(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.

Energy and Chemical Change

Energy- the ability to do work or produce heat 2 Forms: Potential Energy Kinetic Energy

Potential Energy Potential Energy -energy due to the composition or position of an object. Ex: water stored behind a dam depends on composition: 1. the type of atoms 2. the number and type of chemical bonds joining the atoms 3. the way the atoms are arranged.

Kinetic Energy Kinetic Energy – is the energy of motion Ex: water flows from the dam

Chemical systems contain both potential and kinetic energy Potential Kinetic

Heat- represented by symbol Q- energy that is in the process of flowing from a warmer object to a cooler object

Chemical Potential Energy - the energy stored in a substance because of its composition. Composition is the type, number, and arrangement of atoms and bonds.

Thermal energy the energy created by moving particles inside a substance. more movement of particles = more thermal energy

Heat is Thermal energy that is transferred Heat is Transferred in 3 ways Conduction – the way heat moves through solids. (direct transfer) Vibrating molecules pass on heat from molecule to molecule.

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

Radiation Radiation – 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.

Forms of Energy

Wednesday

Phase Changes http://www.youtube.com/watch?feature=player_embedded&v=YG77v1PwQNM

Specific Heat –is the amount of heat required to raise the temperature of one gram of that substance by one degree Celsius. each substance has its own specific heat Table 16-2 pg 492

Heat of Vaporization The amount of heat required to convert unit mass of a liquid into the vapor without a change in temperature.

Heat of Fusion The amount of heat required to convert unit mass of a solid into the liquid without a change in temperature.

Measuring HEAT!!!

Two units for measuring heat calorie - the amount of heat required to raise the temperature of one gram of pure water by one degree Celsius Joule - SI unit of heat and energy

1 calorie = 4.184 joules 1000 calorie = 1 Calorie 1J = 0.2390 calories Table 16-1 Conversion factors and relationships pg 491

Calories are nutritional or food Calories 1 Calorie = 1000 calories 1Calorie = 1 kilocalorie approximates the energy needed to increase the temperature of 1 kilogram of water by 1 °C. The small calorie or gram calorie (symbol: cal)[2] 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)[2] 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.

Calculating Specific Heat Q = m x c x ΔT Q = heat absorbed or released m = mass of the sample in grams c = specific heat of the substance ΔT = difference between final temperature and initial temperature, or Tfinal- Tinitial

16.2 Heat in Chemical Reactions and Processes Measuring Heat Heat changes are measured with a calorimeter

Lab and worksheet The 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 (50.4-25) 114 = 254c C = 114/254 = 0.449 J/goC

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.

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

Calculating Heat Example 125 g water with an Initial temperature of 25.60C 50 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 29.3 0C. Calculate the Heat gained by the water. Example Part A: q = c x m x /\T q water = 4.184 J/(g x0C) x 125 g x (29.30C – 25.60C) q water = 4.184 J/(g x0C) x 125 g X 3.7 0C q water = 1900 J

Calculating Specific Heat Example 50 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 29.3 0C. Water absorbed 1900 J of heat. Example Part B: Calculate the Specific Heat of the Metal c = q___ m x /\T c metal = 1900 J m x /\T c metal = _______1900 J_________ (50.0 g)(1150C – 29.3 0C) c metal = ____1900 J_____ (50.0 g)(85.700C) c metal = 0.44 J/(g x 0C) specific heat of the metal Look at pg 492 at the table. What is this metal?

Thursday- Lab

Friday- Practice worksheet

Monday

16.3 and 16.4 Enthalpy and Enthalpy Changes Enthalpy- (H) the heat content of a system at a constant pressure

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.

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 textbook ΔH rxn = H final – H initial or ΔH rxn = H products – H reactants Example: What is the heat of reaction for the following reaction? H2S + 4F2  2HF + SF6

Endothermic Reaction If 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 + NO3 NH4NO3  NH4 + NO3 ΔH = +27 kJ Energy required to break the bonds in a reactant is less than released after the bonds in the product is formed

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 Fe2O3 + 1625 kJ (energy) 4 Fe + 3O2  2 Fe2O3 ΔH = -1625 kJ Energy needed to break the bond in the reactant is more than energy released after the bonds in the products are formed http://www.youtube.com/watch?v=ksN-t2mmpvM&feature=related http://www.youtube.com/watch?v=ksN-t2mmpvM&feature=related

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Sign of the Enthalpy of Reaction Exothermic reactions have a negative enthalpy Hproducts < Hreactants Endothermic reactions have a positive enthalpy Hproducts > Hreactants

16.3 Thermochemical Equations Enthalpy (heat) of combustion- enthalpy change for the complete burning of one mole of the substance ΔHcomb

Entropy Measure of the disorder or randomness of the particles that make up a system Symbolized by S

Molar Enthalpy (heat) of Vaporization Heat required to vaporize one mole of a liquid ΔHvap Endothermic (positive enthalpy)

Molar Enthalpy (heat) of Fusion The heat required to melt one mole of a solid substance ΔHfus Endothermic (positive enthalpy)

16.5 Reaction Spontaneity Spontaneous process- physical or chemical change that occurs with no outside intervention

Law of Disorder States that spontaneous processes always proceed in such a way that the entropy of the universe increases

Chemical Energy and the Universe Thermochemistry – the study of heat changes that accompany chemical reactions and phase changes.

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

universe – the system plus the surroundings universe = system + surroundings

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 Fe2O3 + 1625 kJ (energy) 4 Fe + 3O2  2 Fe2O3 Heat of rxn = -1625 kJ

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 + NO3 NH4NO3  NH4 + NO3 Heat of rxn = 27 kJ