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Thermochemistry Download Presentation ## Thermochemistry

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1. Thermochemistry Ochran 2014

2. System A system is that part of the universe that is under study. Everything else in the universe is called the surroundings.

3. Chemists have defined three types of systems • an open system can exchange both energy and matter with its surroundings • a closed system can exchange energy but not matter, with its surroundings • an isolated system cannot exchange energy or matter with its surroundings a thermos bottle is a close approximation of an isolated system since it minimizes heat transfer to the surroundings

4. These concepts can be expressed as an equation: universe = system + surroundings Interactions between system and its surrounding typically involve the exchange of energy and matter

5. State Function In thermochemistry, we will encounter enthalpy change,H , entropy change,S , free energy change,G , and energy change,E . The numerical values and mathematical signs of these state functions depend only on the different between the final state and the initial state of the system. Two quantities that are not state functions are heat(q) and work(w). The values for these quantities depend on the sequence of steps used to transform matter from the initial to the final state.

6. Measuring Energy Changes When methane reacts with oxygen in a lab burner, enough heat is transferred to the surroundings to increase the temperature and even to cause a change of state The experimental technique is called calorimetry • it depends on careful measurements of masses and temperature changes

7. q = m • c • T Where • c = the specific heat capacity, the quantity of heat required to raise the temperature of 1g of a substance by 10 C • q = the quantity of heat that flows, varies directly with quantity of substance, mass (m) • T temperature change

8. Practice Problem 1 When a 1.25 kg sample of water was heated in a kettle, its temperature increased from 16.4 0C to 98.8 0C . How much heat did the water absorb? c = 4.19J/g. 0C

9. The First Law of Thermodynamics: Energy is Conserved When a system absorbs energy, the surroundings release it. Similarly, when a system releases energy, the surroundings absorb it. Esystem = ‾Esurroundings

10. Heat Transfer and Enthalpy Change • enthalpy,H, sometimes called the heat content of the system, is defined as: H = E + PV H The change in enthalpy of a system depends only on the initial state and on the final state  H =  E +  (PV)

11. 45o C 25o C magnesium Hydrochloric acid Exothermic Reactions • Magnesium + Hydrochloric acid Heat energy given out Gets hot

12. Exothermic Reactions • If heat is given out this energy must have come from chemical energy in the starting materials (reactants). 45o C 25o C Reactants convert chemical energy to heat energy. The temperature rises.

13. Exothermic Reactions • Almost immediately the hot reaction products start to lose heat to the surroundings and eventually they return to room temperature. 45o C 25o C Chemical energy becomes heat energy. The reaction mixture gets hotter. Eventually this heat is lost to the surroundings. It follows that reaction products have less chemical energy than the reactants had to start with.

14. Reactants have more chemical energy. reactants Some of this is lost as heat which spreads out into the room. Energy / kJ) Products now have less chemical energy than reactants. products Progress of reaction (time) Energy Level Diagram for an Exothermic Reaction

15. H is how much energy is given out reactants H is negative because the products have less energy than the reactants. H=negative Energy / kJ products Progress of reaction Energy Level Diagram for an Exothermic Reaction 2. Energy Level Diagram for an Exothermic Reaction

16. Energy / kJ) Progress of reaction reactants products Exothermic Reaction - Definition Exothermic reactions give out energy. There is a temperature rise and H is negative. His negative

17. Activity

18. Endothermic Reactions Endothermic reactions cause a decrease in temperature. • Endothermic chemical reactions are relatively rare. • A few reactions that give off gases are highly endothermic - get very cold. • Dissolving salts in water is another process that is often endothermic.

19. Heat energy taken in as the mixture returns to room temp. Ammonium nitrate Water Endothermic Reactions Endothermic reactions cause a decrease in temperature. Cools Starts 25°C Cools to 5°C Returns to 25°C

20. Endothermic Reactions • Extra energy is needed in order for endothermic reactions to occur. • This comes from the thermal energy of the reaction mixture which consequently gets colder. 5o C 25o C Reactants convert heat energy into chemical energy as they change into products. The temperature drops.

21. Endothermic Reactions • The cold reaction products start to gain heat from the surroundings and eventually return to room temperature. The reactants gain energy. 25o C 5o C 25o C This comes from the substances used in the reaction and the reaction gets cold. Eventually heat is absorbed from the surroundings and the mixture returns to room temperature. Overall the chemicals have gained energy.

22. This is how much energy is taken in products This is positive because the products have more energy than the reactants. H=+ Energy / kJ) reactants Progress of reaction Energy Level Diagram for an Endothermic Process

23. Energy / kJ Progress of reaction products reactants Endothermic Reaction Definition Endothermic reactions take in energy. There is a temperature drop and H is positive. H=+

24. Energy / kJ) Progress of reaction reactants products Exothermic Reaction - Definition Exothermic reactions give out energy. There is a temperature rise and H is negative. His negative

25. Activity

26. Endothermic Reactions Endothermic reactions cause a decrease in temperature. • Endothermic chemical reactions are relatively rare. • A few reactions that give off gases are highly endothermic - get very cold. • Dissolving salts in water is another process that is often endothermic.

27. Activity exo Are these endothermic or exothermic? • A red glow spread throughout the mixture and the temperature rose. • The mixture bubbled vigorously but the temperature dropped 150C. • Hydrazine and hydrogen peroxide react so explosively and powerfully that they are used to power rockets into space. • The decaying grass in the compost maker was considerably above the outside temperature. endo exo exo

28. Activity Energy / kJ Energy / kJ) reactants H=- Progress of reaction Progress of reaction products H=+ products reactants Sketch the two energy diagrams and label exothermic and endothermic as appropriate.

29. Energy in chemicals Energy needed Breaking chemical bonds • Most chemicals will decompose (break up) if we heat them strongly enough. • This indicates that breaking chemical bonds requires energy – is an endothermic process. Heat taken in Energy needed to overcome the bonding between the atoms

30. Energy in chemicals Energy given out Making chemical bonds • It is reasonable to assume that bond making will be the opposite of bond breaking • Energy will be given out in an exothermic process when bonds are formed. Heat given out Energy given out as bonds form between atoms

31. Energy given out as new bonds form Energy taken in as old bonds break Overall endothermic in this case Energy in chemicals H products reactants Changes to chemical bonds Endothermic Reactions • In most chemical reactions some existing bonds are broken (endothermic) • But new bonds are made (exothermic)

32. Energy taken in as old bonds break Overall exothermic – in this case Energy given out as new bonds form Energy in chemicals H reactants products Changes to chemical bonds Exothermic Reactions • Again some existing bonds are broken (endothermic) • And new bonds are formed (exothermic)

33. Exo Endo Bonds break Bond forming Bonds form Energy in chemicals Energy in chemicals Bonds break products H H reactants reactants products Summary – Bond Changes • Where the energy from bond forming exceeds that needed for bond breaking the reaction is exothermic. • Where the energy for bond breaking exceeds that from bond forming the reaction is endothermic.

34. Calorimetry • The value of ΔH can be determined experimentally by measuring the heat flow accompanying a reaction at constant pressure • When heat flows into or out of a substance, the temp., of the substance changes • Experimentally, we can determine heat flow associated with a chemical rxn by measuring the temp changes it produces • The measurement of heat flow is calorimetry; an apparatus that measures heat flow is a calorimeter

35. Standard Molar Enthalpy • Whenever we write an equation, to represent changes in matter, we usually represent numbers of moles of particles • for example Ca(s) + ½ O2(g)CaO(s) + 635.1 kJ • 1 molof calcium reacts with 0.5 molof oxygen to form 1 molCaO

36. The enthalpy change per mole of a substance undergoing a change is called the molar enthalpy and is represented by ΔHx, Where x indicates the type of change occuringi.e solution, vaporization, combustion etc. ΔHf= -635.1 kJ the enthalpy change may be either endothermic or exothermic Refer to sign convention: • ΔH exothermic + ΔH endothermic

37. To calculate an enthalpy change ΔH for some amount of substance other than a mole, you need to obtain the molar enthalpy value ΔHx from a reference source (Table), and then use the formula: ΔH = n ΔHx

38. Standard enthalpy change of reaction The amount of heat given out or absorbed during a chemical reaction depend on several factors • The nature of the reactants and products • The amount (or concentration) of reactants • The states of the reactants and products • The temperature of the reaction • The pressure at which the reaction is carried out

39. The standard enthalpy change of a reaction, ΔHθ Is defined as the enthalpy change when molar quantities of reactants in their normal states react to form products in their normal states under standard conditions of temperature and pressure. Standard pressure is 101.3 kPa (1atm) Standard temperature is 298 K (250C)

40. The actual amount of heat evolved in an exothermic rxn is normally measured in practice by using the heat given out to the surroundings to increase the temperature of a known mass of water.

41. Calculation of enthalpy changes Energy is defined a the ability to do work: that is to move a force through a distance. Energy = force x distance (J) (N x m) ΔHθvalues are normally given in kJ mol-1 • the actual amount of heat evolved in an exothermic reaction is normally measured by using the heat given out to the surroundings to increase the temp of a known mass of water