Energetics Of Chemical Reaction Class 12 Chemistry Nepal

Homogeneous system and heterogeneous system:

Homogeneous system:

A uniform system throughout is called a homogenous system. In other words, a system is said to be a homogeneous when it has the same chemical composition throughout. For e.g., a mixture of gases, a pure single liquid.

Heterogeneous system:

A heterogeneous system is one that consists of two or more phases. In other words it is not uniform throughout. E.G. ice in contact with water, ice in contact with vapor, etc.

System,boundary and surroundings

System:

The part of the universe under considerationis called system,on whichthe effect of certain variables such as temperature and pressure, is tobe studied.

Depending upon the nature of boundary the system is of three types”

(i) Open system: It exchanges only energy not matter.

(ii) Closed system: It exchanges both matter and energy.

(iii) Isolated: It exchanges neither energy nor matter

Boundary:

The real or imaginary surface separating the system from the surroundings is called boundary.

Surroundings:

Region outside the boundaries of any system is termed surroundings.

Internal energy:

The internal energy of a system is an intrinsic value of the sum of the potential and kinetic energy possessed by a system

State function:A quantity in thermodynamics, such as entropy or enthalpy, that has a unique value for each given state of a system. It is the value that depends on the state of the substance and not as how the state is reached. For e.g: density is a state function because a substance density is not affected by how the substance is obtained. Consider a quantity of H₂O, it doesn’t matter whether that H₂O is obtained from the tap, from a well, or bottle because as long as all three are in the same state, they have the same density. State function depends on three things property, the initial value and not on the objects history or path taken to get from the initial to final value.

State variables:

A state variable is one of the variables used to describe the state of a dynamical system. Each state variable corresponds to one of the coordinates of the underlying state space. An intuitive introduction to state variables is given in the idea of a dynamical system.

Types of thermodynamics

Isothermal and adiabatic process:

A process is said to be isothermal, if the temperature of the system remains constant during each stage of the process. Such type of process may be achieved by placing the system in thermostat.

Here, dT = 0, keeping temperature constant.

A process is said to be adiabatic, if no heat enters or leaves the system during any step of the process. It may be achieved by carrying wt. in a vessel which is completely isolated from its surrounding.

Adiabatic process:

Here, dQ= 0, keeping heat constant.

Isochoric and isobaric process:

A process is said to be isochoric, if the volume of the system remains constant during each step of the process.. Such type of process may be achieved by carrying out in a vessel which is provided with weightless and frictionless piston, if the volume of the system remains constant during each step provided with a fixed piston.

Here, dv = 0.

A process is said to be isobaric if the pressure of the system remains constant during each step of the process. Such type of process may be achieved by carrying out in a vessel which is provided with a weightless and frictionless piston.

Here, dp = 0

Cyclic process:

When a system, after completing series of changes returns to its original state is said to have completed a cycle. The entire process is known as cyclic process.

Here, dE = 0, dH = 0.

First law of thermodynamics:

The first law of thermodynamics is in fact an application of the broad principle known as the Law of conservation of energy to the thermodynamic system. It states that “The total energy of an isolated system remains constant though it may change from one form to another.

When a system is changed from state A to B, it undergoes a change in the internal energy from E_A to E_B. Thus, we can write,

Or, $Δ E = E_{B} - E_{A}$ .

This energy change is brought about by the evolution or absorption of heat and/or by work being done by the system. Because the total energy of the system must remain constant, we can write the mathematical statement of the first law as:

Or, $Δ t$ = q – w.

Where, q is the amount of heat.

w is work done by the system.

Sign convention of heat and work:

When working numerical problems we will quickly become confused if we don’t adopt a universal convention for when we use a positive sign or a negative sign.

Sign Convention for heat, q

1. Heat is transferred into the systemi.e q > 0

2. Heat is transferredout of the systemi.e q < 0

Sign Convention for work, w

3. Work is done upon the system by the surroundings, w> 0

4. Work is done by the system on the surroundings, w< 0

Let’s look at some processes to get a better feel for defining a thermodynamic system and using the proper sign convention.

Example

Hold a piece of ice in your hand until it melts

Solution A

1. System: You

2. Surroundings: Ice + the rest of the universe

3. q< 0 � Heat flows out of the system (you) into the ice.

Solution B

1. System : Ice

2. Surroundings : You + the rest of the universe

3. q> 0 , Heat flows into the system (ice) from you.

You can see that the answer changes depending upon how you define the system, but the physical reality is exactly the same, but both solutions A and B are correct. It doesn’t matter how you define the system as long as you are consistent.

Example

Consider the evaporation of sweat from your body.

Solution A

1. System: The sweat

2. Surroundings: Your body + the rest of the universe

3. q> 0 : Heat flows into the system (sweat) from you in order to raise the kinetic energy of the sweat molecules enough to allow them to go from the liquid phase to the gas phase.

Solution B

1. System: You

2. Surroundings: The sweat + the rest of the universe

3. q< 0 ,Heat flows out of the system (you) into the sweat.

Since heat leaves your body this cools you down. That’s why we sweat after all.

Enthalpy and derivation of $Δ H = Δ E + p Δ v$

Enthalpy:A thermodynamic quantity equivalent to the total heat content of a system. It is equal to the internal energy of the system plus the product of pressure and volume.

The change in enthalpy associated with a particular chemical process. Most compounds have negative enthalpies of formation.

Derivation of $Δ H = Δ E + p Δ v$

To formulate this eqn. consider a gaseous closed system fitted with weightless and frictionless piston.

Let E₁ be the internal energy of the system is changed at state(I). When $Δ$ Q quantity of heat is supplied to the system, the total energy before change = E₁ + $Δ$ Q. But, the system is changed to state (II) by the application of quantity of heat, where the internal energy becomes E₂. In this change $Δ$ W be the amount of work done. After change, the total energy = E₂ + $Δ$ W

Since energy is conserved,

Or, E₁ + $Δ$ Q = E₂ + $Δ$ w

Or, $Δ$ Q = (E₂ – E₁) + $Δ$ W

Or, $Δ$ Q = $Δ$ E + $Δ$ W …(i)

The equation (i) shows that the heat supplied to any gaseous closed system changes internal energy by performing work done.

Case 1:

For cyclic process,

Or, $Δ$ Q = $Δ$ W [for equation (i)]

Case 2:

For adiabatic process

From 1^st law,

Or, $Δ$ Q = $Δ$ E + $Δ$ W

So, 0 = $Δ$ E + $Δ$ W

Or, $Δ$ W = – $Δ$ E

Case 3:

The change of gaseous system at constant pressure is due to the volume expansion

i.e. $Δ$ W = P. $Δ$ v

We know:

From 1^st law expression for gaseous system at constant pressure:

So, $Δ$ Q = $Δ$ E + p $Δ$ V.

Hess’s law of constant heat summation

It states that, “The amount of heat evolved or absorbed in a chemical change is the same whether the process takes place in one step or in several steps”. (i.e. it follows 1st Law of Thermodynamics).

For example, carbon can be oxidized to CO₂ either directly or in two different steps as given below:

1st method: C(s) + O₂(g) à CO₂(g) ∆H = -94.3kcal

2nd method: C(s) + $\frac{1}{2}$ O₂(g) à ∆H₁ = -26 kcal

C(s) + $\frac{1}{2}$ O₂(g) à ∆H₁ = -68.3 kcal

According to Hess’s law ∆Hmust be equal to ∆H₂ + ΔH₂

Heat of neutralization:

It is defined as a the heat of reaction resulting from the neutralization of an acid or base; especially :the quantity produced when a gram equivalent of a base or acid is neutralized with a gram equivalent of an acid or base in dilute solution.

Heat of solution:

The heat of solution is defined as the enthalpy change when one mole of a substance (solute) is dissolved in specified number of moles of a pure solvent. It is commonly known as integral heat of solution.

Heat of combustion :

The heat of combustion (Δ H_c ) is the energy released as heat when a compound undergoes complete combustion with oxygen under standard conditions. The chemical reaction is typically a hydrocarbon reacting with oxygen to form carbon dioxide, water and heat. It may be expressed with the quantities:

1. energy/mole of fuel (kJ/mol)

2. energy/mass of fuel

3. energy/volume of the fuel

Heat of vaporization:

The amount of heat required to vaporize one gram of a liquid at its boiling point with no change in temperature is called heat of vaporization. Usuallyexpressed in J/g. The molar heat ofvaporization is the amount of heat required to vaporize one mole of liquid at its boiling point with no change in temperature and usually expressed ion kJ/mol.

Heat of formation:

The standard enthalpy of formation is defined as the change in enthalpy when one mole of a substance in the standard state (1 atm of pressure and 298.15 K) is formed from its pure elements under the same conditions.

Bond energy:

In chemistry, bond energy (E) is the measure of bond strength in a chemical bond. It is the heat required to break one mole of molecules into their individual atoms

Energetics Of Chemical Reaction

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