Exp 11 Enthalpy
- 1. Enthalpy of Reaction General Chemistry Labs
- 2. Theory • Change in enthalpy, ΔH, is the measurement of the energy flow in a reaction • Chemical reactions absorb or release energy, in the form of heat • Reactions – Exothermic (heat is released, Δ H < 0) – Endothermic (heat is absorbed, Δ H > 0)
- 3. Calorimetry • The absorption or release of heat can be measured by monitoring the change in the temperature of the surroundings. • In order to determine the total amount of energy needed for the reaction, the heat used or released from the reaction must be contained within the reaction solution.
- 4. Calorimeter • A calorimeter is used to prevent the escape of heat. • In this lab, the calorimeter being used is a Styrofoam cup that sets in a glass beaker and has a cardboard lid.
- 5. Determination of heat (q) from Δ T • The relationship between the ΔT and heat is dependent on the properties of the solution • Some solutions will increase in temperature more readily than others even through they are absorbing the same amount of heat. • The specific heat for a substance is a constant that describes how readily the temperature will increase when heat is absorbed. • s = specific heat of water, units of 𝐽 (𝑔∗°𝐶)
- 6. q = msΔT • q = heat, units of Joules • m = mass, units of grams • s = specific heat of water, units of 𝐽 (𝑔∗°𝐶) • Δ T = change in temperature, units of °C
- 7. q and ΔH • q is the amount of heat lost or gained in the total reaction. – The units are Joules and it is an extensive property, – not dependent on the amount of the substance. • Δ H is an intensive property that is typically expressed as the amount of heat per mole, or kJ/mol.
- 8. Relationship of q and Δ H • Therefore, q # moles = Δ H • WATCH the units!!! • Δ H will be expressed as a negative or positive dependent on the direction of heat flow.
- 9. Exothermic or Endothermic • The positive or negative sign shows the direction of the heat flow in a reaction • Negative: heat produced by the reaction, and heats surrounding areas – Exothermic • Positive: heat is consumed by the reaction, and therefore cools the surrounding areas. – Endothermic • The (+) or (-) signs are assigned based on the direction of heat flow not by the mathematical equation, Δ H = 𝑞 𝑚𝑜𝑙 = 𝑚𝑠Δ 𝑇 𝑚𝑜𝑙
- 10. Calculating Enthalpy of Reaction • Indirect Method ΔHrxn = SnΔHf(products) - SmΔHf(reactants) • Δ Hf for most common substances are found in tables
- 11. Hess’s Law – The enthalpy of a reaction does not depend on number of steps involved – If the enthalpy of a set of reactions is known, • combining the enthalpy of these reactions then enthalpy of a reaction of interest can be obtained • Rearrange given reactions to get the overall reactions. Whatever we do to the reaction we also do to the enthalpy
- 12. Reaction Δ H (kJ/mole) C(gr) + O2(g) → CO2(g), Δ H0 rxn1 = -393.5 S(rh) + O2(g) → SO2(g), Δ H0 rxn2= -296.4 CS2(l) + 3O2(g) → CO2(g) + 2SO2(g) Δ H0 rxn3= -1073.6 Determine the enthalpy of the reaction shown using the three reactions given. C(graphite) + 2 S(rhombic) → CS2(l) Hess’s Law : Example
- 13. C(gr) + 2S(rh) CS2(l) C(gr) + O2(g) → CO2(g) S(rh) + O2(g) → SO2(g) CS2(l) + 3O2(g) → CO2(g) + 2SO2(g) Therefore, rewrite equations C(gr) + O2(g) → CO2(g) 2 S(rh) + 2 O2(g) → 2 SO2(g) 2 SO2(g) + CO2(g) → CS2(l) + 3 O2(g) 2 x -1 x
- 14. C(gr) + 2S(rh) CS2(l) Compounds in equal quantities on both sides of the reaction arrow cancel out C(gr) + O2(g) → CO2(g) 2S(rh) + 2O2(g) → 2SO2(g) 2SO2(g) + CO2(g) → CS2(l) + 3O2(g) C(gr) + 2S(rh) → CS2(l) ΔH0 rxn = Δ H0 rxn1 + 2 Δ H0 rxn2 - Δ H0 rxn3 = 87.3 kJ/mol
- 15. Procedure Put cardboard square on top and place temperature probe through hole in cardboard. Calorimeter
- 16. Finding Enthalpy • Mix reactants • Measure temperature change • Calculate heat change • Calculate change in enthalpy (ΔH)
- 17. Calculating Change in Temperature • Include the graphs of temperature vs. time data for all three reactions • Δ T – Initial temperature: Right before solutions are mixed – Final temperature: Maximum temperature in data • Let experiment continue until the temperature is no longer increasing
- 18. Experimental Results • Enthalpy values for reactions 1, 2, and 3 • Two experimental values for reaction 3 – Value from running the reaction – Using Hess’ law from reactions 1 & 2 • Accepted values are from pre-lab
- 19. Data Analysis • Calculate Δ H0 for each reaction, in terms of kJ/mol for one reactant – Use the volume and molarity of one reactant to calculate the number of moles that were actually used in the experiment. – Each of the 3 reactions will have different Δ H0 values
- 20. Data Analysis • Calculate % error in Δ H0 for reaction 3 for both experimental values of Δ H0 – Do not average the two enthalpy values for reaction 3 • The accepted value of Δ H0 for reaction 3 was obtained in the pre-lab exercises
- 21. Practice Question #1 • Consider two metals A and B. Each metal has a mass of 250 g and an initial temperature of 25°C. The specific heat of B is larger than that of A. • Under the same heating conditions, which metal would take longer to reach a temperature of 26°C?
- 22. Practice Question #2 • A 15.0-g sample of nickel metal is heated to 100.00 oC and dropped into 55.0 g of water, initially at 23.0oC. • Assuming that all the heat lost by the nickel is absorbed by the water, calculate the final temperature of the nickel and water. • (The specific heat of nickel is 0.444 J/oC g)
- 23. Practice Question #3 Given the following data: 2O3(g) 3O2(g) Δ H0 = -427kJ O2(g) 2O(g) Δ H0 = +495kJ NO(g) + O3(g) NO2(g) + O2(g) Δ H0 = -199kJ Calculate Δ H for the reaction NO(g) + O(g) NO2(g)