Enthalpy Changes C2.1 Thermochemsitry 21 September 2018

Enthalpy Changes Chemical reactants have a certain amount of chemical energy stored within them. An enthalpy change (∆H) is the difference in energy between the reactants and products.

Exothermic Enthalpy Changes An exothermic chemical reaction is where the products have less energy than the reactants. Energy is given out to the surroundings (temperature goes up). Examples: Combustion, neutralisation, respiration We will draw the enthalpy profile for an exothermic reaction labelling reactants, products, activation energy, overall energy change (H), bond breaking and bond making.

Endothermic Enthalpy Changes An endothermic chemical reaction is where the products have more energy than the reactants. Energy is taken in from the surroundings (temperature goes down). Examples: Electrolysis, thermal decomposition We will draw the enthalpy profile for an endothermic reaction labelling reactants, products, activation energy (energy needed to break the bonds), overall energy change and sign of H, bond breaking and bond making.

Worked Example Draw an enthalpy profile for the combustion of methane. Your profile include: Reactants and Products Overall energy change Activation energy Type of reaction (exothermic or endothermic) Bonds broken Bonds made

Knowledge Check Sketch enthalpy profile diagrams for the following reactions: N2 + 3H2  2NH3 ∆H = -92KJ mol-1 N2O4  2NO2 ∆H = +58KJ mol-1 Your answer should include: Reactants and Products Overall energy change Activation energy Type of reaction (exothermic or endothermic) Bonds broken Bonds made

Standard Enthalpy changes The amount of energy transferred in a given chemical reaction depends on the conditions under which the reaction occurs. Chemists use standard conditions to help them compare data! Standard enthalpy change is the energy change of a particular reaction performed under standard conditions. Standard conditions: 298K (25oC) 100kPa or 1 bar pressure Concentrations of 1 moldm-3

Knowledge Check Define the term standard enthalpy change giving any necessary conditions. [2 marks] Amount of heat energy absorbed or evolved in a reaction performed under standard conditions (1) 298K, 100kPa, 1moldm-3 (1) The complete combustion of methane in air is an exothermic reaction. How does the energy level diagram show that this reaction is exothermic? [2 marks] Energy level of reactant is higher than products (1) So energy/heat given out during reaction (1)

Measuring Enthalpy changes Enthalpy change is measured in joules, J or kilojoules, kJ ∆H = Σ (bond enthalpies of bonds broken) – Σ (bond enthalpies of bonds made) In words, the energy change for a reaction is the sum of the average bonds enthalpies in the reactants minus the sum of the bond enthalpies in the products.

Average bond enthalpy Average bond enthalpy is the energy required to break one mole of a specified type of bond in a gaseous molecule. An input of energy is always required to break bonds. Bond enthalpies are always endothermic. Bond enthalpies are always positive. Bond enthalpies can be used to measure the enthalpy change of a reaction by taking into account the energy of the bonds broken vs. the energy of the bonds made.

Calculating an enthalpy change worked exampleS Using average bond enthalpies, calculate ∆rH for the reaction of propane with oxygen. (Balanced equation  Displayed formula  Number of each bond  Energies of each bond  ∆rH) Bond Bond enthalpy (kJ) C—H 413 C=C 612 C—C 347 C=O 805 C—O 358 N—H 391 O—H 464 H—H 436 O=O 498 Cl—Cl 243 H—Cl 432 I—I 151 H—Br 366 H—I 298 N≡N 945 Br—Br 193

Calculating an enthalpy change worked exampleS Using average bond enthalpies, calculate ∆rH for the reaction of chlorine with ethane in the presence of UV light. (Balanced equation  Displayed formula  Number of each bond  Energies of each bond  ∆rH) Bond Bond enthalpy (kJ) C—H 413 C=C 612 C—C 347 C=O 805 C—O 358 N—H 391 O—H 464 H—H 436 O=O 498 Cl—Cl 243 H—Cl 432 I—I 151 H—Br 366 H—I 298 N≡N 945 Br—Br 193

Knowledge check Complete the worksheet questions, finishing with a previous exam question.

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