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IB DP1 Chemistry Energetics

IB DP1 Chemistry Energetics. Why do chemical reactions get hot or cold?. Topic 5: Energetics ( 8 hours). 5.1 Exothermic and endothermic reactions 5.1.1 Define the terms exothermic reaction, endothermic reaction and standard enthalpy change of reaction ( ∆HÖ ) .

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IB DP1 Chemistry Energetics

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  1. IBDP1 ChemistryEnergetics Why do chemical reactions get hot or cold?

  2. Topic 5: Energetics (8hours) 5.1 Exothermic and endothermic reactions 5.1.1 Define the terms exothermic reaction, endothermic reaction and standard enthalpy change of reaction ( ∆HÖ ) . 5.1.2 State that combustion and neutralization are exothermic processes. 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 reactants and products, and the sign of the enthalpy change for the reaction. 5.2 Calculation of enthalpy changes 5.2.1 Calculate the heat energy change when the temperature of a pure substance is changed. 5.2.2 Design suitable experimental procedures for measuring the heat energy changes of reactions. 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. 5.3 Hess’s law 5.3.1 Determine the enthalpy change of a reaction that is the sum of two or three reactions with known enthalpy changes. 5.4 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.

  3. Thermochemistry • Study of energy changes during chemical reactions • Heat, light, mechanical energy …..

  4. mix ammonium nitrate and water burn ethanol

  5. See Think Wonder

  6. Why do chemical reactions get hot (or cold)? • What is the difference between energy and enthalpy? • What is the difference between temperature and heat? • How much heat does it take to increase the temperature of a substance? • What is the difference between enthalpy change and standard enthalpy change? • How is calorimetry used to measure enthalpy changes? • How is standard enthalpy change calculated from a temperature change?

  7. Exothermic and endothermic reactions Ammonium nitrate and water Iron and oxygen

  8. Cold pack • The temperature of the cold pack decreases during the reaction • The temperature of the cold pack increases during the reaction • The cold pack transfers cold to the person touching it • Heat is taken in from the person to the cold pack • The cold pack has less energy and lower enthalpy after it is used

  9. Exothermic reactions examples NaOH(s) + H2O NaOH(aq) + heat Exothermic HCl + NaOHNaCl + H2O + heat Neutralisation Wood + O2 CO2 + H2O + heat Combustion

  10. Endothermic reaction example Ba(OH)2(s) + 2 NH4SCN(s) + heat  Ba2+(aq) + 2 SCN-(aq) + 2 H2O(l) + NH3(aq) Endothermic

  11. Endothermic or exothermic? In chemical reactions bonds break and form different amounts of energy are in the bonds before and after the reaction • Exothermic reaction: less energy is in the bonds after the reaction: heat is produced • Endothermicreaction = energy is needed

  12. Distribution of speeds of particles number of particles speed

  13. Temperature- average KE per particle higher average speed  higher temperature more particles at same average speed  same temperature

  14. Distribution of speeds

  15. Enthalpy, H • Energy stored in chemical bonds of reactants (in Joules) • PE and KE of particles + energy to make space for substance Image: https://en.wikipedia.org/wiki/File:Ammonium_Nitrate.jpg

  16. Exothermic reaction CH4 + 2 O2  CO2 + 2H2O + heat Energy rich Energy poor DH = (Energy poor) – (Energy rich) negative value Exothermic reactions: DH < O  more stable products Endothermic reactions DH > O  more reactive products.

  17. Exothermic reactions Enthalpy, J reactants ∆H products reaction coordinate

  18. Endothermic reactions Enthalpy, J products ∆H reactants reaction coordinate

  19. Standard enthalpy change of a reaction ∆H° to compare reactions in kJ/mol measured at 298K and 1atm

  20. Calculating standard enthalpy changes 2 Mg + O2 2 MgODH = -1202 kJ/mol Exothermic • The amount of energy released when 0.6 g of Mg is burnt? Mg m 0.6 g M 24.3 g/mol n 0.025 mol • 1202*0.025 = 30 kJ

  21. Calculating standard enthalpy change What is the enthalpy change in kJ per mole if 45kJ are given out when 0.8g of methane is burned? What is the enthalpy change in kJ per mole if 1.6g of methanol is used to heat 200mL water from 20C to 38C?

  22. Using temperature to calculate ∆Hᶱ Heat energy = mass x specific heat capacity x temperature change Q = mc∆T

  23. How much energy is in a cracker? • 1408kJ per 100grams • 7 grams per cracker • How much will a cracker increase the temperature of 100mL of water when we burn it?

  24. Calculating uncertainties How do you calculate uncertainties when: • adding or subtracting variables? • multiplying or dividing variables?

  25. Calculate the standard enthalpy change for the reaction between hydrochloric acid and sodium hydroxide

  26. To calculate the uncertainty in the standard enthalpy for the reaction between zinc and copper sulphate... • absolute uncertainty in temperature change = • percentage uncertainty in temperature change = • percentage uncertainty in mass of liquid = • percentage uncertainty in specific heat capacity = • pecentage uncertainty in enthalpy = • percentage uncertainty in number of moles = • total percentage uncertainty = • total absolute uncertainty in standard enthalpy =

  27. Neutralization- acid-base reaction Write word and chemical equations for the following reactions: • hydrochloric acid + sodium hydroxide • sulphuric acid + sodium hydroxide • nitric acid + potassium hydroxide

  28. Standard enthalpy of neutralization H+ (aq) + OH-(aq) H2O(l) standard enthalpy change of neutralization for a strong acid and base is -55.90 kJ/mol

  29. Bond enthalpy The enthalpy change when one mole of bonds is formed in the gaseous state: X (g) + Y(g) X-Y(g) Forming bonds is exothermic (negative ∆H) Breaking bonds is endothermic (positive ∆H)

  30. Calculate a theoretical standard enthalpy of combustion of methane

  31. Enthalpy of combustion Image: https://en.wikipedia.org/wiki/Burning

  32. How does the theoretical standard enthalpy of combustion depend on the number of carbon atoms for the alkanes? • Compare your answers with experimental values.

  33. Standard enthalpy of combustion of alkanes

  34. Ammonia, NH3 • Draw a Lewis diagram of the molecule • State the bond angles and shape • State whether the N-H bonds are polar, and whether the molecule is polar • Calculate the standard enthalpy of formation from data booklet data

  35. Standard enthalpy change of formation ΔHfᶱ the standard enthalpy change when a compound is formed from its elements • CH4 -74.4kJmol-1 • CO2 -393.5kJ/mol-1 • H2O -285.8kJ/mol-1

  36. Formation of ammonia N2+ 3 H2 2 NH3 N N triple bond944kJmol-1 H-H 436 kJmol-1 N-H 388 kJmol-1 DH = (cost of bond breaking) – (gained for bond forming) = ( 944 + 3*436) – (2*3*388) = -76 kJmol-1 Exothermic

  37. Hess’s Law • conservation of energy applied to chemistry • total enthalpy change is the same whatever the route of a chemical reaction

  38. Combustion of carbonto form Carbondioxide • C + ½ O2 CO DH1= -110kJmol-1 • CO +½ O2 CO2DH2= -283kJmol-1 • C + O2 CO2 DH3 = DH1+DH2=

  39. Plan investigation…

  40. Links • Ionic bonding http://www.teachersdomain.org/asset/lsps07_int_ionicbonding/ • Covalent bonding http://www.teachersdomain.org/asset/lsps07_int_covalentbond/

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