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IB Chemistry HL1

IB Chemistry HL1. Grade 11 : Unit 6- Energetics IB Topic 5. System and Surroundings. The system is the name we give the sample or reaction vessel of interest. The surroundings are everything else in the universe.

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IB Chemistry HL1

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  1. IB Chemistry HL1 Grade 11: Unit 6- Energetics IB Topic 5 2/2013

  2. System and Surroundings • The system is the name we give the sample or reaction vessel of interest. • The surroundings are everything else in the universe. • When a chemical change happens in an open system matter and energy can be exchanged between the system and the surroundings. • In a closed system only energy can be exchanged with the surroundings. 2/2013

  3. System and Surroundings 2/2013

  4. Energy • Energy is defined as the ability to do work. Energy is often converted from one form to another during physical and chemical changes. • Thermochemistry is the study of energy changes associated with chemical reactions. • Chemical energy is the energy stored in chemical bonds. It is a type of potential energy. 2/2013

  5. Thermochemistry • Most reactions absorb or evolve energy usually in the form of heat but chemical reactions can also produce light, electricity and mechanical energy – used to do work. • Energy is measured in joules, J. • 1000 J = 1 kJ. • Physical changes like change of state/phase also have heat energy changes. 2/2013

  6. Exothermic and Endothermic Reactions • 5.1.1 Define the terms exothermic reaction, endothermic reaction and standard enthalpy change of reaction (ΔH°). • Standard enthalpy change is the heat energy transferred under standard conditions—pressure 101.3 kPa, temperature 298 K. Only ΔH can be measured, not H for the initial or final state of a system. 2/2013

  7. Exothermic and Endothermic Reactions • 5.1.3 Apply the relationship between temperature change, enthalpy change and the classification of a reaction as endothermic or exothermic. • 5.1.4 Deduce, from an enthalpy level diagram, the relative stabilities of reactant and products, and the sign of the enthalpy change for the reaction. 2/2013

  8. Enthalpy • Enthalpy is the total energy of a system, some of which is stored as chemical potential energy in the chemical bonds. • Enthalpy is given the symbol H. • Enthalpy is also known as the heat content of a system. • We cannot measure the enthalpy content of a system but we can measure changes in it. 2/2013

  9. Enthalpy Change • In chemical reactions, bonds are broken and made, but the energy absorbed breaking bonds is almost never exactly equal to that released in making new bonds. • All reactions are accompanied by a change in the potential energy of the bonds and hence an ENTHALPY CHANGE. • There is no “absolute zero” for enthalpy so absolute enthalpies cannot be measured only the change in enthalpy that occurs during a reaction. 2/2013

  10. Enthalpy Change • The enthalpy change of a reaction is given the symbol ∆H. • ∆H is the difference in the enthalpy between the products and the reactants. • ∆H = H(products) - H(reactants) when at constant pressure. • Enthalpy level diagrams are used to show the change in enthalpy of a system during a change. 2/2013

  11. Enthalpy Level Diagram 2/2013

  12. Enthalpy Level Diagram 2/2013

  13. Endothermic Reaction • An endothermic reaction is one where energy is transferred from the surroundings to the system. • If energy is absorbed during a reaction then the enthalpy of the products will be higher than that of the reactants. • This means the enthalpy change will have a positive sign. • This reaction will either get cooler or heat will need to be supplied - temperature decreases. 2/2013

  14. Exothermic Reaction • An exothermic reaction is one where energy is transferred from the system to the surroundings. • If energy is released during a reaction then the enthalpy of the products will be lower than that of the reactants. • This means the enthalpy change will have a negative sign. • This reaction will “feel” hotter - temperature increases. 2/2013

  15. Stability • Reactions in chemistry tend towards products that are more stable. • Stability increases as energy decreases so exothermic reactions increase the stability of the substances. • This is why most chemical reactions that occur in nature are exothermic. • Endothermic reactions usually need “help” in the form of energy to allow them to occur. 2/2013

  16. Stability • Remember that lower energy is more stable. We can compare this to standing on top of a high building where you have more potential energy than someone on the ground. • If you fall to the ground you lose some of that potential energy you had on the roof but you are now more stable. 2/2013

  17. Exothermic and Endothermic Reactions • 5.1.2 State that combustion and neutralization are exothermic processes. 2/2013

  18. Combustion • Combustion is the scientific word for burning. • Most hydrocarbons burn easily in excess oxygen. • When they burn they produce carbon dioxide and water. • Ex. CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) • Combustion produces lots of heat which is transferred to the surroundings so it is very EXOthermic. 2/2013

  19. Combustion • When 1 mole of a substance is burned the energy released is called the Enthalpy Change of Combustion, ΔHc°. • The ° sign indicates this was measured under standard conditions in a controlled environment. • Standard conditions for thermochemistry experiments are T = 298K and P = 101.3kPa (= 1 atm). 2/2013

  20. Neutralization • Neutralization reactions involved acids and bases. • If an acid and a base react completely the resulting solution will be pH neutral. • The products are a salt (ionic compound) + water. • Ex. NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l) • When weak acids and bases are involved such as NaHCO3 then CO2(g) may be another product. • These reactions release energy to the surroundings so they are EXOthermic. 2/2013

  21. Negative ΔH • As combustion and neutralization reactions are always exothermic the enthalpy changes will always have negative values. • You can find many ΔH values in scientific literature ex. your textbook, the books in the classroom and online. 2/2013

  22. Spontaneous Reactions • A spontaneous reaction is one that occurs when the reactants are mixed without the need to be heated or have some other outside influence. • Most spontaneous reactions are EXOthermic but there are some spontaneous endothermic reactions ex. Dissolving NH4Cl in water. 2/2013

  23. Summary of Enthalpy Changes 2/2013

  24. Enthalpy Changes • In an exothermic reaction the products are more stable than the reactants so the bonds made are stronger than the bonds broken. • In an endothermic reaction the products are less stable than the reactants so the bonds made are weaker than the bonds broken. 2/2013

  25. Enthalpy Changes • Enthalpy changes are usually written alongside the chemical equation for the process with a positive or negative sign. • State symbols are VERY IMPORTANT as changes of state have their own enthalpy change values. • Enthalpy changes are usually reported per mole so the units are kJ mol-1. • If this is not the case then just kJ is used. 2/2013

  26. Enthalpy Changes • For example: • 2NaHCO3(s)  Na2CO3(s) + H2O(l) + CO2(g) • H = +91.6 kJ mol-1 • The + sign indicates it’s an endothermic reaction. 2/2013

  27. Standard Conditions • To compare enthalpy changes conditions must be the same. • The thermochemical standard conditions are • Temperature = 25°C = 298K (this is room temp) • Pressure = 1 atm = 101.3kPa • Solutions have a concentration of 1 mol dm-3 • Standard conditions are sometimes indicated by the symbol  or °: H orΔH° • Sometimes the temperature is included too:H 298. • The values in bold are the SI units. 2/2013

  28. Learning Check: Do NOW • Write the equation for the formation of chlorine oxide, Cl2O from its elements • What bonds are broken and what bonds are made in this process? • Do the processes in 2. absorb or release energy? • What is an enthalpy change? • In this reaction the bonds made are less strong than those broken. Will the enthalpy change be positive or negative? • Will this be an exothermic or endothermic reaction? 2/2013

  29. Calculation of Enthalpy Changes • 5.2.1 Calculate the heat energy change when the temperature of a pure substance is changed. • Students should be able to calculate the heat energy change for a substance given the mass, specific heat capacity and temperature change using q = mcΔT. • 5.2.2 Design suitable experimental procedures for measuring the heat energy changes of reactions. Students should consider reactions in aqueous solution and combustion reactions. 2/2013

  30. Calculation of enthalpy changes • 5.2.3 Calculate the enthalpy change for a reaction using experimental data on temperature changes, quantities of reactants and mass of water. • 5.2.4 Evaluate the results of experiments to determine enthalpy changes. Students should be aware of the assumptions made and errors due to heat loss. 2/2013

  31. Temperature • Temperature is a measure of the average kinetic energy of the particles in a system. • Units are K or °C. • Note: there is no ° sign used with K! • Heat is a measure of the total energy in a substance. • T°C + 273 = TK • TK – 273 = T°C 2/2013

  32. Specific Heat Capcity • Some substances will conduct heat and therefore change temperature more easily than others. • Ex. A metal pan on a stove will become very hot before the water inside it does. • A measure of how easily something changes temperature is called SPECIFIC HEAT CAPACITY. 2/2013

  33. Specific Heat Capacity • Specific heat capacity is the amount of heat energy required to increase 1 g of a substance by 1K or 1°C. • It is used in the following equation: • q = m x c x T where q is heat energy, m is mass, c is specific heat capacity and T is change in temperature. 2/2013

  34. Specific Heat Capacity • Units of specific heat capacity (c) are J g-1K-1 • Units of heat energy (q) are J or kJ • Units of mass (m) are g or kg • Units of temperature (T) are K or °C 2/2013

  35. Calorimetry • Calorimetry is a method used to measure the enthalpy associated with a particular change. • The temperature change of a liquid is measured inside a well insulated container called a calorimeter. • Often a styrofoam cup is used as it has a very low heat capacity and is a good insulator. 2/2013

  36. Measuring Energy Changes • To measure the enthalpy change of a reaction that occurs in solution you can carry it out in a styrofoam (polystyrene) cup and monitor the temperature during the reaction. • Styrofoam is a good insulator so the amount of heat lost to the surroundings will be reduced. 2/2013

  37. Measuring Energy Changes • Burning substances in a “bomb calorimeter” measures the temperature change in the water surrounding the burning item. • This system is also well insulated to try and reduce heat loss. • There may be some losses from incomplete combustion. 2/2013

  38. Calorimetry • If calorimeters made of other materials are used then the heat absorbed by the calorimeter must be added to that absorbed by the liquid: • Heat absorbed = (mcT)liquid + (mcT)calorimeter • Calorimetry assumes no heat is transferred to or from the surroundings so they must be well insulated. • However this is hard to achieve and is a major source of error in high school labs. 2/2013

  39. Calorimetry • Once you have calculated the energy released from a process then you can calculate how many kJ of energy were released per mole of the reactant. This is usually referred to as ΔH or molar enthalpy in a test question. • ΔH = q/n Where n represents the number of moles of the reactant that is reacting/burning. 2/2013

  40. Sample Problems Specific heat, c of liquid water = 4.18 kJ dm-3K-1 = 4.18J g-1K-1 • How much heat energy is required to increase the temperature of 10 g of nickel (c = 440 J g-1K-1) from 50°C to 70°C? • The enthalpy of combustion of ethanol (C2H5OH) is 1370 kJ mol-1. How much heat is released when 0.200 mol undergo complete combustion? 2/2013

  41. Sample Problems 3. H2(g) + 1/2 O2(g)  H2O(l) ∆H for the reaction above is -286 kJ mol-1. What mass of oxygen must be consumed to produce 1144 kJ of energy? 4. Calculate the molar enthalpy change when excess zinc is added to 50 cm3 of a 1 mol dm-3 solution of CuSO4. The temperature increases from 20°C to 70°C when the zinc is added. Assume the solution has the same density as water = 1.00 g cm-3 2/2013

  42. Bond Enthalpies • 5.4.1 Define the term average bond enthalpy. • 5.4.2 Explain, in terms of average bond enthalpies, why some reactions are exothermic and others are endothermic. 2/2013

  43. Bond Enthalpies • All chemical reactions involve the making and breaking of bonds. • The bonds in the reactants are broken which absorbs energy so this is an endothermic process. • The bonds in the products form which releases energy so this is an exothermic process. 2/2013

  44. Bond Enthalpy • Bond enthalpy is defined as the energy needed to break one mole of bonds in gaseous molecules under standard conditions. • Ex. ½ H2(g) + ½ Cl2(g)  H(g) + Cl(g) • Breaking a bond is endothermic so these values are always positive. • Bond enthalpies depend on the rest of the molecule so values are usually averages. 2/2013

  45. Bond Enthalpy • The energy released when a bond is made is the same value as the bond enthalpy but with a negative sign. • A higher bond enthalpy indicates a stronger bond. 2/2013

  46. Bond Enthalpy • If the bonds being broken (bonds in the reactants) are weaker than those being made (bonds in products) then the reaction will be exothermic. • If the bonds being broken are stronger than those being made then the reaction will be endothermic. 2/2013

  47. Bond Enthalpies • As bond enthalpies are averages and they are only for gases they are not the most accurate way to calculate an enthalpy change but they are usually within about 10% of more accurate values and are a useful tool. • ∆Hºreaction= ∑BEbonds broken- ∑BEbonds made Where BE stands for bond enthalpy 2/2013

  48. Bond Enthalpy • Bond enthalpy values are given in the data booklet. A sample is shown here. When using these values be careful to check if bonds are single or double etc. 2/2013

  49. Bond Enthalpy Calculations • When water is formed from its elements, what bonds are broken and formed? What is the enthalpy change predicted by bond enthalpies? • 2H2(g) + O2(g)  2H2O(l) • Bonds broken: 2(H-H), 1 (O=O) • Bonds formed: 4 (O-H) • ΔH = Σbonds broken – Σbonds formed 2/2013

  50. Bond Enthalpy Calculations • = (2(436) + 498 ) – (4(464)) • = (872 + 498) – (1856) • =1370 – 1856 = -486 kJ • What does this sign tell you about the reaction? • Does this make sense? 2/2013

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