Honors Chem Unit 6

• Energy


  The ability to do work or produce heat
• Forms of energy


  • Potential Energy (PE)- due to composition or position
  • Kinetic Energy (KE)- energy in motion
• Law of conservation of energy


  Also known as the First Law of Thermodynamics, states that in any chemical reaction or physical process, energy can be converted from one form to another, but it is neither created nor destroyed
• Chemical potential energy


  Energy stored in a substance because of its composition, important to chemical reactions
• Heat


  Energy that is in the process of flowing from a warmer object to a cooler object, symbolized as q
• Calorie


  The amount of energy required to raise the temperature of one gram of water one degree Celsius. Food is measured in Calories, or 1000 calories (kilocalorie).
• Joule


  The SI unit of heat and energy, equivalent to 0.2390 calories
• Specific heat (c)


  The amount of heat required to raise 1 gram of a substance 1 °C. Some objects require more heat than others to raise their temperature.
• Calculating heat absorbed and released


  q = c × m × ΔT, where q = heat absorbed or released (J), c = specific heat of substance, m = mass of substance in grams (g), ΔT = change in temperature in Celsius (Tf - Ti)
• Heat, q


  Thermal energy transferred from a hotter system to a cooler system that are in contact
• Temperature


  Measure of the average kinetic energy of the atoms or molecules in the system
• Calorimeter


  Insulated device used for measuring the amount of heat absorbed or released in a chemical reaction or physical process
• Calorimetry scenario


  1. Put 125g of water into a foam-cup calorimeter with initial temp. 25.60°C

  2. Heat a 50.0 g sample of an unknown metal to 115°C

  3. Put the metal sample into the water

  4. Heat flows from the hot metal to the cooler water, and the temp. of the water rises

  5. The flow of heat stops only when the temp. of the metal and the water are equal

  6. Final temp. is 29.30°C

  7. Assuming no heat is lost to the surroundings, heat gained by the water = heat lost by the metal
• Enthalpy


  Heat content of a system at constant pressure
• Enthalpy (heat) of reaction

  Change in enthalpy during a reaction
• Exothermic reactions


  Enthalpy changes are always negative
• Exothermic


  • Heat is released to the surroundings
  • You feel an increase in Temperature
  • Energy is a product
  • Occurs when chemical bonds are formed
  • Enthalpy is negative (-ΔH)
  • ΔH < 0
• Endothermic reactions


  Enthalpy changes are always positive
• Endothermic


  • Heat is absorbed into the system
  • You feel a decrease in Temperature
  • Energy is a reactant
  • Occurs when chemical bonds are broken
  • Enthalpy change is positive (+ΔH)
  • ΔH is > 0
• Thermochemical equation


  Balanced chemical equation that includes the physical states of all reactants and products, and energy change
• Heat of combustion


  Enthalpy change for the complete burning of one mole of a substance
• Heat of vaporization


  Heat required to vaporize one mole of a liquid substance
• Heat of fusion


  Amount of heat required to melt one mole of a solid substance
• Bond Energy


  The energy required to break 1 mole of a chemical bond in the gas phase
• Bond Energy


  • Is energy stored in a bond
  • If negative, it represents the energy released when a bond is formed
  • If positive, it represents the energy required to break a bond
• Compounds with stronger bonds tend to be more chemically stable, and therefore less chemically reactive, than compounds with weaker bonds
• Average bond energies are useful in estimating the enthalpy change in a reaction
• Equation to Calculate the change in standard enthalpy of a reaction using bond energies


  ∆H rxn
• Energy is conserved, meaning it cannot be created or destroyed but only transferred from one form to another.
• Chemical reactions are exothermic if they release heat into their surroundings and endothermic if they absorb heat from their surroundings.
• The energy released during the combustion reaction is equal to the sum of the energies required to break all bonds
• The Law of Conservation of Energy states that energy can neither be created nor destroyed; rather, it changes forms.
• 1J is equal to 0.02390 cal
• To convert from J to small cal use 0.2390 cal divided by 1J
• To convert to J use 1J/0.2390cal or 4.184/ 1cal
• 1 cal is equal to 4.184J
• 1 Calorie is equal to 1kcal
• To convert Calorie to small calorie, multiply by 1000cal/1Cal and do opposite to get big Cal
• Lower specific heat means it gets hotter faster
• Hess's law


  if you add two or more chemical equations to get an overall equation, then you can also add the heat changes (delta H's) to get the overall heat change.