Thermochemistry: Energy and Enthalpy PDF

Basic Chemistry 2 THERMOCHEMISTRY ENERGY AND ENTHALPY Fithri Choirun Nisa THERMOCHEMISTRY Definition: Chemical reaction in the (food) material on energy based Thermochemistry 1st Thermodynamic 2nd Thermodynamic Energy Entropy Heat...

Basic Chemistry 2 THERMOCHEMISTRY ENERGY AND ENTHALPY Fithri Choirun Nisa THERMOCHEMISTRY Definition: Chemical reaction in the (food) material on energy based Thermochemistry 1st Thermodynamic 2nd Thermodynamic Energy Entropy Heat Calorimetry Work Activation energy Internal Energy Enthalpy Thermochemistry in Food Food Processing - Heating (cooking, baking, pasteurization, sterilization) - Cooling and freezing - Drying Thermochemistry in Food (cont.) Food Nutrition Each food with different composition - different energy (water, carbohydrate, protein, fat) ENERGY Definition: Capacity for doing work or supplying heat An object can have energy: - Kinetic energy (KE): energy possessed by an object when they move - Potential energy (PE): energy possessed by an object due to attraction or repulsion Total energy: KE + PE ENERGY MEASUREMENT (CALORIMETRY) International Unit (IU): Joule (J) 1 J = 1 kg.m2/sec2 1 erg = 1 x 10-7 J Calorie (cal): the amount of heat required to raise the temperature of 1 gram of water at an original temperature of 15C by 1C Kilocalories (Cal): energy contained in food 1 cal = 4.184 J 1 kcal = 1 Cal = 4184 J Conservation of energy (First Law of Thermodynamics) Energy cannot be created or destroyed. Energy can only be changed from one form to another KE  PE To add energy to one object, it must be taken from another object First Law of Thermodynamics Energy = internal energy (U) E = Efinal - Einitial Internal energy = kinetic and potential energy Esys = -Eenv In chemistry → E system E = q + w q = the amount of heat exchanged between the system and environment w = work done on / by the system First Law of Thermodynamics (cont.) Sign agreement for work and heat Process Sign Proses Tanda Work is done by the system on the - environment Work is done on the system by the + environment Heat is absorbed by the system from the + environment (endothermic process) Heat is absorbed by the surroundings from the - system (exothermic process) Work and Heat Work (w) = mechanical work (expansion of gas) w = -PV P = pressure V = change in volume If there is no change in volume (constant volume), then: E = q If there is a change in volume, then E = q + w E = q - PV Energy in atom and molecule Atom, molecul, ion also have same energy form: kinetic energy and potential energy Energy in atom and molecule (cont.) Kinetic energy Particles in atomic size move and collide with each other Atoms and molecules in motion have kinetic energy: - moving slowly - small kinetic energy - moving fast - large kinetic energy The average value of kinetic energy of atom directly proportional to the absolute temperature (K) of the object Energy in atom and molecule (cont.) Potential energy = Chemical energy Atoms are made of electrically charged particles that either attract or repel Potential energy changes with distance If a substance reacts, there is a change in the nature of attraction (chemical bonds) between the atoms → change in potential energy Heat energy Heat is kinetic energy (of atoms and molecules) In a hot substance, the value of its kinetic energy is large and the heat it contains is large and vice versa If two hot and cold substances are brought close to each other, their collision will cause them to have the same temperature ENERGY CHANGES IN CHEMICAL REACTION Chemical reaction in an enclosed space: no energy goes in or out: total energy (KE and PE) remains constant Chemical reactions are not closed (in general) - exothermic change: the temperature of the reaction mixture will increase and potential energy decrease - endothermic change: the temperature of the reaction mixture will decrease and the potential energy increase ENERGY CHANGES IN CHEMICAL REACTION Heat capacity and specific heat Calorie → the change in temperature that water undergoes when it takes in or releases heat Heat capacity: the amount of heat required to change the temperature of an object by 1C - extensive (the value depends on the amount) - units: J/C or cal/C Specific heat: the amount of heat required to raise the temperature of 1 g of a substance by 1C - intensive - unit: J/g.C or cal/g.C Calorimeter Constant volume calorimeter : heat of combustion → records the rise in water temperature → isolated system: no heat and mass lost to the environment Constant pressure calorimeter : determine the heat change for reactions other than combustion (heat of neutralization, heat of ionization, heat of fusion, heat of vapor, heat of reaction) HEAT OF REACTION AND THERMOCHEMISTRY Thermochemistry is a branch of chemistry that studies the relationship between reactions and heat System and environment - The system isolated from the environment: adiabatic change: system temperature varies (exothermic reaction and endothermic reaction) - The system not isolated from the environment: isotherm change: system temperature remains constant HEAT OF REACTION AND THERMOCHEMISTRY (cont.) If the reaction is exothermic or endothermic, the chemical substances involved will change in potential energy The heat of reaction is measured as the change in potential energy (PE) PE = PE final - PE initial Exothermic change → PE negative Endothermic change → PE positive HEAT OF REACTION AND THERMOCHEMISTRY (cont.) Exothermic reaction a reaction which accompanied by the release of heat. The heat of reaction is written with a negative sign. Example : N2 (g) + 3H2 (g) → 2NH3 (g) - 26,78 Cal Endothermic reaction reactions that require heat. The heat of reaction is written with a positive sign Example 2NH3 (g) → N2 (g) + 3H2 (g) + 26,78 Cal Enthalpy and Enthalpy Changes Reaction heat measurement - constant volume - changing pressure - constant pressure Most reactions are in an open container so they correspond to a constant external pressure of the atmosphere - constant pressure Heat of reaction at constant pressure: enthalpy change of reaction (H) H = H final - H initial Enthalpy Constant pressure process E = q + w = qp - PV or qp = E + PV H = E + PV H = enthalpy H = E + (PV) Constant pressure: H = E + PV Constant pressure process, qp = H Constant volume process, qv = E HESS’S LAW (TOTAL HEAT) Enthalpy is a state function → H of the reaction is independent on the path of reactants pass to form products The overall change as a result of the sequence of steps and the value of H for the whole process is the sum of the enthalpy changes that occur throughout the process → thermochemical equation HESS’S LAW (TOTAL HEAT) (cont.) Example: H2O (l) → H2 (g) + 1/2 O2 (g) H = +283 kJ H2 (g) + 1/2 O2 (g) → H2O (g) H = -242 kJ so H2O (l) → H2O (g) H = + 41 kJ The use of thermochemical equations Changes in the thermochemical equation before combining → changes in the value of H Multiplication or division of the coefficient by several factors Example: H2O (l) → H2 (g) + 1/2 O2 (g) H = +283 kJ so 2 H2O (l) → 2 H2 (g) + O2 (g) H = +566 kJ The use of thermochemical equations (cont.) Changing the direction of the chemical equation Example: H2O (l) → H2 (g) + 1/2 O2 (g) H = +283 kJ so H2 (g) + 1/2 O2 (g) → H2O (l) H = -283 kJ Problem Using Hess’s Law by combining chemical equations (1) 2 C2H2 (g) + 5 O2 (g) → 4 CO2 (g) + 2 H2O(l) H = -2602 kJ (2) 2 C2H6 (g) + 7 O2 (g) → 4 CO2 (g) + 6 H2O (l) H = -3123 kJ (3) H2 (g) + 1/2 O2 (g) → H2O (l) H = -286 kJ All data are valid at 25C and 1 atm pressure. What is the value of H for the following reaction: (4) C2H2 (g) + 2 H2 (g) → C2H6 (g) Key Concepts Bond energy = Bond enthalpy : energy required per mole of a compound to break a particular bond to produce fragments at 25C and 1 atmosphere pressure - Breaking chemical bonds requires an input of energy (H is positive = endothermic reaction) - Making chemical bonds releases energy (H is negative = exothermic reaction) - Bond energy can be used to indicate the stability of the compound Heat of formation (enthalpy of formation) : enthalpy change associated with formation of one mole of a compound from its elements (Hf) Example: Formation of water and water vapor at a temperature of 100C and 1 atm H2 (g) + 1/2 O2 (g) → H2O (l) Hf = -283 kJ H2 (g) + 1/2 O2 (g) → H2O (g) Hf = -242 kJ  Heat of formation (enthalpy of formation) Heat of vaporization of water? H2O(l) → H2O(g) Heat of reaction: H = Hf H2O (g) - Hf H2O (l) H = -242 kJ - (-283) = +41 kJ STANDARD STATE The value of H depends on the temperature, pressure, and the physical state (gas, liquid, solid) of the reactants and products. Different conditions → different value of H → standard state: 25C and 1 atm (H) (The standard value of heat of formation for each element = 0) Standard heat of formation table → standard heat of reaction Problem Calculate the H of a reaction from heat of formation at standard state Baking soda can be used to extinguish the fire from its decomposition products: 2 NaHCO3 (s) → Na2CO3 (s) + H2O (g) + CO2 (g) Calculate the H for the reaction Thermochemistry and Nutrition https://saylordotorg.github.io/text_general-chemistry-principles-patterns-and- applications-v1.0/s09-04-thermochemistry-and-nutrition.html#averill_1.0- ch05_s04_t03