Presentation on theme: "(11) Science concepts. The student understands the energy changes that occur in chemical reactions. The student is expected to: (A) understand energy and."— Presentation transcript:
1 (11) Science concepts. The student understands the energy changes that occur in chemical reactions. The student is expected to:(A) understand energy and its forms, including kinetic, potential, chemical, and thermal energies;(B) understand the law of conservation of energy and the processes of heat transfer;(C) use thermochemical equations to calculate energy changes that occur in chemical reactions and classify reactions as exothermic or endothermic;(D) perform calculations involving heat, mass, temperature change, and specific heat; and(E) use calorimetry to calculate the heat of a chemical process.
3 Energy- the ability to do work or produce heat 2 Forms: Potential EnergyKinetic Energy
4 Potential EnergyPotential Energy -energy due to the composition or position of an object.Ex: water stored behind a damdepends on composition:1. the type of atoms2. the number and type of chemical bonds joining the atoms3. the way the atoms are arranged.
5 Kinetic EnergyKinetic Energy – is the energy of motionEx: water flows from the dam
6 Chemical systems contain both potential and kinetic energy Potential Kinetic
7 Heat- represented by symbol Q- energy that is in the process of flowing from a warmer object to a cooler object
8 Chemical Potential Energy - the energy stored in a substance because of its composition.Composition is the type, number, and arrangement of atoms and bonds.
9 Thermal energythe energy created by moving particles inside a substance.more movement of particles = more thermal energy
10 Heat is Thermal energy that is transferred Heat is Transferred in 3 waysConduction – the way heat moves through solids. (direct transfer)Vibrating molecules pass on heat from molecule to molecule.
11 Convection – the way heat moves through gases and liquids. Heated molecules move AWAY from the heat and cooler molecules take their place.Ex: Hot air rises and cool air sinks
12 RadiationRadiation – the way heat moves through empty space.Does not need atoms or molecules to work.Electromagnetic radiation – light and heat from the sun, visible light, microwaves, X-rays, etc.
24 Calories are nutritional or food Calories 1 Calorie = 1000 calories1Calorie = 1 kilocalorieapproximates the energy needed to increase the temperature of 1 kilogram of water by 1 °C.The small calorie or gram calorie (symbol: cal) approximates the energy needed to increase the temperature of 1 gram of water by 1 °C. This is about 4.2 joules.The large calorie, kilogram calorie, dietary calorie, or food calorie (symbol: Cal) approximates the energy needed to increase the temperature of 1 kilogram of water by 1 °C. This is exactly 1,000 small calories or about 4.2 kilojoules. It is also called the nutritionist's calorie.
25 Calculating Specific Heat Q = m x c x ΔTQ = heat absorbed or releasedm = mass of the sample in gramsc = specific heat of the substanceΔT = difference between final temperature and initial temperature, or Tfinal- Tinitial
26 16.2 Heat in Chemical Reactions and Processes Measuring HeatHeat changes are measured with a calorimeter
27 Lab and worksheetThe temperature of a sample of iron has a mass of 10.0g changed from 50.4oC to 25.0oC with the release of 114 J of heat. What is the specific heat of iron? Q = mc∆T 114 = 10 x c x ( ) 114 = 254c C = 114/254 = J/goC
28 Calorimeter – an insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process.Data is the change in temperature of this mass of the substance.
29 Determining Specific Heat Place a hot metal into water.Heat flows from the hot metal to the cooler water until the temperature of the metal and water are equal.The heat gained by the water is equal to the heat lost by the metal
30 Calculating Heat Example 125 g water with an Initial temperature of 25.60C50 g metal at 1150C is placed in the water.Heat flows from the hot metal to the cooler water until the temperature of the metal and water are equal. Both have a final temperature of C. Calculate the Heat gained by the water.Example Part A:q = c x m x /\Tq water = J/(g x0C) x 125 g x (29.30C – 25.60C)q water = J/(g x0C) x 125 g X 3.7 0Cq water = 1900 J
31 Calculating Specific Heat Example50 g metal at 1150C is placed in the water.Heat flows from the hot metal to the cooler water until the temperature of the metal and water are equal. Both have a final temperature of C.Water absorbed 1900 J of heat.Example Part B: Calculate the Specific Heat of the Metalc = q___m x /\Tc metal = 1900 Jm x /\Tc metal = _______1900 J_________(50.0 g)(1150C – C)c metal = ____1900 J_____(50.0 g)(85.700C)c metal = 0.44 J/(g x 0C) specific heat of the metalLook at pg 492 at the table. What is this metal?
35 16.3 and 16.4 Enthalpy and Enthalpy Changes Enthalpy- (H) the heat content of a system at a constant pressure
36 A thermochemical equation is a balanced chemical equation that includes the physical states of all reactants and products and the energy change expressed as the change in enthalpy, ∆H.
37 Use the table on pg. 510 in your textbook You can’t measure actual enthalpy, but you can measure change in enthalpy, which is called enthalpy (heat) of reaction (ΔH rxn)Use the table on pg. 510 in your textbookΔH rxn = H final – H initial orΔH rxn = H products – H reactantsExample:What is the heat of reaction for the following reaction? H2S + 4F2 2HF + SF6
38 Endothermic ReactionIf the ∆H is shown on the reactants side, it is endothermic (gaining energy)The heat of the reaction will be positive.(energy) 27 kJ + NH4NO3 NH4 + NO3NH4NO3 NH4 + NO3 ΔH = +27 kJEnergy required to break the bonds in a reactant is less than released after the bonds in the product is formed
39 Exothermic Reaction The heat of the reaction will be negative. If the ∆H is shown on the products side, it is exothermic (losing energy)The heat of the reaction will be negative.4 Fe + 3O2 2 Fe2O kJ (energy)4 Fe + 3O2 2 Fe2O3 ΔH = kJEnergy needed to break the bond in the reactant is more than energy released after the bonds in the products are formed
41 Sign of the Enthalpy of Reaction Exothermic reactions have a negative enthalpyHproducts < HreactantsEndothermic reactions have a positive enthalpyHproducts > Hreactants
42 16.3 Thermochemical Equations Enthalpy (heat) of combustion- enthalpy change for the complete burning of one mole of the substanceΔHcomb
43 EntropyMeasure of the disorder or randomness of the particles that make up a systemSymbolized by S
44 Molar Enthalpy (heat) of Vaporization Heat required to vaporize one mole of a liquidΔHvapEndothermic (positive enthalpy)
45 Molar Enthalpy (heat) of Fusion The heat required to melt one mole of a solid substanceΔHfusEndothermic (positive enthalpy)
46 16.5 Reaction SpontaneitySpontaneous process- physical or chemical change that occurs with no outside intervention
47 Law of DisorderStates that spontaneous processes always proceed in such a way that the entropy of the universe increases
48 Chemical Energy and the Universe Thermochemistry – the study of heat changes that accompany chemical reactions and phase changes.
49 system – the specific part of the universe that contains the reaction or process you wish to study. surroundings – everything in the universe other than the system
50 universe – the system plus the surroundings universe = system + surroundings
51 Example: Using a heat pack to warm your hands Heat flows from the heat pack (the system) to your cold hands (surroundings)Exothermic - If energy is shown as a product it means that heat is released. The heat of the reaction will be negative.4 Fe + 3O2 2 Fe2O kJ (energy)4 Fe + 3O2 2 Fe2O3 Heat of rxn = kJ
52 Example: Using a cold pack on an injured knee Heat flows from the knee (the surroundings) to the cold pack (the system)Endothermic – If energy is shown as a reactant it means that energy is absorbed.The heat of the reaction will be positive.(energy) 27 kJ + NH4NO3 NH4 + NO3NH4NO3 NH4 + NO3 Heat of rxn = 27 kJ