Energy and Rates

Chapter 6 said that a balanced equation does not tell you how fast a reaction goes or whether it happens at all, and that most reactions can run in both directions. This chapter delivers on that. It separates three different questions: does a reaction give out or take in energy, how fast does it go, and does it run both ways until a balance is reached. Each fact is attributed to the Wikipedia page that states it, and my own arithmetic or general knowledge is marked "(check)". Any energy numbers in the exercises are invented teaching values, not real data.

Exothermic and Endothermic

The Exothermic process article defines an exothermic process as a "thermodynamic process that releases energy to its surroundings", usually as heat, though it says light, electricity or sound are also possible. It also describes the IUPAC wording: a reaction whose overall standard enthalpy change is negative. The Endothermic process article gives the opposite: "An endothermic process is a chemical or physical process that absorbs heat from its surroundings", and says an endothermic reaction "generally leads to a decrease in the temperature of the surroundings". So, from the point of view of the surroundings (the flask, the water, your hand), exothermic means it warms up and endothermic means it cools down (my restatement of those two pages (check)).

Notice that acid-base reactions are on the exothermic list. That fits Chapter 8, where the safety note said mixing strong bases with acid spills can cause a violent exothermic reaction. Combustion, from Chapter 6, is the everyday exothermic example, and the Combustion wording used there ("often producing energy in the form of heat or light") says the same thing. Melting and evaporation are endothermic, which is why a change of state (Chapter 1) needs heat to be supplied (my link between the pages and Chapter 1 (check)).

Bonds: Energy to Break, Energy Released When Made

Chapter 5 described covalent bonds as shared electrons. The Bond-dissociation energy article defines bond dissociation energy as "the standard enthalpy change when A−B is cleaved by homolysis to give fragments A and B". The page gives example values with their units: for H−H, 435.780 kJ/mol at 298 K in the main text and 431 kJ/mol in its table (the page gives both, and I did not resolve the difference); for O=O, 498 kJ/mol in its table. That page does not state in so many words that breaking bonds takes in energy; the values are positive enthalpy changes for splitting a bond, which implies it (the page's own limitation). The Collision theory article does say that colliding particles must have enough energy "to break existing bonds and form new ones".

The usual textbook picture is that a reaction breaks bonds in the reactants, which takes in energy, and makes bonds in the products, which gives energy out, and the overall result depends on which is larger. The bond-energy page does not state that sum as a rule in the part I read, so treat it as general knowledge (check). The per-mole unit, kJ/mol, builds on Chapter 7's mole (check). The sign convention used in this chapter is that a negative energy change means energy leaves the system, matching the negative enthalpy change the Exothermic process article describes.

Activation Energy and an Energy Profile

A reaction can give out energy overall and still not start by itself, because the reactants must first get over a barrier. The Activation energy article defines it as "the minimum energy barrier that reactant molecules must overcome to transform into products". It describes an energy profile that plots energy against reaction coordinate, in which "the highest energy position (peak position) represents the transition state". The next table is such a profile written as text. Every number is an invented teaching value, not a measured one.

From these invented levels, the activation energy is 60 − 0 = 60 kJ/mol, and the overall energy change is −250 − 0 = −250 kJ/mol (Python). The two numbers answer different questions: 60 is how high the barrier is, and −250 is where you end up. The reverse reaction, products back to reactants, would face a barrier of 60 − (−250) = 310 kJ/mol (Python), which is my own reading of the picture (check).

What Changes the Rate

The Reaction rate article lists the factors, and I quote what its summary of the page gave. The page's length was not stated.

Applying the page's rule of thumb: a rise of 20 °C is two steps of 10 °C, so the rate constant roughly doubles twice, a factor of 2² = 4 (Python: 4.0). A rise of 30 °C gives 2³ = 8 (Python: 8). This is a rough rule, as the page's word "approximately" says, and it is not exact for every reaction (check).

Collision Theory in Plain Terms

The Collision theory article says that for a reaction, "suitable particles of the reactant hit each other with the correct orientation", and "the successful collisions must have enough energy, also known as activation energy". It also says increasing concentration "brings about more collisions and hence more successful collisions", and that increasing temperature "increases the average kinetic energy of the molecules in a solution". Put together (my plain-language summary (check)):

• Particles must meet, so more crowded particles (higher concentration, or more exposed surface) mean more meetings.
• They must meet the right way round, with enough energy to get over the barrier.
• Heating gives more particles enough energy to react, so more of the collisions succeed.

Most collisions do nothing. The page says only some of the impacts produce a noticeable change.

Catalysts

The Catalysis article defines catalysis as "the increase in rate of a chemical reaction due to an added substance known as a catalyst", and says "Catalysts are not consumed by the reaction and remain unchanged after the reaction." It says the catalyst "provides an alternative reaction mechanism (reaction pathway) having a lower activation energy". The Activation energy article adds that a catalyst "does not change the energies of the original reactants or products", so it lowers the peak of the profile but not the start and end levels. The Catalysis article's examples are enzymes ("protein-based catalysts in metabolism and catabolism") and catalytic converters, "typically composed of platinum and rhodium", which "break down some of the more harmful byproducts of automobile exhaust".

Using the invented profile from earlier, a catalyst that lowered the peak from +60 to +40 would cut the forward barrier from 60 to 40 kJ/mol, and leave the overall change at −250 kJ/mol (Python). It would give no extra energy out. That is an invented example of what the pages describe, not a measured one. Because the catalyst is not used up, it does not appear as a reactant or product in the balanced equation of Chapter 6; the Chemical equation article, as Chapter 6 noted, shows catalysts only above or below the arrow.

Reversible Reactions and Equilibrium: A First Look

Chapter 6 reported that most reactions "can and do run in both directions". The Chemical equilibrium article describes the state in which reactants and products are both present, with no net change in their concentrations: "This state results when the forward reaction proceeds at the same rate as the reverse reaction", known as dynamic equilibrium. It says arrows pointing both ways in an equation indicate equilibrium, and credits Jacobus van 't Hoff with proposing the double arrow, ⇌, in 1884. The word "dynamic" matters: both reactions are still going on, only their rates match (my plain restatement (check)).

Chapter 8's weak acids are the example the end of that chapter pointed to: a weak acid only partly dissociates, so both the acid and its conjugate base are present at equilibrium (Acid article, as quoted in Chapter 8).

The Le Chatelier's principle article says that if the equilibrium of a system is disturbed by a change in one or more of the determining factors, the system adjusts to counter the change. The words I could see on the page were only part of the statement, so this is a simplified version (check). From the page's summary:

• Concentration. Adding carbon monoxide to the methanol synthesis, CO + 2 H₂ ⇌ CH₃OH, pushes the reaction toward more methanol, and removing a product can drive the reaction forward.
• Temperature. The direction depends on whether the reaction is exothermic or endothermic. The page says for an exothermic reaction "an increase in temperature decreases the equilibrium constant".
• Example. The reaction 2NO₂ ⇌ N₂O₄ is exothermic. Cooling the flask shifts it toward N₂O₄ and the flask becomes colourless; heating it shifts the reaction back toward NO₂ and the flask turns reddish-brown.
• Pressure. The page says the equilibrium concentrations "do not directly depend on the total pressure"; pressure matters mainly through partial pressures, and has no effect when the gas moles are equal on both sides.

Equilibrium constants and calculations are beyond this course. Equilibrium also is not "equal amounts": the page says only that concentrations stop changing (my reading (check)).

Where to Read More

• Encyclopaedias. The Wikipedia articles above are the sources used here; the Activation energy article covers the Arrhenius relation, which this chapter leaves out.
• A school textbook. Look for enthalpy changes, energy-profile diagrams, rate laws and equilibrium constants.
• Practical. Only with supervision and a proper kit: compare how fast a safe, dilute reaction goes warm and cold (my suggestion (check)).

Exercises