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Understanding Energy in Chemical and Physical Changes: Thermochemistry Basics

This guide covers key concepts in thermochemistry, including energy types, chemical reactions, heat transfer, and measurements. It differentiates between potential and kinetic energy while explaining the law of conservation of energy. The section on exothermic and endothermic reactions highlights the energy changes that occur during these processes. The guide also explores temperature and heat differences, specific heat calculations, and calorimetry for measuring heat transfer during reactions and phase changes. Essential for students and chemistry enthusiasts alike.

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Understanding Energy in Chemical and Physical Changes: Thermochemistry Basics

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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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