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Energy in Chemical & Physical Changes

Energy in Chemical & Physical Changes. Thermochemistry. Study of changes that accompany chemical reactions and phase changes The Universe is considered to be made of 2 parts: 1. System: part that contains the reaction or process 2. Surroundings: everything else. ENERGY.

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Energy in Chemical & Physical Changes

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  1. Energy in Chemical & Physical Changes

  2. Thermochemistry • Study of changes that accompany chemical reactions and phase changes • The Universe is considered to be made of 2 parts: 1. System: part that contains the reaction or process 2. Surroundings: everything else

  3. ENERGY • defined as the ability to do work or transfer heat energy. 2 types of energy • Potential Energy (PE): Energy at rest due to the position of an object; chemical potential energy is the energy stored in a substance’s bonds. • 2. Kinetic energy (KE): Energy of the motion of particles in a substance and is directly proportional to temperature. As temperature increases, KE also increases.

  4. Law of Conservation • Law of Conservation of Energy states that energy is neither created nor destroyed, just changed in form C8H18 + O2 H2O + CO2 + Energy • Stored PE converts to 25% work and 75% heat (ENERGY)

  5. Exothermic Reactions • HOT PACK • An exothermic reaction is when the system releases energy; heat flows out of a reaction and the surroundings get warmer. They have a NEGATIVEH. • H products < H reactants 4Fe + 3 O2 2 Fe2O3 + 1625 kJ OR 4Fe + 3 O2 2 Fe2O3H = - 1625 kJ

  6. Endothermic Reactions • COLD PACK • An endothermic reaction is when the system absorbs energy; heat flows into a reaction and the surroundings get cooler. They have a POSITIVEH • H products > Hreactants 27kJ + NH4NO3(s) NH4(aq)+1+NO3(aq)-1 OR NH4NO3(s) NH4(aq)+1+ NO3(aq)-1H = + 27 kJ

  7. Reaction Co-ordinates

  8. What is the difference between Temperature & Heat? Temperature • Instrument: thermometer • Units: Celsius, Fahrenheit, Kelvin • Definition: • A measure of the average kinetic energy of the molecules in a substance • A measure of the motions of the molecules • A measure of how hot or cold something is

  9. What is the difference between Temperature & Heat? Heat • Instrument: calorimeter • Units: calories, joules • Definition: • The total amount of energy in a substance. • A form of energy that is transferred between objects because one is warner than the other. • Heat transfer is always from hot to cold • Depends on 3 things: 1. amount of substance (mass) 2. Temperature change 3. type of material (specific heat)

  10. Units of Heat Energy • A calorie is defined as the amount of heat needed to raise the temperature of 1 g of water by 1 C 1 cal= 4.184 J • Food “Calories” are kilocalories. 1kcal = 1000 calories.

  11. Temperature ≠ Heat Greater Thermal Energy

  12. Specific Heat • Amount of heat required to raise the temperature of 1 g of a substance by 1 C • Different substances have different specific heats. Water has a specific heat of 4.184 J/gC. Iron(Fe) has a specific heat of .449 J/gC. Gold (Au) has a specific heat of .129 J/gC. • The higher the specific heat the more energy it takes to change its temperature.

  13. Calculating Heat c= specific heat q = heat in joules or galories m= mass T = change in temperature = Tf – Ti c= q_ mT

  14. Example • A 155 g sample of an unknown substance was heated from 25.0 C to 40.0 C. The substance absorbed 5696 J of energy. What is the specific heat?

  15. Example • How much heat is needed to change the temperature of 12.0 g of silver with a specific heat of 0.057 cal/g°C from 25.0°C to 83.0 °C?

  16. Measuring Heat in a Calorimeter • A coffee cup calorimeter measures heat at constant pressure; works on the premise that the amount of heat released in a reaction(-q) or physical change is equal to the amount of heat absorbed by the water(+q) - q = +q  • Rearrange the specific heat equation: q = m x c x T

  17. Example • A piece of unknown metal with mass 17.19 g is heated to an initial temperature of 92.50 °C and dropped into 25.00 g of water (with an initial temperature of 24.50 °C) in a calorimeter. The final temperature of the system is 30.05°C. What is the specific heat of the metal? Specific heat of water = 4.184 J/g°C

  18. Example • A 32.07 gram sample of vanadium was heated to 75.00 °C (its initial temperature). It was then dumped into a calorimeter. The initial temperature of the calorimeter’s water was 22.50 °C. After the metal was allowed to release all its heat to the calorimeter’s water, 26.30 °C was the final temperature. What mass of distilled water was in the calorimeter? • Specific heat of vanadium = .4886 J/gC Specific heat of water = 4.184 J/g°C

  19. Measuring Heat during Phases Changes

  20. Heat of Fusion/Solidification • Heat of fusion (Hfus ) is the heat energy required to melt one gram of a solid at its melting point For water, Hfus = 334 J/g q = Hfus x mass • Heat of solidification (Hsolid ) is the heat energy lost when one gram of a liquid freezes to a solid at its freezing point For water, Hsolid = -334 J/g q = Hsolid x mass

  21. Heat of Vaporization/Condensation • Heat of vaporization (Hvap) is the heat to vaporize one gram of a liquid at its normal boiling point For water, Hvap= 2260 J/g q = Hvap x mass • Heat of condensation (Hcond ) is the heat energy released when one gram of a liquid forms from its vapor For water, Hcond = -2260 J/g q = Hcond x mass

  22. Example • How much heat is needed to melt 500.0g of ice at 0 C?

  23. Example • How much heat is evolved when 1255 g of water condenses to a liquid at 100°C?

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