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Chemical Kinetics – Example 1

Chemical reactions happen at hugely different rates. Imagine these two examples. In one case, you could have a balloon with hydrogen in a room with oxygen in the environment and you put a match to it and it explodes. That reaction happens in fractions of a second. On the other hand imagine a nail left out in the environment. Due to the oxygen in the air and the humidity, water molecules, maybe help from some salt, that nail would eventually rust. That reaction could take years to happen.

Chemical Kinetics – Example 2A

Determining the units of a rate constant is a tricky process because they’re not always the same. Different reactions are going to have different rate laws. And according to the rate law, that is what’s going to determine the units for the rate constant. This is the way to set about it. On the left hand side of the equation, we have the rate, r. That is set equal to, on the right hand side, the rate constant, k, times a series of concentrations to a series of powers. The rate itself on the left hand side of the equation is always given in moalrity per second. On the right hand side, you can have varying orders of molarity according to what those small numbers are, m, n, can have values of 0,1,2, so on and so forth. So the units of the rate constant must be such that when you look at over all the units on the left hand side and the right hand side, they’re equal to each other. For instance the rate constant could be given in one over seconds, in one over molarity times seconds, and so on and so forth depending on those orders of reaction.

Chemical Kinetics – Example 2B

When you have a chemical reaction, consider plotting the concentration of one given reactant as a function of time. As the reaction goes forward, that concentration is going to decrease. Usually that decrease will be plotted as some form of curve. At any point of that curve, if you take the tangent, that will determine the instantaneous rate of reaction.

Chemical Kinetics – Example 3

The Arrhenius equation gives an expression for k, the rate constant and recognizes three important factors. One of these is the fraction of molecules with an Energy of EA, the activation energy, or greater. The second one is the number of collisions actually occurring per second. And the third one is the fraction of collisions that have the appropriate molecular orientation.

Chemical Kinetics – Example 4

A catalyst is a substance that will speed up a chemical reaction without being consumed in the process Catalysts are very common. Think for instance like the catalytic converter in your automobile, a man made example. Or enzymes in living systems, a natural example,

Chemical Equilibrium – Example 5

Chemical equilibrium is a dynamic state of equilibrium. Think of a similar example corresponding to the picture in your text book. Imagine a bridge with traffic, cars, traveling in both directions. If over a period of time, the same amounts of cars have gone in one direction as in the opposite direction, the cars on both sides of the bridge will not be the same ones anymore, but the quantity will be constant. That would be a case of dynamic equilibrium. Similarly, in a chemical reaction, what’s changing is the amount of atoms and molecules according to the particular chemical reaction. But what will stay constant over a period of time is the overall composition of the system.

Chemical Equilibrium – Example 6

Not all chemical reactions happen in aqueous solutions. Consider, for example, a solution of carbonate ions exposed to an acid, let’s say hydrochloric acid. As a result of adding the hydrochloric acid to the solution a gas will evolve, carbon dioxide. In a closed container at some point there’ll be sufficient carbon dioxide and carbonate ions to establish equilibrium. This would be an example of a heterogeneous equilibrium where one of the components is in the gas phase, the carbon dioxides, and one of the components is in aqueous solution, the carbonate ions.

Chemical Equilibrium – Example 7

When discussing chemical equilibrium, we frequently want to be able to calculate an equilibrium concentration of a reactant or product. For this we use ICE charts. ICE is just an acronym for the three rows in the chart, the initial row, the change row, and the equilibrium row.

To complete the initial row, first we must write the chemical equation. Consider, for instance, the reaction of hydrogen gas plus iodine gas that combine to give two moles of hydrogen iodide. In most problems, initial concentrations will be given for the reactants. We transcribe these in the initial row. Sometimes there is also a product at the beginning of the reaction. Should this be the case, we put the value under the product. In most cases, though, there is no product to start with and therefore we complete that part of the initial row with a zero. Let’s look at the case, for instance, given in your text book, on page 647, for this very reaction. As you’ll see there we start with a one molar concentration of hydrogen gas, a two molar concentration of iodine gas, and zero for the product, hydrogen iodide.

The next row we need to address is the change. The change is simply what happens as a function of time. And frequently, we don’t know the answer for this at the beginning of the problem so we’ll indicate this change with a variable, x. If a substance is disappearing, as will be the case of a reactant, we’ll write minus x. If the substance is a product and therefore being formed as part of the reaction, we’ll write plus x. There’s an addition level of complication. Reactions happen according to a defined stoichiometry. Therefore we must also include the stoichiometric coefficients. In this particular reaction that we’re looking at, we’ll write minus x for hydrogen, as the coefficient was one, minus x for iodine as the coefficient was one, and plus two x for hydrogen iodide as it’s a product and therefore positive, and two moles of hydrogen iodide are formed for each mole of reactant.

This leads us to the third row which is the equilibrium row. To complete this simply add up the values from the first two rows. Therefore, for the equilibrium value for hydrogen gas, you’ll have one molar minus x, for iodine it’ll be two molar minus x, and for hydrogen iodide, it will simply be two x.

Chemical Equilibrium – Example 8

When we consider the effect o f heat on a chemical reaction, according to le Chatelier's principle, it’s important to know if this reaction is exothermic or endothermic. In an exothermic reaction, heat evolves as a product of the reaction. To keep track of what’s happening, you may write reactants go to product plus heat. Therefore, if we increase the temperature in one of these reactions, it’s like increasing the products. Le Chatelier's principle tells us then that the equilibrium will shift towards the reactants.

Acid Base Equilibria – Example 9

Water, H20, is a fascinating molecule. Among other properties, it can act sometimes as a base and sometimes as an acid. Look at the reaction of nitrous acid with water, HNO2 + H20. Its products are nitrate ions, NO2 minus, and hydronium ions, H3O plus. In this case, water has accepted a hydrogen ion from the nitrous acid going from H2O to H3O plus. In this case, water has acted as a base.

Acid Base Equilibria – Example 10

Indicators are used to determine the end point of a titration. If you look at figure 16.7 of your text book, there’s a series of common chemical indicators and the ranges at which the can indicate a change in pH. A very common one is methyl orange. Methyl orange looks red at pHs lower than 4, more acidic that 4, and yellow at pHs greater than this value. Therefore at a pH of 8, the solution would look yellow.

Acid Base Equilibria – Example 11

When comparing the acidity of weak acids, you shouldn’t try and memorize the information. You should look for the equilibrium constants for the acidic reaction in appendix D of the textbook under aqueous equilibrium constants. In particular if you look for ascorbic, citric, and oxalic acid, their KAs are in the order of ten to the minus five, ten to the minus four, and ten to the minus two respectively. Therefore oxalic acid would be the most acidic, citric acid would follow and the least acidic of the three would be ascorbic acid.

Acid Base Equilibria – Example 12

If you’re trying to decide which of a group of bases is the weakest one, first discard all bases that are known as strong bases. Then the remaining, the weak bases, must be compared as to their KB values. KB values can be found either in table D2, directly for the base, or if the information is not available there, look for the conjugate acid and find the KA value in table D1. Remember KA times KB equals KW. that is 10 to the minus fourteen.

Acid Base Equilibria – Example 13

The molecular structure of a compound will determine its acidity. Considering binary acids, that is acids composed of one element and hydrogen, if you compare them along a period in the periodic table the acidity will increase from left to right. This is because the electro negativity also increases from left to right, polarizing the bond between the hydrogen and the remaining atom.