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Energy - PowerPoint PPT Presentation


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Energy. Energy. The capacity to do work Many types of energy and one type can be transformed into other types. Example: Law of Conservation of Energy: Energy cannot be created nor destroyed, only transformed. Energy. Potential Energy: Energy stored in an object due to its position

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energy1
Energy
  • The capacity to do work
  • Many types of energy and one type can be transformed into other types.
    • Example:
  • Law of Conservation of Energy:
    • Energy cannot be created nor destroyed, only transformed
energy2
Energy
  • Potential Energy:
    • Energy stored in an object due to its position
  • Kinetic Energy:
    • Energy of motion
potential energy
Potential Energy
  • Chemical Potential Energy
  • Gravitational Potential Energy
  • Elastic Potential Energy
kinetic energy
Kinetic Energy
  • Thermal Kinetic Energy
  • Sound Kinetic Energy
  • Mechanical Kinetic Energy
  • Electrical Kinetic Energy
  • Radiant Kinetic Energy
heat temperature
Heat & Temperature
  • Heat is not the same thing as temperature!!
    • The addition of heat to a system often causes an increase in temperature – and vice versa – but the two are not synonymous
heat temperature1
Heat & Temperature
  • Temperature
    • A measure of the average kinetic energy of the particles in a sample
  • Heat
    • A measure of the total kinetic energy of the particles in a sample
heat temperature2
Heat & Temperature
  • Both cups of water are at the same temperature…
    • but which has a bigger heat content?
heat temperature3
Heat & Temperature
  • Compare a match and a warm bath…
    • Which has higher temperature?
    • Which has a bigger heat content?
heat transfer big ideas
Heat Transfer – BIG IDEAS
  • Energy (heat) transfers occur during both physical and chemical changes
heat transfer big ideas1
Heat Transfer – BIG IDEAS
  • Heat will always be spontaneously transferred from matter at a higher temperature to matter at a lower temperature
    • When you place your hand on a cold glass, it feels cold – but what is really happening in terms of heat transfer?
      • Heat is being transferred from your hand to the glass.
heat transfer
Heat Transfer
  • Endothermic
    • If a system absorbs heat, the system undergoes an endothermic change
    • Heat is being absorbed from the surroundings, which makes it feel cold!
heat transfer1
Heat Transfer
  • Exothermic
    • If a system releases/loses heat, the system undergoes an exothermic change
    • Heat is being released into the surroundings, which makes it feel warm!
enthalpy
Enthalpy
  • Thermochemistry is the study of heat exchange during physical and chemical changes
  • Enthalpy (H) is a measure of the heat contained by a system
    • Remember, though, that we cannot measure heat directly, so instead we measure changes in enthalpy (ΔH)
    • An enthalpy change is the amount of heat absorbed/lost by a system as heat
enthalpy changes
Enthalpy Changes
  • If a system is gaining energy (+ΔH) in an endothermic change, why does the system feel cold to the touch?
    • Because the system is absorbing heat from the environment (i.e., your hand)
    • Similarly, when a system undergoes an exothermic change, you feel the heat being released by the system as warmth
enthalpy of reaction1
Enthalpy of Reaction
  • The quantity of energy transferred during a chemical reaction is referred to as the enthalpy of reaction (or “heat of reaction”)
  • The reaction system includes all reactants and products
    • The enthalpy of reaction is calculated by:
enthalpy of reaction2
Enthalpy of Reaction
  • If the amount of energy stored in the bonds of the reactants is greater than the amount of energy stored in the bonds of the products, then the system has lost energy
    • - ΔH
    • Exothermic
enthalpy of reaction3
Enthalpy of Reaction
  • If the amount of energy stored in the bonds of the reactants is less than the amount of energy stored in the bonds of the products, then the system has gained/absorbed energy
    • + ΔH
    • Endothermic
calculating enthalpy of reaction
Calculating Enthalpy of Reaction
  • The enthalpy of reaction can be calculated using the individuals heats of formation of each product/reactant in a chemical reaction
  • Example:
    • Calculate the standard reaction enthalpy, DHoRXN for the following equation:
    • Is this reaction exothermic or endothermic?
      • = -890.2 kJ
      • Exothermic
thermochemical equations
Thermochemical Equations
  • A thermochemical equation includes the ΔH of the reaction

EXOTHERMIC

ENDOTHERMIC

thermochemical equations1
Thermochemical Equations

EXOTHERMIC

ENDOTHERMIC

thermochemical stoichiometry
ThermochemicalStoichiometry
  • Sample Problem:

How much heat will be released if 1.00 g of hydrogen peroxide decomposes in a bombardier beetle to produce a steam spray?

chemistry question
Chemistry Question
  • Calculate the standard enthalpy of reaction for the combustion of ethyne (C2H6)
    • Answer = -3121.2 kJ
    • Compare this answer to the heat of combustion of ethane in Table A-5. What do you notice?
    • Why are these values different?
terms
Terms
  • Enthalpy of reaction
    • Quantity of energy transferred during a chemical reaction
  • Molar enthalpy of formation
    • Enthalpy change that occurs when one mole of a compound is formed
    • Defined as one mole of product
  • Molar enthalpy of combustion
    • Enthalpy change that occurs during the complete combustion of one mole of a substance
    • Defined as one mole of reactant
standard states
Standard States
  • The degree sign is used to indicate a measurement made at the standard state of a substance
    • Typically 1 atm and room temperature (~298 K)
    • ΔH°f for water is the standard enthalpy of formation for liquid water (since water would be liquid under standard conditions)
bonus assignment for today
Bonus Assignment for Today
  • A) ΔH°RXN for combustion of ethyne (C2H2)
  • B) The absolute value of the amount of energy released when 1.000 pound of ethyne is completely combusted (1 lb = 453.6 g)
  • C) The mass of ethyne combusted if ΔH = -1200. kJ
  • D) The absolute value of the amount of energy released when 14 g of ethyne are mixed with 25 g of O2
heat transfer2
Heat Transfer
  • Heat is always transferred from a “warm” object to a “cold” object – until an equilibrium is reached
    • Say you place some ice at into a beaker of hot water at. What factors will affect the final equilibrium temperature of the system?
energy3
Energy
  • calorie
    • the amount of energy required to raise the temperature of one gram of water by one degree Celsius.
  • Joule
    • SI unit of energy
    • 1 cal = 4.184 J
energy4
Energy
  • Express 60.1 cal of energy in units of joules.
calculating energy requirements
Calculating Energy Requirements
  • How does the amount of substance heated affect the energy required?
    • Determine the amount of energy in joules required to raise the temperature of 7.40 g of water from 29 °C to 46 °C.
specific heat
Specific Heat
  • The amount of heat required to raise the temperature of 1 g of any substance by 1oC
    • For liquid water, s = 4.184
specific heat1
Specific Heat
  • Intensive Property!
  • Specific heat describes how well an object retains heat
    • A substance with a low specific heat is quickly heated, but also quickly cools
    • A substance with a high specific heat takes a long time to warm up, but will also retain that heat for a longer period
specific heat2
Specific Heat
  • Styrofoam is a very poor conductor of heat; it is a good insulator. It has a high specific heat.
  • Metals are good conductors of heat. They have low specific heats.
heat transfer3
Heat Transfer
  • Say you place some ice at into a beaker of hot water at. What factors will affect the final equilibrium temperature of the system?
    • Mass of ice and water
    • Initial temperature of the water/ice
    • The specific heat of the ice/water
specific heat3
Specific Heat
  • What quantity of energy (in joules) is required to heat a piece of iron weighing 1.3 g from 25 °C to 46 ° C?
  • The specific heat of iron is 0.45
specific heat4
Specific Heat
  • Q = s x m xΔT
    • Q = energy required
    • s = specific heat capacity
    • m = mass of sample in grams
    • ΔT = change in temperature in degrees Celsius
specific heat5
Specific Heat
  • A 1.6 g sample of a metal that has the appearance of gold requires 5.8 J of energy to change its temperature from 23 °C to 41 °C. Is the metal pure gold?
  • Specific heat capacity of gold = 0.13
heat transfer4
Heat Transfer
  • The Law of Conservation of Energy says that energy can never be created or destroyed
    • So, if one object is absorbing heat, another object must be losing it, and vice versa
    • We can use this idea to solve problems involving two objects.
example
Example
  • A 5.0 g sample of an unknown metal is placed into a beaker of water containing 500. g of water. The metal has an initial temperature of 45.0°C and the water has an initial temperature of 20.0°C. When equilibrium is reached, the final temperature is 33.5°C. What is the specific heat of the metal?
example1
Example
  • A 20.0 g piece of aluminum (cp = 0.902 J/g·°C) is placed into a beaker containing 50.0 mL of water at 10.0°C. After a few minutes, an equilibrium temperature of 25.0°C is reached. What was the initial temperature of the aluminum metal?
heat transfer5
Heat Transfer
  • The equation q = mcΔt can be used to calculate heat transfer when temperature is changing.
    • BUT remember that when a substance is undergoing a phase change, its temperature remains constant
      • The energy being added is instead being used to overcome intermolecular forces
heat transfer6
Heat Transfer
  • To find the total amount of energy transferred during a phase change, we must also consider the heat of fusion and the heat of vaporization of the substance.
phase changes
Phase Changes
  • During solidification or melting

DH = Kfm

Kf = heat of fusion, the amount of heat needed to melt/freeze 1g of a substance

phase changes1
Phase Changes
  • During boiling or condensing
    • DH = Kvm
    • Kv = heat of vaporization, the amount of heat needed to boil/condense 1g of a substance
example2
Example
  • How much heat must be added to raise the temperature of a 10.0 g sample of ice at 0.0°C to 15.0°C?
example3
Example
  • How much heat must be added to change a 10.0 g sample of ice at 0.0°C to steam at 150.0°C? The specific heat of water vapor is 2.03 J/g·°C
hess law1
Hess’ Law
  • The overall enthalpy change in a reaction is equal to the sum of enthalpy changes for the individual steps in the process
  • We can use this law to calculate the total enthalpy change for a reaction by using the measured values of each individual step
example4
Example
  • A 50.0 g iron sample is heated to a temperature of 123.0 oC, then dropped into 40.0 mL of water. The initial temperature of the water is 98.0 oC. Assuming no heat is lost to the surroundings, how many grams of water boil away? The specific heat of iron is 0.451 J/goC and the Kv of water is 2260 J/g.
example5
Example
  • A 10.0 g sample of liquid water is heated from 15.0°C to 45.0°C. How much energy was added?