Chemical Reactions, Stoichiometry, and Chemical kinetics

1. TYPES OF CHEMICAL REACTION AND THEIR STOICHIOMETRY

There are thousands of different chemical reactions. It would be impossible to memorize them all. However, most chemical reactions can be categorised into five major groups. Understanding these categories of reactions can help you predict how reactants will react and what products will form.

Chemical reactions are processes in which substances (reactants) are transformed into new substances (products).

There are five major types of chemical reactions:

1. Combination (Synthesis) Reactions
2. Decomposition Reactions
3. Single Displacement (Replacement) Reactions
4. Double Displacement (Metathesis) Reactions
5. Combustion Reactions

 

Combination (Synthesis) Reactions

A synthesis reaction is a type of chemical reaction in which two or more different substances (elements or compounds) combine and form one compound. Synthesis means “putting together”. You can recognise a synthesis reaction because two or more reactants form only one product.

The general format of a synthesis reaction is shown below.

A + B → AB

Examples:

·         2H₂ ​+ O₂ ​→ 2H_2​O

·         CaO + CO₂ ​→ CaCO₃

·         2Na + Cl₂ → 2NaCl

 

Decomposition Reactions

In a decomposition reaction, one compound breaks down into two or more simpler substances. Notice that decomposition is the reverse of synthesis. You can recognise a decomposition reaction because one reactant forms two or more products. The general format of a decomposition reaction is shown below.

AB → A + B

Examples:

• Decomposition of water by electrolysis
  2H₂O → 2H₂ + O₂​
• Decomposition of calcium carbonate
  CaCO₃ → CaO + CO_2​
• Decomposition of hydrogen peroxide
  2H₂O₂ → 2H₂O + O₂​

 

Single Displacement (Replacement) Reactions

In a single replacement (also called single displacement), one element replaces another element in a compound. In this type of reaction, an element and a compound react to form a different element and a different compound. The general format of a decomposition reaction is shown below.

A + BC → AC + B
(Where A is a more reactive element that displaces B)

Examples:

·         Zinc replaces hydrogen from hydrochloric acid
Zn + 2HCl → ZnCl₂ + H₂​

·         Iron replaces copper from copper (II) sulfate
Fe + CuSO₄ → FeSO₄ ​+ Cu

Note: Reactivity is based on the activity series of metals.

 

Double Displacement (Metathesis) Reactions

In a double replacement (also called double displacement), the positive ions in two compounds switch places, forming two new compounds. In this type of reaction, two compounds react and form two new compounds. The general format of a double replacement reaction is shown below.

AB + CD → AD + CB

Examples:

• Reaction between sodium sulfate and barium chloride

Na₂SO₄ + BaCl₂ → BaSO₄(s) + 2NaCl

• Neutralisation of hydrochloric acid with sodium hydroxide

HCl + NaOH → NaCl + H₂O

 

Combustion Reactions

Combustion is a chemical reaction in which a substance reacts with oxygen and releases energy. This energy is usually released as thermal energy and light energy. For example, burning is a common combustion reaction. The general format of a combustion reaction of a hydrocarbon (a compound made of hydrogen and carbon) is shown below. The products of the combustion of a hydrocarbon are always CO₂ and H₂O.

Cx​Hy​ + O₂ ​→ CO₂​ + H_2​O

Examples:

• Combustion of methane
  CH₄ + 2O₂ → CO₂ + 2H₂O
• Combustion of propane
  C₃H₈ + 5O₂ → 3CO₂ + 4H₂O 

 

Special Cases and Additional Types of Reaction

• Redox Reactions: Involve transfer of electrons (oxidation-reduction).
• Precipitation Reactions: A solid forms from the mixing of two aqueous solutions.
• Acid-Base Reactions: A specific type of double displacement involving proton transfer.

    

                            REACTION STOICHIOMETRY

A balanced chemical equation provides a great deal of information in a very succinct format. Chemical formulas provide the identities of the reactants and products involved in the chemical change, allowing classification of the reaction. Coefficients provide the relative numbers of these chemical species, allowing a quantitative assessment of the relationships between the amounts of substances consumed and produced by the reaction. These quantitative relationships are known as the reaction’s stoichiometry; a term derived from the Greek words stoicheion (meaning “element”) and metron (meaning “measure”).

Stoichiometry: It is the study of the quantitative relationship conveyed by a chemical equation.

You must master the following chemistry concepts to solve stoichiometry problems:

1. Balancing a chemical equation

2. Converting between grams and moles

3. Calculating molar mass

4. Calculating mole ratio

 

Stoichiometric coefficient: Also known as the Stoichiometric number, is the number of molecules participating in the reaction. In simplicity, it is the number in front of atoms, molecules, and ions.

Note: Stoichiometric number can be a whole number or a fraction.

 

 

 

 

 

CHEMICAL ANALYSIS

Definition of terms used in volumetric analysis.

1. Mass concentration: The mass concentration of a solution is the amount of solute present in a given volume of the solution. It is expressed in g/dm³ or g/dm³.

Thus,

Mass concentration = molar concentration X molar mass

2. Molar concentration: The molar concentration (M) of a solution is the number of moles of solutes per dm³ of solution. It can also be defined as the concentration of a solution in moles per dm³.

The following equations are useful in calculations involving molar concentration.

·         Number of moles of a substance = Number of particles    mol^-1

         6.02 x 10²³

·         Number of moles of a substance = Mass of substance    mol^-1

        Molar mass

 

In volume per cm³ of solution, we have

·         Number of moles of solute = Volume    X molar concentration

          1000

3. Standard solution: A standard solution is a solution of known concentration.

4. Molar solution:  A molar solution of a compound contains one mole of the molar mass of the compound in one dm³ of the solution.

 

STOICHIOMETRIC PLAY QUESTIONS

Mass‐Mass relationship

1) What mass of barium chloride is required to react completely with 10.0 g of aluminium sulfate?

2) What mass of chlorine gas is required to react with 10.0 g of aluminium metal?

3) Reaction of hydrochloric acid with a sample of zinc hydroxide gave 0.555 g of zinc chloride. What was the mass of the zinc hydroxide sample?

4) When a barium chloride solution is mixed with a solution containing excess aluminium sulfate, 0.888 g of barium sulfate is obtained. What mass of barium chloride was contained in the solution?

5) 1.00 g of aluminium metal is treated with a solution containing 7.00 g of zinc chloride. What mass of metallic zinc will form?

 

Mass-mole\mole-mass relationship

6) Tin metal reacts with hydrogen fluoride to produce tin (II) fluoride and hydrogen gas, according to the following balanced equation.

Sn(s)+2HF(g)→SnF₂(s)+H₂(g)

How many moles of hydrogen fluoride are required to react completely with 75.0g of tin?

7) Hydrogen sulfide gas burns in oxygen to produce sulfur dioxide and water vapour: 2H₂S(g)+3O₂(g)→2SO₂(g)+2H₂O(g)

What mass of oxygen gas is consumed in a reaction that produces 4.60 mol SO₂?

8) 1.50 mol of KClO₃ decomposes according to this equation (2KClO₃→2KCl + 3O₂). How many grams of O₂ will be produced?

9) If 80.0 grams of O₂ were produced, how many moles of KClO₃ decomposed?

10) How many grams of H₂O are produced when 2.50 moles of oxygen are used?

2H₂ + O₂ → 2H₂O 

Mole-mole relationship

11) When 2.00 mol of N₂ reacts with sufficient H₂ (N₂ + 3H₂ → 2NH₃), how many moles of NH₃ will be produced?

12) Suppose 6.00 mol of H₂ reacted with sufficient nitrogen. How many moles of ammonia would be produced?

13) How many moles of water are produced when 3 moles of methane (CH₄) are combusted?

14) If 3 moles of aluminium react, how many moles of aluminium chloride (AlCl₃) are formed?

(2Al+3Cl₂​→2AlCl₃​)

 

15) How many moles of hydrogen gas will be produced from 10 moles of zinc?

(Zn+2HCl→ZnCl₂​+H₂​)

Mole-Volume relationship

16) What volume of oxygen gas is needed to completely combust 5 moles of methane according to this equation: CH₄​ + 2O₂ ​→ CO₂ ​+ 2H₂O

17) What volume of oxygen gas is produced when 4 moles of potassium chlorate decompose?

18) Na + H₂O → NaOH + H₂​ from the given equation, what volume of hydrogen gas is produced from 1.2 moles of sodium?

19) If 2.0 moles of sodium react with water, what volume of hydrogen gas is released at STP?

 

Volume-Mole relationship

20) You collected 22.4 L of oxygen gas from the decomposition of potassium chlorate. How many moles of O₂ did you collect?

21) You collected 67.2 L of CO₂ gas at STP during the reaction. How many moles of calcium carbonate (CaCO₃) decomposed? (CaCO₃ → CaO+CO₂)

2. CHEMICAL KINETICS (Rates of Reactions)

The branch of chemistry that deals with the study of reaction rates is known as Chemical Kinetics or Reaction Kinetics.

These include:

·         Rate of reaction

·         Mechanism/sequence of steps by which a reaction occurs

·         Factors influencing the rate of reaction

 

Rates of reaction (ROR) are a measure of how fast a reaction will occur.

The rate of a chemical reaction is the number of moles of reactants converted or products formed per unit time.

Usually, the rate of reaction is determined experimentally by measuring the change in concentration of one

of the components in the reaction with time.

Thus,

Rate of reaction = Change in concentration of reactant or product (mol/dm³)

Time taken for the change (seconds)

 

Consider the reaction: A → B

Rate R = -d[A]/dt = d[B]/dt or -∆[A]/∆t = ∆[B]/∆t

where ∆ = change, t= time, A = reactant, B = product

 

The unit of the rate of reaction is mol/dm^-3S^-1 or gdm^-3S^-1.

 

The rate of reaction can also be expressed as:

Rate of reaction = Change in number of moles or mass of reactant or product

Time taken for the change

 

Then the unit of rate is molS^-1 or gS^-1

 

EXAMPLES:

1. When 0.5g of calcium trioxocarbonate (IV) was added to excess dilute hydrochloric acid, carbon (IV) oxide was evolved. The complete reaction took 5 minutes. What was the rate of reaction?

SOLUTION:

Rate of reaction = Change in number of moles or mass of reactant or product

Time taken for the change

 

Rate of reaction = (0.5-0) g     = 1.67 x 10^-3 gS^-1

                             (5 x 60) S

 

2. Consider the following reaction:

H₂O₂ (aq) + 3I^– (aq) + 2H⁺ (aq) → I₃^– (aq) + 2H₂O (l)

In the first 10.0 seconds of the reaction, the concentration of I– dropped from 1.000 M to 0.868 M.

Calculate the average rate of this reaction in this time interval.

 

Solution:

Average Rate = −Δ[Reactant]​

                                    Δt

 

Given Data:

• Initial concentration of I⁻: 1.000 M
• Final concentration of I⁻: 0.868 M
• Time interval: 10.0 seconds

Eqn: H₂O₂ (aq) + 3I^– (aq) + 2H⁺ (aq) → I₃^– (aq) + 2H₂O (l)

 

Rate _I− ​= −Δ[I⁻]​ = - (-0.132M)    = 0.0132 M/s

                    Δt            (10s)

           

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                             COLLISION THEORY

Reactions happen when molecules, elements, or atoms collide with the proper orientation and sufficient energy, but not all collisions result in a reaction.

 

The collision theory states that for a chemical reaction to occur, the reactant particles must collide and they must collide with a certain minimum amount of energy known as activation energy.

 

Collisions, which result in chemical reactions, are called EFFECTIVE COLLISIONS. The minimum amount of energy required by reacting particles for a chemical reaction to occur is called ACTIVATION ENERGY. Activation energy is the ENERGY BARRIER that the reactants must overcome for the reaction to occur. It is the minimum energy required for bond breaking for a chemical reaction to occur.

Chemical reactions occur only when the energy of the colliding reactant particles is equal to or greater than the activation energy. Activation energy must be equal to the energy barriers, also for a chemical reaction to occur.

 

Note: Every reaction has its energy of activation. Reactions with low activation energy have a high rate of reaction and occur spontaneously. Reactions with high activation energy have a low rate of reaction and are not spontaneous.

 

FACTORS AFFECTING THE RATE OF REACTION

From the collision theory, it can be seen that the rates of reaction depend on the following features.

1. The energy of the particle.

2. The frequency of collision of the reaction.

3. The activation energy of the reaction.

 

These features of a chemical reaction are, in turn, affected by some factors, which can cause them to change and consequently affect the rate of reaction. These are factors that affect the rate of reactions.

Some important ones are:

1. Nature of reactants.

2. Concentration/pressure (for gases) of reactants.

3. Surface area of reactants

4. Temperature of the reaction mixture

5. Presence of light

6. Presence of catalysts

To study the effect of any one of these factors on the rate of reaction, all other factors must be kept constant.

 

EFFECT OF NATURE OF REACTANTS

If all other factors are kept constant, different substances will have different rates of reaction with dilute HCl, for example. When dilute HCl reacts with zinc, iron, and gold under the same conditions, hydrogen gas is evolved fast with zinc, slow with iron, and no gas is evolved with gold.

The difference in rate of reaction is due to the chemical nature of the elements as they naturally possesses a different amount of energy content.

 

EFFECT OF CONCENTRATION OF REACTANTS

The frequency of collision among particles is high when the particles are crowded in a small space, i.e high concentration. This leads to highly effective collision and thus a high rate of reaction. An increase or decrease in the concentration of the reactants will result in a corresponding increase or decrease in effective collisions of the reactants and hence the reaction rate.

 

EFFECT OF SURFACE AREA OF REACTANTS

This is a very important factor to be considered when a solid is involved in a chemical reaction. Lumped solids offer a small surface area of contact for reaction, while powdered solids offer a large surface area for reaction. The rate of reaction is slow with lumped solids but high with powdered solids.

 

EFFECT OF TEMPERATURE

Increasing the temperature of a system can lead to an increase in reaction rate in two ways. When heat is raised, energy in the form of heat is supplied to the reactant particles, so that

1. The number of particles with energy equal to or greater than the activation energy increases.

2. The velocity of all the reactant particles increases due to the greater kinetic energy, leading to

a higher frequency of collision.

As a result, the number of effective collisions increases and the reaction proceeds at a faster rate. Decreases in temperature lead to a decrease rate of reactions.

 

EFFECT OF LIGHT

Some reactions are influenced by light. The rate of reaction is high when the light intensity is high, low when the intensity is low and does not proceed at all in the absence of light. Such reactions are known as photochemical reactions. Examples of photochemical reactions include.

1. Reaction between hydrogen and chlorine and

2. Decomposition of hydrogen peroxide

3. Reactions between methane and chlorine

4. Photosynthesis in plants

5. Conversion of silver halides to grey metallic silver

 

EFFECT OF CATALYST

A Catalyst is a substance which alters the rate of a reaction, but itself does not undergo any change at the end of the reaction.

A positive catalyst increases the rate of reaction by lowering the activation energy of the reaction whereas, the one which increases the activation energy is known as a negative catalyst or an inhibitor.

  

 

RATE LAW

The rate law (also called the rate equation) expresses the rate of a reaction as a function of the concentration of reactants, each raised to a power called the order of the reaction. Rate law tries to explain the dependence of reaction rate on concentration.

 

For a general reaction:

aA + bB → Products

The rate law is:

Rate = k[A]^m[B]ⁿ         

Rate = Reaction rate (usually in mol·L⁻¹·s⁻¹)

• k = Rate constant (depends on temperature and nature of reaction)
• [A], [B] = Concentrations of reactants A and B
• m, n = Reaction orders with respect to A and B (determined experimentally)

 

REACTION ORDER

The order of a reaction with respect to a reactant tells us how the rate is affected by changes in that reactant’s concentration.

·         If m = 1, the rate is directly proportional to [A] → First order

·         If m = 2, the rate is proportional to [A]² → Second order

·         If m = 0, the rate is independent of [A] → Zero order

·         Overall order = m + n + …

 

Overall Reaction Order

 The overall reaction order is the sum of the orders of the reactant species

For a reaction: A + 2B → AB₂

Rate = k[A]¹[B]²     

 

Overall Reaction Order = m + n = 1 + 1 = 2

Reaction is of the Second Order Overall

 

 

 

 

 

Examples

 

 

EXPERIMENTAL DETERMINATION OF RATE LAW

 

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