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Energy

Energy. Relationships in Chemical Reactions. The nature of Energy and Types of Energy. Energy – The capacity to do work Chemists define work as directed energy change resulting from a process Radiant Energy – energy produced from the sun

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Energy

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  1. Energy Relationships in Chemical Reactions

  2. The nature of Energy and Types of Energy • Energy – The capacity to do work • Chemists define work as directed energy change resulting from a process • Radiant Energy – energy produced from the sun • Thermal Energy – the energy associated with the random motion of atoms and molecules • Increase in thermal energy will increase the temperature but temperature does not define the amount of thermal energy • For example a cup of coffee at 70 degree verus a bathtub of water at 40 degrees. The bathtub will have more thermal energy due to the number of molecules

  3. Types of Energy • Chemical Energy – the stored energy within the structural units of the substance • Kinetic Energy – energy produce by a moving object • Potential Energy – the amount on energy an object possibly has due to the composition or position

  4. Conservation of Energy • It is deducted that nothing can be created or destroyed by the law of conservation of mass • Therefore energy can be transfer into a different type of energy • THE TOTAL QUANTITY OF ENERGY IN THE UNIVERSE IS ASSUMED CONSTANT

  5. Energy Changes in Chemical Reactions • Almost all chemical reactions absorb or produce (release) energy • Heat – the transfer of thermal energy between two bodies at different temperatures • Welcome to Thermochemistry

  6. Thermochemistry • To analyze the energy changes associated with chemical reactions we must define the system in relationship to the rest of the universe • Systems – open, closed or isolated

  7. Systems • Open – can exchange mass and energy with its surroundings (usually in the form of Heat) • Closed – transfer of energy but not mass • Isolated – does not allow transfer of mass or energy

  8. Introduction to Thermodynamics • The study of interconversion of heat and other kinds of energy • Comparing the State of the System • V = Vf - Vi

  9. First Law of Thermodynamics • Energy can be converted from one form to another but cannot be created or destroyed • The change of internal energy in denoted by U • U = Uf - Ui

  10. Example of 1st Law • Consider the reaction • S(s) + O2(g) SO2(g) • U = U(products) – U(reactants) • = energy content of 1 moles ofSO2(g) – energy content of [1moleS(s) + 1moleO2(g) ] • The energy of the products is less than that of the reactants and U is negative • Therefore the reaction gives off heat

  11. Interpreting the release of heat • The release of heat is a transfer of energy to the surrounding universe • Based on the 1st law the total energy of the universe is not changed • Therefore the total sum of the energy changes must be zero • U sys + U surr = 0

  12. Common Chemical Equation for 1st Law • U = q + w • U of the system is the sum of heat exchange q between the system and the work done on the system w. • Positive q – endothermic (heat absorbed by the system) • Negative q - exothermic (heat released into the surroundings) • Positive w – work done by the system on the surroundings • Negative w – work done on the system by the surroundings

  13. Work equation • W = -P V • Example • A certain gas expands in volume from 2L to 6L at constant temperature. Calculate the work done by the gas if it expands against a pressure of 1.2 atm

  14. Converting to Joules • In the equation for work the units will be L*atm • To convert to Joules

  15. Work problem • A gas expands from 264ml to 971ml at a constant temperature. Calculate the work done in joules by the gas if it expands against a pressure of 4atm

  16. Heat (q) • Energy can be gained by adding it directly to the system • Bunsen burner • Stirring

  17. Heat Problem • The work done when a gas is compressed in a cylinder is 384J. During this process there is a 152J heat transfer to the surroundings . Calculate the energy change for this system • U = q + w • Hint: must figure out signs of q and w

  18. Enthalpy of Chemical Reactions • Enthalpy of a reaction or energy change of a reaction • H, is the amount of energy or heat absorbed in a reaction • H is calculated in a pressure constant system • H = H(products) – H(reactants)

  19. H2O (l) Heat absorbed by the system from the surroundings H = 6.01 kJ/mol CH4 + 2O2 H2O (s) Heat given off by the system from the surroundings H = -890.4 kJ/mol CO2 + 2H2O

  20. Thermochemical Relationship between U and H • U = H - P V Consider the reaction between carbon monoxide and oxygen (write out on your paper) • U = H - n(RT)

  21. Calorimetry • Measurement of heat changes • Specific Heat (s) – the amount of heat required to raise the temperature of one gram of the substance by one degree Celcius • Heat Capacity (C) – the amount of heat required to raise the temperature of a given quantity of the substance by one degree Celcius

  22. C=ms • What is the heat capacity of 60g of water is the specific heat is 4.184J/g*C

  23. Calculating the amount of heat lost or gained (q) • q = ms T or q = C T • Reminder – • Pos (q) = Endothermic • Neg (q) = Exothermic

  24. Constant Volume Calorimetry • q sym = q cal + q rxn • = O • q cal = Ccal T Pg.192

  25. Example: It is known that 1g of benzoic acid releases 26.42kJ of heat. If the temp rises 4.673 then what is the heat capacity of the calorimeter

  26. Pressure Constant Calorimetry • Much simpler device to determine heat changes for non-combustion reactions • Can use two Styrofoam cups to contain heat for measurement

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