1. Energy changes in reactions Senior Chemistry R Slider

2. Chemical Energy- Enthalpy Kinetic energy is the energy of motion Particles that make up matter are constantly in motionParticles that make up matter are constantly in motion Potential energy is stored energy (no motion) Every bond between atoms and ions has stored energy within the bondsEvery bond between atoms and ions has stored energy within the bonds This stored energy is known as chemical potential energyThis stored energy is known as chemical potential energy Total energy = Kinetic energy +Potential Energy (known as Enthalpy or heat content– symbol H)

3. Change in Enthalpy (ΔH) Because it is difficult to measure the amount of energy in individual reactants or products, Chemists measure the change in enthalpy ( ΔH) that occurs when a reaction takes place. ΔH is the difference in the enthalpy of the products and the enthalpy of the reactants and is measured in kJ/mol (molar enthalpy) or sometimes kJ/g. Enthalpy changes result from bonds breaking and new bonds reforming. This breaking and reforming of bonds results in energy changes within the system which are measured as temperature changes This breaking and reforming of bonds results in energy changes within the system which are measured as temperature changes Note: Standard enthalpy changes are measured under standard temperature (298K) and pressure (101.3kPa)

4. Endothermic or Exothermic? Endothermic Some chemical reactions absorb heat from the surroundings so the products of the reaction contain more heat. The surroundings feel colder, ΔH = + Examples: photosynthesis, cold packs Exothermic Some chemical reactions release heat to the surroundings so the products have less heat than the reactants The surroundings feel hotter, ΔH = - Examples: combustion, neutralisation reactions

5. Heat exchange Another way to visualise ΔH is to think about energy being exchanged between the chemical system and the surroundings Source diagrams: http://www.chemhume.co.uk/ASCHEM/Unit%203/13%20Enthalpy/13%20Enthalpyc.htm http://www.chemhume.co.uk/ASCHEM/Unit%203/13%20Enthalpy/13%20Enthalpyc.htm Heat loss in a chemical system = Heat gain to the surroundings (Temperature increases) Heat loss in a chemical system = Heat gain to the surroundings (Temperature increases) Heat gain in a chemical system = Heat loss to the surroundings (Temperature decreases)

6. Respiration is exothermic – energy is released Source: http://image.slidesharecdn.com/cellular-respiration-23666/95/cellular-respiration-8-728.jpg?cb=1171448210 http://image.slidesharecdn.com/cellular-respiration-23666/95/cellular-respiration-8-728.jpg?cb=1171448210

7. Neutralisation is exothermic AcidBaseSaltWater Acid + base reactions are called neutralisation reactions because the low pH ( 7) of a base, results in a pH closer to neutral (7) These reactions all release heat and are exothermic. For example: HCl + NaOH  NaCl + H 2 0 + 58 kJ/mol

8. Combustion is exothermic Combustion is the process of burning Most often, the combustion of a material involves its combination with oxygen gas. Because combustion reactions release energy in the form of heat and light, they are exothermic. Fire in the Penola Forest (SA), Ash Wednesday 1983 Fire in the Penola Forest (SA), Ash Wednesday 1983 (Source: www.austehc.unimelb.edu.au/ fam/1611_image.html) www.austehc.unimelb.edu.au/ fam/1611_image.htmlwww.austehc.unimelb.edu.au/ fam/1611_image.html Example: The combustion of ethanol in oxygen: C 2 H 5 OH + 3O 2  2CO 2 + 3H 2 O ∆ H rxn = -1368 kJ/mol (exothermic)

9. Combustion: why exothermic? Energy is required to break bonds(endothermic)Energy is required to break bonds(endothermic) Energy is released when bonds form (exothermic)Energy is released when bonds form (exothermic) In summary: ∆H rxn = H (products) – H (reactants) ∆ H rxn = H (products) – H (reactants) This reaction releases more energy than it absorbs resulting in a negative value for ∆Hrxn This reaction releases more energy than it absorbs resulting in a negative value for ∆ Hrxn H reactants > H products Breaking bonds Absorbs energy Forming bonds Releases energy Energy required to break bonds Energy released when bonds formed In the simple example involving the combustion of methane: CH 4 + 2O 2  CO 2 + 2H 2 O ∆ H = -832kJ A thermochemical equation

10. Photosynthesis is endothermic – absorbs energy Source: http://tomatosphere.org/teachers/guide/images/photosynthesis-equation.jpg http://tomatosphere.org/teachers/guide/images/photosynthesis-equation.jpg

11. Energy changes in chemical reactions – Enthalpy diagrams Exothermic reactions The enthalpy of the reactants is higher than the products. Reactants ReactantsProducts Endothermic reactions The enthalpy of the reactants is lower than the products. Products Products Reactants Reactants Enthalpy (H) ∆ ∆ H is -ve ∆ ∆ H is +ve Heat released Heat absorbed Reaction progression

12. Examples The combustion of methane is exothermic and releases energy to the surroundings The decomposition of calcium carbonate is endothermic and absorbs energy from the surroundings Source diagrams: http://www.chemhume.co.uk/ASCHEM/Unit%203/13%20Enthalpy/13%20Enthalpyc.htm http://www.chemhume.co.uk/ASCHEM/Unit%203/13%20Enthalpy/13%20Enthalpyc.htm

13. Exercises 1.Classify each of the reactions below as endothermic or exothermic a)C 6 H 12 O 6(aq) + 6O 2(g)  6CO 2(g) + 6H 2 O (l) Δ H = -2801 kJ mol -1 b)6CO 2(g) + 6H 2 O (l)  C 6 H 12 O 6(aq) + 6O 2(g) Δ H = +2801 kJ mol -1 c)H 2(g) + 1/2 O 2(g)  H 2 O (l) + 286 kJmol -1 2.When a spark is introduced into a vessel containing a mixture of hydrogen and oxygen gases, they react explosively. a)What is the role of the spark in this reaction? b)What bonds are broken when this reaction is initiated? c)What bonds are generated when the products are formed? d)Is the reaction endothermic or exothermic? Justify your answer. 3.10.0g of ammonium nitrate is dissolved in 100cm 3 of water and the temperature of the solution decreases from 19.0 0 C to 10.50 0 C. a)Which is greater, the enthalpy of the reactants or the products? b)Is the reaction endothermic or exothermic? c)What chemical bonds were broken when the ammonium nitrate dissolved in water?

14. Exercises - Answers 1.a) exothermic b) endothermic c) exothermic 2.a)spark is to overcome the activation energy b) H-H and O=O covalent bonds broken c) O-H bonds are made during the reaction d) exothermic as energy is released 3.a) products>reactants b) endothermic c) ionic bonds between NH 4 + and NO 3 -

15. Measuring Enthalpy Changes q Recall that changes in temperature are related to changes in the amount of heat energy that is being absorbed or released in a system. This quantity of heat energy, q in joules (J), can be calculated using the following equation and often involves water being heated by reactions: q = m C ∆T Heat energy released or absorbed in Joules (J) Mass (g) Specific heat Capacity (4.18 J g -1 K -1 for water) Change in Temperature (final – initial) Example: What quantity of energy is required to raise the temperature of 0.5L of water from 20 0 C to 100 0 C? q = m C ∆T = 500g X 4.18 J g -1 K -1 X 80 0 C or K = 500g X 4.18 J g -1 K -1 X 80 0 C or K = 167,200 J = 167,200 J = 167 kJNB: 0 C = K - 273 = 167 kJNB: 0 C = K - 273

16. Heat Capacity Values Substance Specific Heat Capacity J K -1 g -1 Substance Specific Heat Capacity J K -1 g -1 Water4.18Ethanol2.41 Ethylene glycol 2.39Hexane2.26 50/50 water/E. glycol 2.86Chloroform0.96 Acetone2.17Aluminium0.90 Iron0.448Copper0.386 Specific heat capacity (C), is the amount of heat energy in Joules (J), required to raise the temperature of 1g of a substance by 1 Kelvin (K). Heat capacity is a measure of how much energy a substance can absorb without changing temperature. These values vary greatly among substances.

17. Practice Questions 1.A Bunsen burner was used to heat 100 cm 3 of water for 10 min. The temperature of the water increased from 15.0 0 C to 80.0 0 C. Determine the heat energy change of the water. (note: the density of water is 1.0 g cm -3 ) 2. The same Bunsen burner was used to heat a 500g block of copper for 10 minutes. Assuming the same amount of energy is transferred from the Bunsen burner to the block as in question 1, determine the highest temperature that the block could reach if the starting temperature was 15.0 0 C. 3.Explain the difference in the temperatures considering the same amount of energy was put into the water and the copper Solutions: 1)27.2kJ2) 156 0 C3) The water has a much higher heat capacity than the copper which means the water can absorb much more heat energy before changing temperature.

18. Calorimetry What is calorimetry Water can be used to measure the change in heat energy in a chemical reaction due to its ability to absorb heat. This is known as calorimetry. A calorimeter is a device that is used to measure the enthalpy change that occurs during a chemical reaction. Measuring heat in the laboratory If a known quantity of water is placed in the calorimeter and a reaction is carried out, the change in temperature due to the reaction is transferred to the water and is measured using a thermometer placed in a hole in the lid. The heat energy change is: q = m C ∆T The amount of heat lost or gained in the reaction is equal in size but opposite in sign to the amount of heat lost or gained by the water. Coffee cup calorimeter The simplest calorimeter makes use of two polystyrene (Styrofoam) cups, one inside the other with a lid on top. This minimises the amount of heat lost to the surroundings. This is the most significant source of error.

19. Calorimetry example When a student added excess Mg is to 100cm 3 of 2.00 mol dm -3 CuSO 4, the temperature rose from 20.0 o C to 65 o C. i) The energy change: Q = m x c x Δ T Q = 100 x 4.18 x (65 - 20) Q = 100 x 4.18 x 45 Q = 18810 J Q = 18.81 kJ Convert the energy calculated into KJ (divide by 1000) Q = 18.81 kJ Convert the energy calculated into KJ (divide by 1000) ii) Calculate the number of moles used: No moles = c x V No moles = c x V 1000 1000 No moles = 2 x 100 No moles = 2 x 100 1000 1000 No moles = 0.2 No moles = 0.2 iii) Calculate the amount of energy exchanged per mole, this is the enthalpy: Enthalpy = Energy Enthalpy = Energy moles moles Enthalpy = - 18.81 Enthalpy = - 18.81 0.2 0.2 Enthalpy = - 94.05 kJ mol -1 Enthalpy = - 94.05 kJ mol -1 iv) Finally write the equation with the enthalpy change: Mg (s) + CuSO 4(aq)  MgSO 4(aq) + Cu (s) Δ H = - 94.05 kjmol -1 It is exothermic, therefore negative

20. Enthalpy of solution In order for a solute to dissolve in a solvent, the solute particles and the solvent particles must separate, which involves energy changes. If one mole is dissolved, it is the molar enthalpy of solution In addition, solute particles interact with solvent particles, which also can result in energy changes. Overall: ΔH soln = ΔH 1 + ΔH 2 + ΔH 3 The energy profile diagrams on the right show net exothermic and endothermic dissolutions Question: How are these processes similar to a chemical reaction? How are they different? Diagrams Source: http://wps.prenhall.com/wps/media/objects/3312/3391718/blb1301.html http://wps.prenhall.com/wps/media/objects/3312/3391718/blb1301.html

21. Enthalpy of Combustion (ΔH c ) The diagram to the right shows a typical school laboratory setup to test the enthalpy of combustion of a fuel such as ethanol. 1.A known quantity of water is placed in a flask or beaker (calorimeter) 2.A thermometer is positioned with bulb near the middle of the volume of water 3.A known quantity of fuel, such as an alcohol (alkanol), is placed in the spirit burner 4.The initial temperature of the water is measured and recorded (T i ) 5.The wick on the spirit burner is lit, burning the fuel, and heating the water 6.When the temperature has risen an appreciable amount, the spirit burner is extinguished and the final temperature recorded (T f ) 7.The final quantity of fuel is measured and recorded Source: http://www.ausetute.com.au/heatcomb.html http://www.ausetute.com.au/heatcomb.html

22. Enthalpy of Combustion example initial water temperature (T i ) = 20 o Cintial mass burner + ethanol = 37.25 g final water temperature (T f )= 75 o Cfinal mass burner + ethanol = 35.50 g change in temperature = T f - T i = 55 o C mass ethanol used = 1.75g A student used the apparatus on the previous slide to determine the heat of combustion of ethanol. Their results for one trial are below: Use the results above to determine the enthalpy of combustion for ethanol in kJ/mol assuming 200 cm 3 of water is used in the calorimeter Also, note MM of ethanol is 46 g.mol -1. (Solution on the next slide)

23. Enthalpy of Combustion example 1.Calculate moles (n) of fuel used molecular mass (MM) of ethanol = 46.1g/mol mass ethanol used = 1.75g n = mass ÷ MM = 1.75 ÷ 46.1 = 0.0380 mol 2. Calculate energy required to change temperature of water energy = mass of water x specific heat capacity of water x change in water temp (mCΔT) energy = 200g x 4.184 JK -1 g -1 x 55 o C = 46024 J = 46.024 kJ 3. Calculate the heat of combustion of ethanol Assume all the heat produced from burning ethanol has gone into heating the water, ie, no heat has been wasted. 0.0380 mole ethanol produced 46.024 kJ of heat. Therefore 1 mole of ethanol would produce 46.024 kJ ÷ 0.0380 mol = 1211 kJ/mol The heat of combustion of ethanol is 1211 kJ/mol The experimentally determined value for the heat of combustion of ethanol is usually less than the accepted value of 1368 kJ/mol because some heat is always lost to the atmosphere and in heating the vessel.

24. Energy changes in reactions Compiled by: Robert Slider (updated 2016) Please share this resource with others