Acids:pH varies from 1slightly less than 7.Conductivity variesSpeed of reactivity varies

Whether an acid is strong or weak depends on its equilibrium:

Strong Acids eg. nitric, hydrochloric, sulfuric acids-strong acid ionizes QUANTITATIVELY in water to form H+ (or H3O+)

ions.

Weak acids eg. acetic acid- weak electroytes (low conductivity)-react slower-pH closer to 7- only partially ionizes. (usually less than 50% eg. acetic acid 1.3%at SATP)

This means that if you had the same concentration of a strong and a weak acid that the weak acid will have fewer H+ ions than the strong one.

Methods of expressing strength:

1. %ionization:[H3O + (aq)] x100 = %ionization/dissociation[acid]

[H3O+(aq)] = %ionization x [acid]

100

[H3O+(aq)] = [H+

(aq)] which can be obtained from pH, pOH, or hydroxide conc (using Kw).

Restrictions: 1. Using % is only valid for the concentration given and not at any other

concentration since it varies too much.2. Not valid with ions that can be either an acid or a base eg. HCO3

-

2. Ionization Constants You will define Kw , Ka , Kb and use these to determine pH, pOH, [H3O+ ] and [OH– ] of

acidic and basic solutions You will calculate equilibrium constants and concentrations for homogeneous systems

and Brønsted–Lowry acids and bases (excluding buffers) when o concentrations at equilibrium are knowno initial concentrations and one equilibrium concentration are known o the equilibrium constant and one equilibrium concentration are known.

Note: Examples that require the application of the quadratic equation are excluded; however, students may use this method when responding to open-ended questions.

Ionization of weak acids have a state of equilibrium.Use equilibrium law (constant), known as acid ionization constant, Ka in your data bookletAdvantage: Can accurately be used over a range of concentrations to predict [H+].

Eg. find [H+] and pH of following:

0.10 M CH3COOH

0.10 M HCl

0.10 M H3PO4

**the top few weak acids you need to be cautious of the short cut method; It depends on the initial conc whether the short cut will be accurate!!!** weak polyprotic acids: only need to consider the removal of the FIRST hydrogen – the others are insignificant.**However, there is one strong polyprotic acid (sulfuric) in which you could consider the removal of BOTH hydrogen!

BASESArrhenius definition: ionizes to produce OH-

Modified definition; reacts with water to produce OH-

The only ‘strong’ bases are ionic hydroxides that dissociate quantitatively into OH-. The amount of [OH-] in solution determines the strength Ba(OH)2 has a higher pH than NaOH with the same concentration because the barium hydroxide dissociates to form 2 hydroxides for every one compound.

We have all the Ka values available to us in the data booklet but not the Kb’s for the bases. However, the Kb can be obtained using the Ka for the conjugate acid.

Amphiprotic substances – How can you tell if they are more acidic or basic? (pg746)

Bronsted-Lowry Theory of Acids and Bases: You will describe Brønsted–Lowry acids as proton donors and bases as proton acceptors You will identify conjugate pairs and amphiprotic substances

An acid-base reaction involves two substances-one which behaves like an acid and the other like a base. It involves the transfer of a proton (H+)between these two substances. It involves the use of a table that is designed by relative strengths of acids and bases. *An acid is a substance that donates a proton (H+ ) to a base.*A base is a substance that accepts a proton from an acid.An acid must therefore have H in its formula, but any negative ion can be a base (not just OH-)

Eg.

HCl + H2O H3O+ + Cl-

HNO3 + OCl-

B-L Predicting Reactions: You will write Brønsted–Lowry equations, including indicators, and predict whether

reactants or products are favoured for acid-base equilibrium reactions for monoprotic and polyprotic acids and bases

Rules:*change and write the strong acids as H3O+ ( and their negative ion). -because this is how they exist in water. (Ionic aqueous substances are also written as ions)*Include water in your list. *pick the strongest acid (SA) and strongest base (SB).*Equilibrium arrow (reaction arrow) rules:

If the acid is higher than the base, the reaction favors products. If the acid is lower than the base, the reaction favors reactants. If either the acid or base is ‘strong’ then it will be quantitative toward

the products if the other substance is relatively strong. (far apart on the table)

The biggest advantage of this theory over the Arrhenius theory is that it allows you to predict the reactions of different substances as acids or bases with each other and not just with water. To do this you must have a table of relative strengths such as the one in your data booklet.

Eg. 1. Predict the reaction when nitrous acid is added to sodium cyanide. Will it

favor products or reactants?

2. Predict the reaction of sodium hydrogen carbonate with hydrochloric acid. ls the hydrogen carbonate an acid or base in the reaction?

*what might indicate if a significant amount of product is formed???*note that hydrogen carbonate is amphiprotic and yet you predicted its reaction

with something other than water.Amphiprotic substances: capable of acting like an acid or base (usually have a

H and a negative charge – except water) eg.

Monoprotic acids have one hydrogen to donate. Eg.

Polyprotic acids have more than one hydrogen to donate. Eg.

Predicting steps (731)

INDICATORS

READ PG 751-754

-The range given in the data booklet is the range over which the color changes from one to the other. Therefore in that range, the color is a mixture of the two forms or colors.-writing reactions of acids/bases with indicators can demonstrate the color change that occurs.

-only a small amount of indicator should be used since, it too is reacting. You don’t want to alter the concentration of available acid or base by having it react with too much indicator.

Acid/Base titration curves: You will interpret, qualitatively, titration curves of monoprotic and polyprotic acids and

bases for strong acid–weak base and weak acid–strong base combinations, and identify buffering regions

Definitions:1. titration – progressive addition of one reagent to another. You can begin

with the acid and add base or vice versa.2. Equivalence point – the point in a titration where you have chemically

equivalent amounts of reactants and no excess of either. This is where there are stoichiometric quantities. The midpoint of the drop on the pH (titration) curve.

3. Endpoint – visual observed color change occurs because of an indicator. It should match an equivalence point for stoich reasons.

4. Indicator – coloured acids or bases that are present in trace amounts that give a visible indication when a certain pH is reached.

5. Titration curve (pH curve) – a graph with pH on the vertical axis and volume of the titrant (substance in buret) on the horizontal axis.

pH curves

The pH at the middle of the colour change range for bromothymol blue is 6.8, which very closely matches the equivalence point pH; so, a titration analysis endpoint for this reaction, as indicated by bromothymol blue, should give accurate results.

Note that only a strong acid with strong base will have an equivalence point of pH7.

A. -initial addition of titrant to the sample does not produce large changes in pH of the solution. –this is where buffering action occurs. It can be a buffering region if there is a conjugate weak acid/base pair present. Is there a buffering region in the above graph?

B. –very rapid changes in pH for a small addition of titrant. –midpoint of this drop indicates the equivalence point of titration (stoichiometric quantities).–remember, the pH at this point should match the indicator endpoint in order to signal the end of the titration.(otherwise a pH meter can be used and the titration continues until the pH levels off.)-the volume of titrant used to reach endpoint is called the equivalence point volume. This volume is used for stoichiometric calculations since it is the volume where all initial moles of reactants have gone to products (no excess of any reactant).

C. – at this stage only excess titrant is being added until it reaches the titrant’s pH. It is not a buffering region.

General Rule: Only quantitative reactions produce detectable endpoints in a titration. Therefore the # of equivalence points indicates the number of quantitative reactions that have occurred.

Weak acid titrated with strong base:

Is there a buffering region? Where?

POLYPROTIC ACIDS and continuous addition

-any acid with more than 1 proton to release. Eg.-the first H+ comes off the easiest. The remaining are harder to remove.-the fact that more than one hydrogen can be removed means several steps may occur in a reaction. Previously, we only considered the first step when writing the reaction. Now, we will use our further knowledge of acids and bases and determine the number of steps to include with each separate case. Some steps may go to completion while others may not. Only the quantitative steps need to be considered.

Eg. Draw the curve of NaOH continuously added to HOOCCOOH (oxalic acid) with two equivalence points.

Eg.2. Below is a curve of KOH continuously added to sulfuric acid. How many equivalence points are there? How many buffering regions?

pH

Volume of KOH

Eg. 3. Draw the pH curve of NaOH added to carbonic acid with one detectable equivalence point. (both solutions are 0.10mol/L)

POLYBASIC SPECIES and continuous addition

-any base that can accept more than one proton. Eg.

Eg. Draw the curve of HCl continuously added to the carbonate ion.

What would a pH curve look like for hydrochloric acid added to the sulfate ion?

*****Acid/base reactions are quick but not all are stoichiometric.*****Redox reactions are not all quick, but they are all stoichiometric.

Buffers: You will define a buffer as relatively large amounts of a weak acid or base and its

conjugate in equilibrium that maintain a relatively constant pH when small amounts of acid or base are added.

A buffer solution is a chemical system that maintains a nearly constant pH when an acid or base is added-resists a change in pH. They are used to calibrate pH meters and to control the speed and extent of reactions that are pH sensitive. Many examples of buffering exist in biochemistry. For instance, human blood pH must be 7.30 to 7.40 regardless of the food ingested and its pH. If the pH level in human blood were not buffered at 7.35, the acid absorbed from a glass of orange juice would be lethal.One buffer within the cells is the H2PO4

- , HPO42- system. A buffer within the

fluids is H2CO3, HCO3- system.

Eg1: Write the reaction of the citric acid from oranges with carbonic acid-hydrogen carbonate ion system:

*They may not react quantitatively with these substances, but if the reactant is weak to begin with it’s not as serious. .Eg 2: Write the reaction of stomach acid (HCl) with carbonic acid-hydrogen carbonate ion system:

A buffer solution should completely neutralize any STRONG acid or base that is added in small amounts since these are the harmful substances. That means the buffer system can’t be too strong or too weak of an acid/base pair.

To make a good buffer solution, it must consist of fairly equal amounts of unreacted WEAK acid and its conjugate base. (middle of the chart)

When looking at pH curves of acid/base titrations, you may recall the buffering regions. (nearly horizontal areas) these were classified as buffering regions because a small addition made little change to the pH. The point in the buffering

region where you have the best buffering solution is when the volume of the titrant is ½ of the equivalence point volume; or for polysubstances, you are ½ way between endpoints. This is because the solution contains half acid and half conjugate base (or vice versa) making it good for maintaining pH.

* buffering solutions only work for SMALL additions and not continuous additions of an acid or base*use the pH curve and identify the equivalence point, equivalence point volume(s), the buffering region (s), and the volume of each substance required to form a good buffering solution-for which pH.

**Ka for the acid in the buffer system is the approximate [H3O+](and therefore the exponent is the approximate pH) of a buffer solution with equal amounts of acid/base conjugate pair.

You are asked to prepare a good buffer solution for a pH of 7.0 using 50.0 mL of 0.10mol/L NaHSO3 and unlimited 0.10mol/L NaOH. Describe how you would prepare the solution.