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Chapter 6

Chapter 6. Thermochemistry -study of heat changes that occur during chemical reactions, study of relationships between chemistry and energy energy - ability to do work or supply heat -unlike matter, energy is weightless, odorless, tasteless -only way to detect energy is by its effects

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Chapter 6

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  1. Chapter 6 Thermochemistry -study of heat changes that occur during chemical reactions, study of relationships between chemistry and energy energy- ability to do work or supply heat -unlike matter, energy is weightless, odorless, tasteless -only way to detect energy is by its effects work- result of a force acting through a distance

  2. heat- flow of energy from one object to another because of a temperature difference *energy is something that an object has *heat and work are ways in which objects exchange energy kinetic energy(KE)- energy of motion -faster = more kinetic energy potential energy(PE)- energy of position -a rock at the top of a hill has more PE than a rock at the bottom

  3. thermal energy- energy associated with temperature of an object -higher temp = more thermal energy -type of KE because it arises from motion of particles Law of Conservation of Energy -energy can neither be created nor destroyed -energy can be transferred from one object to another and from one form to another

  4. All chemical reactions involve the release or absorption of heat: system- what is being studied surroundings- everything else around the system in nature universe- system and surroundings together *the system either loses energy to or gains energy from the surroundings

  5. Heat Flow two directions • surroundings  system endothermic- system gains heat as surroundings cool, system absorbs heat • system  surroundings exothermic- system loses heat as surroundings heat up, system releases heat Units of Energy joule(J) - SI unit for heat/energy -can use kJ 1kJ = 1000J calorie(cal) - amount of energy required to raise the temp of 1g of pure water 1°C

  6. Calorie(Cal) - nutritional unit **1Cal = 1kcal = 1000cal **1J = 0.2390 cal 4.184J = 1cal Convert: • 1656.70J cal • 483.12cal  J • 0.56721Cal  J

  7. thermodynamics- general study of energy and its interconversions The First Law of Thermodynamics -states that the total energy of the universe is constant internal energy(E)- sum of the potential and kinetic energies of all the particles that compose the system **E = PE + KE -internal energy is a state function -value depends only on the state of the system (initial and final values) and not on how the system arrived at that state

  8. -Example on page 241 ∆E = Efinal - Einitial ∆ = change in -in a chemical reaction: ∆E = Eproducts - Ereactants *If the reactants have a higher internal energy than the products, ∆Esys is negative and the energy flows out of the system into the surroundings *If the reactants have a lower internal energy than the products, ∆Esys is positive and energy flows into the system from the surroundings

  9. -According to the 1st Law of Thermodynamics: ∆E = the amount of heat transferred and the work done ∆E = q + w q= heat w= work

  10. Ex- Identify each of the following energy exchanges as heat or work and determine whether the sign of heat or work (relative to the system) is positive or negative. • An ice cube melts and cools the surrounding beverage. (ice cube is the system) *heat, sign is positive 2) A metal cylinder is rolled up a ramp. (the cylinder is the system; assume no friction) *work, sign is positive • Steam condenses on skin, causing a burn. (the steam is the system) *heat, sign is negative

  11. Ex- A cylinder and piston assembly (defined as the system) is warmed by an external flame. The contents of the cylinder expand, doing work on the surroundings pushing the piston outward. If the system absorbs 732J of heat and 628J of work, what is the change in external energy? ∆E = q + w q= +732J w= -628J ∆E = +732J + (-628J) = +104.00J Try page 246- Ex 6.1 and For Practice 6.1

  12. *Notice the distinction between temperature and heat -temperature is a measure of thermal energy -heat is the transfer of thermal energy thermal equilibrium- when there is no additional heat transfer because system and surroundings are same temperature heat capacity (C) - the quantity of heat required to change the temperature of a system 1°C specific heat capacity (Cs) - the amount of heat needed to raise the temperature of 1g of a substance by 1°C

  13. -the higher the specific heat capacity, the more energy is required to raise the temperature -metals tend to have fairly low specific heats because they heat up easily -substances like water have very high specific heats because they take a long time to heat up q = (m)(Cs)(ΔT) q = heat (J) m = mass (g) Cs = specific heat capacity J/g°C *will be given to you if not solving for- page 247 ΔT = temp change Tfinal – Tinitial °C

  14. calorimetry- the accurate and precise measurement of heat change for chemical and physical processes In calorimetry: heat released = heat absorbed -to measure changes accurately and precisely, processes must be carried out in an insulated container (calorimeter)ex- foam cup bomb calorimeter- equipment designed to measure ∆E for combustion reactions, volume is constant page 251 figure 6.6

  15. -when a reaction is performed inside a bomb calorimeter, the following equation is used: qcal= (Ccal)(∆T) qcal= heat absorbed by calorimeter Ccal= heat capacity of entire calorimeter, usually given **the amount of heat gained by the calorimeter must equal the amount of heat released by the reaction -equal in magnitude, but opposite in sign qcal= -qrxn

  16. -since the reaction inside a calorimeter is occurring under conditions of constant volume, then: qrxn= ∆Erxn page 252

  17. -most chemical reactions and physical changes occur at constant pressure in a lab enthalpy (H)- heat content of a system at constant pressure H = ∆H = q -enthalpy change for chemical reaction = ΔHrxn -called heat of reaction ∆H is + for endothermic reaction -feels cold to the touch ∆H is – for exothermic reaction -feels warm to the touch

  18. Examples: Identify each as endo or exothermic and tell the sign on ∆H. • sweat evaporating from skin endothermic, +∆H • water freezing in a freezer exothermic, - ∆H • wood burning in a fire exothermic, - ∆H

  19. Try these!! • an ice cube melting endothermic, +∆H • nail polish remover quickly evaporating after it is accidently spilled on the skin endothermic, +∆H • a chemical hand warmer emitting heat after the mixing of substances within a small handheld package exothermic, -∆H

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