Problem #18: Factors Affecting Change of Matter [Solved]

Daniel Iyamuremye -

Problem 17 was about the factors that affect the change process of matter, can you expand a little more about how this works?

Solution:

To understand how the different factors that drive the changes of matter work, let’s start by a comparison. In physics, potential energy (PE), is the energy of an object due to its position.

We all know that rivers flow naturally from high altitude to low altitude. At high altitude, rivers have high potential energy, PE1, at low altitude rivers have a low potential energy, PE2. Therefore rivers at high altitude, with high potential energy, PE1 > PE2, flow naturally from high potential energy position to low potential energy position at low altitude. This is a natural process that goes by itself.

Contrary to the situation where a river flows naturally from high altitude to low altitude, you may want to bring water from a source at low altitude to a point higher in altitude, PE1 < PE2; in that situation you need electrical energy or another kind of force to help bring water from low altitude to the high altitude point. The flowing is not natural, it must be assisted by an external source of energy.

In this example of a river, the potential energy is a key factor that determine if it flows naturally from a point to another or must be assisted.

The same happens to many natural phenomena, and more specifically to phenomena related to changes of matter such as physical changes and chemical reactions.

Changing liquid water into vapor needs external energy (evaporation and vaporization), because liquid water has lower energy than vapor. On the contrary, changing vapor into liquid water (condensation, formation of rain) doesn’t need external energy.

When dealing with changes of matter, keep that comparison in mind.

In a chemical reaction we are dealing with reactants that transform into products:

A + B → C + D (1)

Chemical reactions consist into the breaking of chemical bonds in the reactants, this requires energy, and the formation of new chemical bonds in the products, and this latter stage releases energy.

  • How do those two processes compare each other?
  • Another factor to consider, as said in problem 17, is the entropy, does the entropy increase or decrease during the change?

Let’s answer to those two questions:

1. System: what’s a system? You can consider any object subjected to your observation or study as a system.

Examples:
When you go to see a doctor, they first of all take your temperature, weight, blood pressure, etc.In front of the doctor, your body is a system. The doctor is assessing your body system.

When you are cooking your supper, you put the food and the ingredients in a casserole and start cooking by supplying heat. The casserole and its content constitute a system, and the system exchanges heat with its surroundings.

There are three types of systems:

  • Open system: an open system is a system that can exchange matter and heat with its surroundings.
    Example: Cooking in an open casserole or carrying out an experiment in an open beaker.
  • Closed system: a closed system is a system that can exchange heat but not matter with its surroundings.
    Example: An under pressure cooking casserole and its content. A closed glass or plastic bottle and its content.
  • Isolated system: an isolated system is a system that cannot exchange neither matter nor energy with its surroundings.
    Example: A thermos flask and its content.

2. Internal Energy of a System (U): one of the factors that drives chemical changes is the Internal Energy of the System, U. The internal energy is the total of kinetic energies of the atoms and molecules of which a system consists and the potential energies associated with their mutual interaction. One major component of U is the Potential Energy stored in chemical bonds and static interactions such as: covalent bonds, electrostatic forces or ionic bonds, and nuclear forces.

U = H – pV

where H = enthalpy or heat exchanged between the system and its surroundings,     p = pressure, V = volume.

At constant pressure: ΔU = ΔH – pΔV = ΔH – W

where pΔV = W = work done by the system on the surroundings.

The units of U are: Joules

3.  Enthalpy of a System (H): enthalpy generally represents the heat exchange (taken in or taken out the system) during the chemical reaction. This property varies with conditions of the experiment: temperature, pressure, concentration of solutions, etc... This thermodynamic property is represented by the mathematical formula, H = U + pV

In other words, it is the sum of internal energy and the product of the pressure and volume.

The units of H are: Joules/mol.

During a chemical reaction at constant pressure, there is a change in enthalpy, represented by the mathematical formula:

ΔH = ΔU  +  pΔV = ΔU + W

As you can notice, the change in Enthalpy, ΔH, is equal to the change in Internal Energy, ΔU (can be negative or positive), plus the work, W, that the system can do to its surroundings or that the surroundings can do to the system.

When ΔH < 0, negative value, the reaction is Exothermic; it releases heat in its surroundings.

If ΔH > 0, positive value, the reaction is Endothermic; it takes heat from its surroundings.

Exothermic reactions (ΔH < 0) are naturally more favored.

(See Fig. below: energy absorbed = positive ΔH, energy released = negative ΔH)

Enthalpy and Activation Energy of Chemical reactions (Source: researchgate.net)

4. Heat of reaction (Q) or (q): the heat of the reaction is the amount of heat that must be added or removed during a chemical reaction in order to keep all the substances present at the same temperature. If the pressure in the vessel containing the reacting system is kept at constant value, which is the case in many chemical reactions, then Q and ΔH are equal.

Although Q and ΔH are often used alternatively, it is wrong to take them as identical concepts; it is only when a reaction is carried out at constant pressure that Q and ΔH are equal.

5. Entropy (S): as seen previously, in a wider sense entropy can be interpreted as a measure of disorder, randomness; the higher the entropy the greater the disorder. S = Q/T

The units of S: Joules/K

During a chemical change, the change in entropy is represented by the mathematical formula:

ΔS = ΔQ/T, where ΔQ = heat transferred to the system, T = temperature, at which the reaction takes place

Reactions with positive change of entropy, ΔS > 0, are naturally more favored (as the example of the students on the school playground given in the previous Problem).

6. Temperature: for any reaction, the temperature at which the reaction is taking place must be indicated. When the temperature is not indicated, it is assumed to be at room temperature. When calculations are involved, the Temperature Unit must be Kelvin (K).

7. Gibbs Energy (G): Gibbs Energy, also called Gibbs Free Energy or simply Free Energy, is the combination of enthalpy and entropy into a single value. Gibbs free energy is a measure of a System’s ability to do work. It is represented by the mathematical formula:

G = H – TS

G = U + pV – TS

And the change in Gibbs Free Energy during a chemical reaction is represented by the mathematical formula: ΔG = ΔH – TΔS    (2)

The change in Free Energy, ΔG, is an indicator if the chemical reaction will occur or not in the given conditions.

If ΔG < 0, negative value, the reaction is qualified thermodynamically Spontaneous, in other words, the reaction would take place, going on by itself. It can be initiated but after initiation, it continues by itself. Spontaneous reactions are also called Exergonic. Notice that in those conditions: ΔH < TΔS

Fuel combustion is a spontaneous reaction, but it must be initiated by a match flame or a spark; once initiated it continues by itself.

Examples of a spontaneous reaction:

  1. 2H2(g) + O2(g) → 2H2O(g) + Release of Heat (negative ΔH), Exothermic reaction

  2. NH4NO3(s) + H2O(l0 → NH4+(aq) + NO3-(aq) Absorption of heat (positive ΔH), This is an endothermic reaction, but spontaneous due to its high entropy change so that ΔH < TΔS

  3. Ice (solid H2O)at room temperature changes to H2O(l) Endothermic process, but spontaneous due to its high entropy change so that ΔH < TΔS

N.B: Notice that all exothermic reactions/processes are spontaneous; but all spontaneous reactions/processes are not exothermic. Endothermic reactions/processes where ΔH < TΔS as in examples (2) and (3) are also spontaneous.

If ΔG > 0, positive value, the reaction is qualified thermodynamically Non-spontaneous.  A non-spontaneous reaction is a reaction which can neither take place by itself or by initiation. It needs a continuous supply of heat or other types of energy all along the reaction. They are also called Endergonic. Notice that for those reactions, ΔH > TΔS.

Example of a non-spontaneous reactions:

Decomposition of calcium carbonate (Limestone) by heat:
CaCO3(s) + Heat(2000oC) → CaO(s) + CO2(g) Endothermic and non-spontaneous.

Decomposition of water into hydrogen and oxygen by electricity (electrolysis):
2H2O(l) + electricity → 2H2(g) + O2(g) Electrolysis, non-spontaneous.

N.B: Remember that all endothermic reactions are not non-spontaneous, as in the examples above.

If ΔG = O, this indicates that the reaction has reached equilibrium, where the rate of formation of the products from the reactants (forward reaction), and the rate of formation of the reactants from the products (backward reaction) are equal.

8. Activation Energy (Ea): there is a barrier of energy that needs to be overcome for a reaction to take place; this barrier is called the “Activation Energy”. The activation energy is the minimum energy needed to start the reaction (see previous figure or below).

9. Catalyst: a catalyst is a substance that increases the rate (speed) of the reaction without itself undergoing any permanent chemical change. Its role consists in lowering the Activation Energy. Not all chemical reactions need a catalyst (see Fig. below).

Action of Catalyst (Source: thefactfactor.com)

Conclusion:

All the factors that contribute in driving a chemical reaction can be combined in one factor, the change in Free Energy, ΔG, of a reaction or system. A negative value indicates that the reaction or change is spontaneous, i.e. can occur naturally and sustain itself. A positive value indicates that the reaction or change needs a continuous supply of external energy to take place. A zero ΔG indicates that the system is at equilibrium.

Use & Application:

It is good to know all those factors that affect a chemical reaction and how those factors operate, for theoretical purpose and advancement of knowledge. But are these worthy for any practical purpose? Yes.

That knowledge is very important not only for theoretical purposes and advancement of knowledge, but also for practical purposes, particularly in the chemical industry. No need to talk about the importance of the chemical industry in our modern society, we live with its products everywhere around us and in almost all of our activities.

For any industrialist to venture into the production of any chemical product, s/he must first  find  of all answers the following questions:

  • Is the envisaged production feasible?
  • If yes, will the chemical reaction be spontaneous, i.e. no need of external energy, or nonspontaneous, i.e. will need external energy?
  • If the reaction will need external energy, what kind of energy and what’s the cost of that energy?
  • What is the best available technology to manufacture the product?
  • Is the reaction a catalyzed or uncatalyzed one?

It’s the answers to all these questions and others that will determine the decision to invest in the venture. And as you can see, many of those questions are related to the factors mentioned in the answer to the problem above.