# Rate of Reaction | Basic Concept and Formulas

Hi, guys. Now, we’re gonna discuss about ‘Rate of Reaction’. I’m sure you are familiar with the term of ‘rate’. Yeah, we usually use that term in physics. However, do you know that in chemistry, we also use that term? Yep, we use the ‘rate of reaction’ to define the loss of reactants or the increase of products in chemical reaction over time.

Anyway, how can reactants change into product?

Now, we introduce you another term, collision theory. The theory says that, to undergo a reaction, particles need to come into contact with one another or in the other words collide with sufficient energy and correct orientation. Effective collision will change reactants into products. If particles do not have enough energy, they won’t change into reactant even though there is a collision. That is called ineffective collision.

So, how much energy do the particles need in order to collide effectively? There is another term, known as activation energy, symbolised by Ea which is the minimum energy required for the particles to collide effectively so the reactant will turn into product.

Figure below shows activation energy for exothermic and endothermic reaction,

exothermic

endothermic

see more about exothermic and endothermic reaction.

Based on collision theory, there are several factors that effect rate of reaction. What are they? Let’s check one by one!

## The Effect of Concentration on Rate of Reaction

A higher concentration implies a greater quantity of solute particles dissolved within a specific volume of solvent. In reactions that involve solutions, more concentrated reactants exhibit a swifter reaction rate due to the increased frequency of collisions between reacting particles. If we see the figure above, the left picture has more particles means the solution is more concentrated than the right one. Thus, the collision happened more frequent for the solution on the left. In conclusion, the increase of concentration will increase rate of reaction because the frequency of collision increases.

## The Effect of Temperature on Rate of Reaction

When the temperature of a reaction mixture is elevated, it leads to an increase in the average kinetic energy of the particles. This results in faster movement of particles in both solutions and gases, leading to a higher frequency of collisions. The likelihood of successful collisions, meaning those resulting in a reaction, rises due to the increase in particles surpassing the activation energy threshold. This aspect holds greater significance in influencing the outcome.

## Catalysts Effect

Catalysts enhance the reaction rate by guiding the reaction along an Catalysts enhance the reaction rate by facilitating an alternate reaction pathway or mechanism with a lower activation energy compared to the uncatalyzed reaction.

From the picture above, we see that there is the coloured area under the curve, that was the area which the particles or molecules have greater energy than activation energy. Without catalyst, that area is the darker one. With the present of catalyst, there are more particles or molecules that have energy greater than activation energy (the total of dark and light purple area). Therefore, more effective collusions happen and consequently, rate of reaction increases.

## Useful Formulas

After we talked about the concept of rate of reaction and some factors that affect it, now let’s discuss about the formulas we use to calculate rate of reaction.  Before we go there, let’s discuss about some useful terms. Look at the table and the graph below!

The table shows concentration of cyclopropane and rate of rection. From the graph, we can see the trend, when the concentration of cyclopropane doubled, the rate of reaction was also doubled. Likewise, when the concentration of cyclopropane decreased one third, the rate of the reaction also reduced one third. Let’s, look at the table, in the third column, there are constant numbers which we get by dividing rate with the concentration. That numbers indicates, the rate of reaction is constantly change proportional to the concentration of reactants. We call that numbers as rate constant, symbolised by k. Unit of k varies, depending on order of the reaction.

So, what is order?

The reaction order indicates how the rate of a reaction is influenced by the concentration of a specific reactant. It represents the exponent to which the concentration of that reactant is raised in the rate equation.

Look at the rate equations below,

A + B -> C + D

Rate = k [A]m [B]n

m and n are the order of reaction, m is the order for reactant A and n is the order for reactant B. m + n is the total order of the reaction

To make uou more understand, see the example below,

Rate = k [H2] [NO]2

From that equation, we know that the order with respect to H2 is 1 and the order with respect to NO is 2. The total order of the reaction is 1 + 2 = 3.

Units of rate constant and the other formulas summarised in the following table.

Source: Chemistry by Cambridge International As and A Level 3rd Edition.

# ENTHALPY CHANGES | Chemistry Cambridge As Level

Hello, everyone, my name’s Putri Dwi Safitri. You can call me Putri. I graduated from IPB university, majoring in Chemistry. Then, I continued my master degree at the University of Edinburgh in Medicinal and Biological Chemistry. I’ve just finished my master degree in August this year. Okay. That’s all about myself. Now, let’s move one. Today, we’re gonna discuss about Enthalpy Changes.

Firstly, before we jump to this topic, let’s find out the difference between exothermic and endothermic reaction. Those terms related to the heat energy transfer during chemical reaction.

## What is exothermic and endothermic reaction? enthalpy changes

Exothermic reaction is the reaction that requires heat energy is transferred to the surrounding. As the result, the temperature of the surrounding increases.

For example, when you react magnesium with sulfuric acid in a test tube, you will feel the test tube becomes warm. It indicates, the temperature of the surrounding (or in this case, the test tube) increases.

That’s it for the exothermic reaction. It’s easy to understand, right?

Exothermic reaction is the reaction that requires heat energy absorption from the surrounding. As the result, the temperature of the surrounding decreases.

For example, when sodium hydrogencarbonate reacts with an aqueous solution of citric acid in a test tube, you’ll feel the test tube becomes cold which means, there is a transfer of heat energy from test tube and the air to the chemical reaction.

Now, you’re able to distinguish between exothermic and endothermic reaction.

Let’s go back to Enthalpy Changes. Read previous article!

## What is enthalpy changes?

Enthalpy is the total energy associated with the material that react. The symbol of enthalpy is H. It must be capital. We cannot measure enthalpy, but we can measure enthalpy changes.

Enthalpy changes is the exchange of the energy between the chemical reactions and its surroundings at constant pressure. The symbol of enthalpy changes is ΔH. The unit for enthalpy change is kilojoule per mole (kJ mol-1). The enthalpy change can be calculated by this formula,

ΔH = H products – H reactants

By using this formula, if the result is negative, means the reaction is exothermic. On the contrary, if the result is positive, the reaction is endothermic.

We can draw a reaction pathway diagram to show enthalpy changes.

The reaction pathway usually also shows the activation energy, Ea, which means, the minimum energy needed for a reaction to happen. For exothermic reaction, the energy is released to the surrounding, so the enthalpy of reactants must be greater than the enthalpy of the products. We can write down the reaction,

Remember, the negative sign for ΔH shows the reaction is exothermic.

For endothermic reaction, the energy is absorbed from the surrounding by the substance, so the enthalpy of the products must be grate than the enthalpy of the reactants.

We can write the reaction,

Always, remember, the positive sign for ΔH means the reaction is endothermic.

After we discussed about exothermic and endothermic reactions, enthalpy change, and how to construct reaction pathway diagram, let’s discuss about,

## A variety of Enthalpy Changes

But firstly, we need to know about standard enthalpy changes which means the reaction is carried out under standard condition which are, a pressure of 101 kPa or 1 atm and a temperature of 298 K.

The symbol used for standard enthalpy change is ΔH°. We can read it delta H standard or some people called delta H nought.

Okay. That’s for standard enthalpy changes.

Now, let’s identify the varieties of enthalpy changes.

According to the type of chemical reaction taking place, we can describe the enthalpy change. For example,

• Enthalpy change of formation
• Enthalpy change of combustion
• Enthalpy change of neutralisation

And in more general case, we can use

• Enthalpy change of reaction

Let’s discuss one by one

### Standard enthalpy change of reaction

Standard enthalpy change of reaction is the enthalpy change when the amounts of reactants shown in stoichiometric equation react to give product under standard conditions. he symbol for this is ΔHr°. For example,

This equation shows us when 1 mol of water is formed from hydrogen and oxygen, 286 kJ of energy is released. However, if we add more hydrogen and oxygen, the energy that is released from this reaction is twice, because 2 mol of water is formed.

### Standard enthalpy change of formation

Standard enthalpy change of formation is the enthalpy change when one mol of compound is formed from its element under standard conditions. The symbol for this is ΔHf°. For example,

### Standard enthalpy change of combustion

Standard enthalpy change of combustion is the enthalpy change when one mol of a substance is burnt in excess of oxygen under standard conditions. The symbol for this is ΔHc°. For example,

### Standard enthalpy change of neutralisation

Standard enthalpy of neutralisation is the enthalpy change when one mol of water is formed from the reaction of an acid and an alkali under standard conditions. The symbol for this is ΔHneut°. For example,

Now that we knew the varieties of enthalpy changes, let’s move on to

## How to measure enthalpy changes

There are at least two ways to measure enthalpy changes, those are

1. Using Hess’s law
2. Bond energies

## Hess’s Law

Drawing Hess’s cycle can help us illustrate Hess’s Law.

This figure shows 3 routes. First, direct route, the reactants A and B combine directly to form C. Indirect route 1, where F and G are formed as intermediates to form product C. Or indirect route 2, where two steps of intermediates formed to form the final product, C.

Hess’s law tells us that enthalpy change of reaction both for direct and indirect route are the same, no matter how many steps are in the indirect route.

You can use this model below to draw an enthalpy cycle for calculating an enthalpy change of reaction from enthalpy change of formation. Dash lines shows us the indirect routes.

When we use Hess’s law, we can see that

ΔH2 (direct route) = ΔH1 + ΔHr (indirect route)

So,

ΔHr = ΔH2 – ΔH1

While the picture below can be used to draw an enthalpy cycle for calculating enthalpy change of formation from enthalpy change of combustion.

When we use Hess’s law, we can see that

ΔH1 (direct route) = ΔHF + ΔH2 (indirect route)

So,

ΔHF = ΔH1 – ΔH2

## Bond energies

Enthalpy change caused by breaking and forming bond. Breaking bond needs energy while forming bond releases the energy.

In chemical reaction,

• If the energy needed to break bonds less than the energy released to form new bonds, the reaction will release energy and is exothermic.
• If the energy needed to break bonds more than the energy released to form new bonds, the reaction will absorb energy and is endothermic.

To calculate the enthalpy changes using average bond energies, we can use this formula,

ΔHr = ΔHbonds broken – ΔHbonds formed

Well, it’s the end for our discussion about Enthalpy Changes today. We’ve learned about exothermic and endothermic reaction, type of enthalpy changes, and how to calculate enthalpy change. Now, it’s time to discuss some exercises about this topic.

## PROBLEM SOLVING

1. Draw full reaction pathway diagrams including activation energy for: (a) The combustion of sulfur to form sulfur dioxide (exothermic); (b) The endothermic reaction of H2O (g) + C (s) à H2 (g) + CO (g)

(a)

(b)

2. (a) Draw the enthalpy cycle for the reaction: 2Al (s) + Fe2O3 (s) à 2Fe (s) + Al2O3 (s); (b)Draw an enthalpy cycle of the formation of ethanol, C2H5OH using enthalpy change of combustion.

(a)

(b)

3. Calculate the standard enthalpy change for the reaction:

2NaHCO3 (s) à Na2CO3 (s) + CO2 (g) + H2O (l)

The relevant enthalpy changes of formation are:

Apply Hess’s law!

4. Use the average bond energies that follow to calculate a value for the enthalpy change for the reaction:

H2 (g) + I2 (g) à 2HI (aq)

5. Propanone is a liquid. It has the structure,

The equation for the complete combustion of propanone is:

CH3COCH3 (l) + 4O2 (g) à 3CO2 (g) + 3H2O (l)

Use the following bond energies (in kJ mol-1) to calculate the value for the standard enthalpy change.

Source: Chemistry Cambridge As and A Level 3rd Edition

# CHEMICAL BONDING

Ionic bonding occurs when positive ions (cations) and negative ions (anions) are attracted to each other through electrostatic forces within an ionic crystal lattice. Covalent bonding, on the other hand, are established by the sharing of outer electrons between two atoms. These resulting ionic or covalent bonds are typically quite robust and require a significant amount of energy to break. Another type of chemical bonding is metallic bonding.

## Ionic Bonding Chemical bonding

When an atom loses one or more electrons, it becomes a positive ion. This is a common occurrence in metal atoms, which tend to lose electrons and transform into positive ions. Conversely, when an atom gains one or more electrons, it becomes a negative ion. This process is typical for non-metal atoms, as they often gain electrons and turn into negative ions. As mentioned previously, cations and anions interaction form ionic bonds or sometimes called electrovalent bond to have the configuration of noble gas.

From Figure 1, we can see that sodium has one electron in its outer shell. Meanwhile, chlorine lacks of an electron in its outer shell to have the configuration of noble gas. This conditions causes the electron from sodium transferred to chlorine. As the result, sodium formed positive ion with charge +1 and chlorine form negative ion with charge -1. Interaction between those two ions called ionic bonding.

We employ dots and crosses in diagrams to represent the electronic arrangement of chloride and sodium ions. This method assists us in tracing the origin of electrons. It’s important to note that the transferred electron is no different from the others; these diagrams are referred to as dot-and-cross diagrams. Please note that this dot-and-cross diagrams only shows the outer electron shells. To make it clearer, let’s see the example below,

From the picture above, we can see that ‘x’ indicates the electrons in the outer shell of magnesium atom while ‘dot’ insicates the electrons in the outer shell of oxygen atom. Because of the electron transfer, oxygen has more electrons which are written by ‘x’, means those electrons are from magnesium. The ion’s charge typically written in the top right-hand corner. From Figure 2, we know that magnesium has charge +2 while oxygen has charge -2.

## Covalent Bonding

There are three types of covalent bonding, including single covalent bonds, muliple covalent bonds, and co-ordinate bonding. Let’s find out one by one!

### Single Covalent Bonds

Single covalent bonds occurs when two non-metal atoms combine by sharing one or more pair electrons. A single covalent bond is symbolized by a single line connecting the atoms, for example Cl─Cl. The picture below will make us more understand about that.

Figure 3 shows chlorine has 7 electron in its outer shell. To distinguish the electrons from each chlorine, electrons in the first chlorine written by ‘x’ while electrons from the second chlorine written by ‘dot’. It is differenet from the ionic bonds. Instead of transfer electrons, the pair of electrons are used together. Other examples of single covalent bondings provides by the figures below,

In certain situations, the electrons surrounding a central atom may not possess the electron arrangement typical of noble gases, which is often referred to as an octet of electrons. For instance:

Boron trifluoride has only 6 electrons around Boron atom. We call this electron deficient.

Phosphorus(V) chloride has 10 electrons around phosphor atom. We call this expanded octet. Other examples of expanded octet are sulfur hexafluoride and sulfur dioxide.

### Multiple Covalent Bonds

Some atoms bond together by sharing two pair of electrons which we call double covalent bonds. Double lines between the two atoms symbolises this double covalent bond, for example O=O.

From figure 4, we know that oxygen needs two electrons to reach octet. When two oxygens combine, each oxygen shares one pair of electron. So, there are two pair of electrons that being used together.

Some atoms can also bond together by sharing three pairs of electorns or we call that triple covalent bond. The figure above is the example for this. Each nitrogen atom needs three electron to reach octet so that when two nitrogen atoms join together, they share three pair of electrons. Just like the picture above, thre lines between the two atoms represents triple covalent bond.

### Co-ordinate Bonding

The last type of covalent bonding we will discuss is co-ordinate bonding. In this case, an atom share a pair of electron to be used together with another atom.

From Figure 6, we can see that ammonia share a pair of electron to be used together with hydrogen ion (proton). The arrow indicates coordinate bond. the head of the arrow points to the atom that needs the electrons.