Endothermic and Exothermic Processes Unit 6.1

Endothermic and Exothermic Processes Unit 6.1

by Jennifer Haycraft

• ​Enthalpy is the heat (q) exchanged at constant pressure.

• Changes in heat energy in a system is represented with the letter Q

  • In a constant pressure system, the net energy change is called the Enthalpy and is represented as ΔH°

​Breaking and Forming Bonds

• ​Breaking bonds takes energy

  • Think of trying to pull apart very strong magnets: it required a great deal of energy on your part to pull them apart

  • Because it requires energy to break these bonds, the process is energetically unfavorable

  • When bonds are broken, our ΔH° is POSITIVE

​Forming Bonds

• ​Making bonds release energy

  • Let's think of those magnets again: when strong magnets are put next to each other, they snap back together. You spend no energy

  • Because energy is given off, this process is considered energetically favorable

  • When bonds are created our ΔH° is NEGATIVE

​Endothermic vs Exothermic Reaction

• ​

System Change is ambiguous.​

• Physical Change

  • ​The salt can be retrieved

  • ​Evaporation is a physical process

• ​Chemical Change

  • ​Dissolving is often accompanied by a chemical change.

  • ​Ionic bonds are broken which is an indicator of a chemical change.

​Do or do not, there is no try.

• ​NO Trying to destroy energy.

• • Laws of thermodynamics:

    • The first law, also known as the Law of Conservation of Energy, states that energy cannot be created or destroyed

    • The second law of thermodynamics states that the entropy of any isolated system always increases.

    • The third law of thermodynamics states that the entropy of a system approaches a constant value as the temperature approaches absolute zero.

​So where does the energy go or come from?

• ​System vs. Surroundings

  System:

  • Where the reaction takes place

• ​Surroundings:

  • ​Whatever doesn't take place in the system (basically everything else)

​Endo/Exothermic Rxns

• Endothermic:

  • Heat is absorbed from the surroundings.​

• ​Exothermic:

  • ​ Heat is released to the surroundings.

Enthalpy (H) vs. Heat (Q)

​

We often use the terms enthalpy and heat interchangeably, but there are important differences:

• enthalpy (H) is a thermodynamic quantity equivalent to the total energy (heat content) of a system and cannot be measured directly.

• enthalpy change (∆H) describes amount of heat transferred during a chemical reaction at constant pressure and this can be measured (or calculated from measurements).

  whereas heat is a form of energy that can be transferred from one place to another and can be measured directly from changes in temperature

q = m C ∆T where q = heat transferred

C = specific heat capacity

m = mass

∆T = change in temperature

when a chemical change is happening in such a way there is no energy being used to do work, then we can assume:

q = mC ∆T = ∆H

(heat transferred to/from surroundings) = (heat transferred from/to chemicals) can be measured can be calculated

q = mC ∆T and q = ∆H

Heat can refer to any quantity of reaction - it is not directly dependent on the amount or form of the reactants

Enthalpies are defined in terms of specific states and specific amounts (usually 1 mole of something).

Having calculated q (the heat energy transferred into the water), we can now calulate ∆H on the assumption that:

the heat energy transferred into the water = the chemical energy released by the reaction

q = ∆H

However, since the water is heating up (gaining energy ∆T = +ve value q = +ve value) the chemical reaction is losing energy (exothermic) and ∆H = -ve value. (Similarly, if the water is cooling down (losing energy ∆T = -ve value q = -ve value) the chemical reaction is gaining energy (endothermic) and ∆H = +ve value.)

So,

∆H = - q

∆H = - 2299 J

The ∆H calculated at this stage is simply an 'Enthalpy of Reaction', ∆Hrxn

Similarly, 'Heat of .........' can be different from 'Enthalpy of .......' as the former can be defined for whatever quantity you want, whereas the latter must always refer to 1 mole.

Vaporization refers to the process of a liquid (at its boiling point) changing into a gas.

e.g. Heat of vaporization of water = 2250 J g-1 or 2.25 kJ g-1 or 2250 kJ kg-1

Enthalpy of vaporization of water = 40.53 kJ mol-1 (1 mol of H2

O = 18g)

Open Ended

When 500.0 mL of 2.00 mol/L Ba(NO3)2 solution at 22.50°C is combined in a coffee cup calorimeter with 500.0 mL of 2.00 mol/L Na2SO4 solution also at 22.50°C, a white

precipitate forms and the temperature of the mixture rises to 25.60°C.

Assume the calorimeter materials absorb only a negligible quantity of heat and the final

solution’s density and specific heat capacity are identical to those of water.

Calculate the molar enthalpy of precipitation of BaSO4(s).

Open Ended

 A coffee cup calorimeter contains 50.00 g of water at 20.73°C. When 2.13 g of NH4NO3

pellets are stirred into the water, the temperature falls to 17.41°C.

Assume the heat capacity of the resulting solution is the same as that of water and that no

energy is absorbed or released from or to the surroundings. Calculate the molar enthalpy of dissolution of ammonium nitrate, a chemical commonly used in the production of cold packs.

Open Ended

Given: 2 HCl(aq) + Ba(OH)2(aq) ⟶ BaCl2(aq) + 2 H2O(l) ∆H = - 118 kJ mol-1

Calculate the heat released when 300.0 mL of 0.500 mol/L hydrochloric acid is combined with 100.0 mL of 1.00 mol/L of barium hydroxide. Assume that the temperature of both solutions is initially 21.55°C. The final mixture has a mass of 400.0 g and a heat capacity the same as that of water with negligible heat “leakage” to or from the system. What is the final temperature of the mixture?