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Section 10.1 Energy, Temperature, and Heat • Thermodynamics – study of energy • First law of thermodynamics – Energy of the universe is constant – Energy can not be created or destroyed Section 10.1 Energy, Temperature, and Heat • Energy is the ability to do work or produce heat. • 2 types of Energy Potential energy Energy of position Kinetic energy Energy of motion E = ½ mv2 Section 10.1 Energy, Temperature, and Heat • Law of conservation of energy – Energy can not be created or destroyed but can be converted from one form to another – Ex chemical energy (gas) can be converted to mechanical and thermal energy Section 10.1 Energy, Temperature, and Heat • Internal energy = E “sum” of kinetic & potential energies of all “particles” in a system Internal energy can be transferred by two types of energy flow: • Heat (q) • Work (w) (q = quantity of heat) E = q + w Change of energy in a system Section 10.1 Energy, Temperature, and Heat Measuring Energy Changes • common energy units for heat are the calorie and the joule. – Calorie – the amount of energy (heat) required to raise the temperature of one gram of water 1oC. – Joule – heat required to raise the temperature of 1 g of water by 0.24 K 1 calorie = 4.184 joules Section 10.1 Energy, Temperature, and Heat • System – part of the universe on which we focus attention • Surroundings – everything else in the universe • Ex. Burning a match is the system releases heat to surroundings Section 10.1 Energy, Temperature, and Heat Exothermic – energy flows out of system Endothermic – energy flows into system Section 10.1 Energy, Temperature, and Heat • C- Specific heat capacity energy required to change the mass of one gram of a substance by one degree Celsius. [joule/gram °C] • Like energy storage bank Section 10.1 Energy, Temperature, and Heat • C- Specific heat capacity energy required to change the mass of one gram of a substance by one degree Celsius. [joule/gram °C] • Like energy storage bank • Water has a high specific heat capacity that means it takes a lot of heat (joules) to increase the temperature of water compared to metals (water better at storing energy than metals Section 10.1 Energy, Temperature, and Heat B. Measuring Energy Changes • To calculate the energy required for a reaction: Q = C m t Section 10.1 Energy, Temperature, and Heat Q = C m t How much heat is required to warm 145 g of water from 25 *C to 65 *C Section 10.1 Energy, Temperature, and Heat Q = m x C t How much heat is required to warm 145 g of water from 25 *C to 65 *C the specific heat of liquid water is 4.184 J/g *C Q = m x C t Q= 145 g x 4.184 J x (65 *C - 25 *C) g *C Section 10.1 Energy, Temperature, and Heat Q = m x C t How much heat is required to warm 145 g of water from 25 *C to 65 *C the specific heat of liquid water is 4.184 J/g *C Q = C m t Q= 145 g x 4.184 J x (65 *C - 25 *C) g *C Q = 145 x 4.184 J x 40 Q = 24,267 joules Section 10.1 Energy, Temperature, and Heat A quantity of water is heated from 20.0 *C to 48.3 *C by absorbing 623 Joules of heat energy, what is the mass of the water? Specific heat of water is 4.184 J/g *C Q = C m t __Q__ = m C x t Section 10.1 Energy, Temperature, and Heat A quantity of water is heated from 20.0 *C to 48.3 *C by absorbing 623 Joules of heat energy, what is the mass of the water? Specific heat of water is 4.184 J/g *C Q = C m t __Q__ = m C x t 623 J = 4.184 J /g *C x (48.3 *C - 20 *C) m Section 10.1 Energy, Temperature, and Heat A quantity of water is heated from 20.0 *C to 48.3 *C by absorbing 623 Joules of heat energy, what is the mass of the water? Specific heat of water is 4.184 J/g *C Q = C m t __Q__ = m C x t 623 J = m 4.184 J /g *C x (48.3 *C - 20 *C) 623 J 4.184 J /g *C x (28.3 *C) 5.26 g = m = m Section 10.1 Energy, Temperature, and Heat In a household radiator a 25.5 g of steam at 100 *C condenses (changes from gas to liquid) How much heat is released? Heat of vaporization is 2260 J/g Q = m Hvap Section 10.1 Energy, Temperature, and Heat In a household radiator a 25.5 g of steam at 100 *C condenses (changes from gas to liquid) How much heat is released? Heat of vaporization is 2260 J/g Q = m Hvap Q = 25.5 g x 2260 J/g Q = 57,630 J Section 10.1 Energy, Temperature, and Heat If an endothermic reaction absorbs 482 J how many calories are absorbed? 482 J x 1 calorie = 115 calories 4.184 J Section 10.1 Energy, Temperature, and Heat Q = m x C t How much heat is required to warm 275 g of water from 76 *C to 87 *C the specific heat of liquid water is 4.184 J/g *C Q = C m t Q= 275 g x 4.184 J x (87 *C - 76 *C) g *C Q = 275 x 4.184 J x 11 Q = 12,657 joules Section 10.1 Energy, Temperature, and Heat In a household radiator a 1000 g of steam at 100 *C condenses (changes from gas to liquid) How much heat is released? Heat of vaporization is 2260 J/g Q = m Hvap Q = 1,000 g x 2260 J/g Q = 2,260,000 J Section 10.1 Energy, Temperature, and Heat Objectives 1. To consider the heat (enthalpy) of chemical reactions 2. To understand Hess’s Law Section 10.1 Energy, Temperature, and Heat A. Thermochemistry (Enthalpy) • Enthalpy, H – amount of heat content of a system absorbed or released at constant pressure. • H = heat = q ΔH = q (valid for constant pressure ONLY!) That is the Enthalpy ΔH (heat content) = q (quantity of heat) The heat of formation is the enthalpy gain when a compound is CREATED from its elements at their conditions at some STP. The heat of reaction involves CHEMICAL REACTIONS OF A SUBSTANCE WITH OTHER SUBSTANCES, and is the relative enthalpy of the products compared to the reactants at some STP. Section 10.1 Energy, Temperature, and Heat • Enthalpy, H – • Since H = q • And q = m x C x ΔT then… ΔH = mCΔT ΔH = mΔHvap ΔH = mΔHfus Section 10.1 Energy, Temperature, and Heat A. Thermochemistry Calorimeter • Enthalpy, H is measured using a calorimeter. Section 10.1 Energy, Temperature, and Heat Temperature- average kinetic energy Heat- flow of energy due to a temperature difference Enthalpy- heat content of a system Entropy- measure of disorder or randomness in a system Section 10.1 Energy, Temperature, and Heat Hess’s Law • the change in enthalpy is the same whether the reaction takes place in one step or a series of steps. • Example: N2(g) + 2O2(g) 2NO2(g) H1 = 68 kJ Since H = + 68 kJ then this is a endothermic reaction Section 10.1 Energy, Temperature, and Heat ΔHrxn = Heat of Reaction (represents change in enthalpy) ΔHrxn = (ΔHproducts -ΔHreactants) –ΔH = Exothermic +ΔH = Endothermic How much heat is required to raise the temperature of 445 g of zinc from 25.0*C to 74.5*C if the specific heat of zinc is 0.390 J/g*C How much heat is required to raise the temperature of 445 g of zinc from 25.0*C to 74.5*C if the specific heat of zinc is 0.390 J/g*C ΔH = mCΔT How much heat is required to raise the temperature of 445 g of zinc from 25.0*C to 74.5*C if the specific heat of zinc is 0.390 J/g*C ΔH = mCΔT ΔH = 445g x 0.390 J/g*C x 49.5*C ΔH = 8,591 Joules Determine the heat of reaction of the following equation CaCO3 CO2 + CaO Heat of formation for calcium carbonate is -1207 KJ/mol for 1 mole Heat of formation for carbon dioxide is -394 KJ/mol for 1 mole Heat of formation for calcium oxide is -635 KJ/mol for 1 mole Hreaction = SUM (Hproducts) - SUM (Hreactants) Determine the heat of reaction of the following equation CaCO3 CO2 + CaO Heat of formation for calcium carbonate is -1207 for 1 mole Heat of formation for carbon dioxide is -394 for 1 mole Heat of formation for calcium oxide is -635 for 1 mole Hreaction = SUM (Hproducts) - SUM (Hreactants) Hreaction = (-394 – 635) – (-1207) Determine the heat of reaction of the following equation CaCO3 CO2 + CaO Heat of formation for calcium carbonate is -1207 for 1 mole Heat of formation for carbon dioxide is -394 for 1 mole Heat of formation for calcium oxide is -635 for 1 mole Hreaction = SUM (Hproducts) - SUM (Hreactants) Hreaction = (-394 – 635) – (-1207) Hrxn = (-1029) – (-1207) Hrxn = +178 Determine the heat of reaction of the following equation CaCO3 CO2 + CaO Heat of formation for calcium carbonate is -1207 for 1 mole Heat of formation for carbon dioxide is -394 for 1 mole Heat of formation for calcium oxide is -635 for 1 mole Hreaction = SUM (Hproducts) - SUM (Hreactants) Hreaction = (-394 – 635) – (-1207) Hrxn = (-1029) – (-1207) Hrxn = +178 ENDOTHERMIC The Heat of Formation for ALL Diatomic molecules is zero "All elements in their standard states (oxygen gas, solid carbon in the form of graphite, etc.) have a standard enthalpy of formation of zero, as there is no change involved in their formation." The problem is that we can only measure CHANGES in the Enthalpy of the system, and have no way to determine the ABSOLUTE Enthalpy. We have to define a convenient ZERO of Enthalpy for all the chemical compounds that we can make. We do this by defining the enthalpy of the Standard State of the Elements, this should cover all possible matter since all matter is made of elements we only need about 100 descriptions of Standard Elements Section 10.1 Energy, Temperature, and Heat Objectives 1. To understand how the quality of energy changes as it is used 2. To consider the energy resources of our world 3. To understand energy as a driving force for natural processes Section 10.1 Energy, Temperature, and Heat A. Quality Versus Quantity of Energy • When we use energy to do work we degrade its usefulness. Section 10.1 Energy, Temperature, and Heat A. Quality Versus Quantity of Energy • Petroleum as energy Section 10.1 Energy, Temperature, and Heat B. Energy and Our World • Fossil fuel – carbon based molecules from decomposing plants and animals – Energy source for United States Section 10.1 Energy, Temperature, and Heat B. Energy and Our World • Petroleum – thick liquids composed of mainly hydrocarbons – Hydrocarbon – compound composed of C and H Section 10.1 Energy, Temperature, and Heat B. Energy and Our World • Natural gas – gas composed of hydrocarbons Section 10.1 Energy, Temperature, and Heat B. Energy and Our World • Coal – formed from the remains of plants under high pressure and heat over time Section 10.1 Energy, Temperature, and Heat B. Energy and Our World • Effects of carbon dioxide on climate • Greenhouse effect Section 10.1 Energy, Temperature, and Heat B. Energy and Our World • Effects of carbon dioxide on climate • Atmospheric CO2 – Controlled by water cycle – Could increase temperature by 10oC Section 10.1 Energy, Temperature, and Heat B. Energy and Our World • New energy sources – Solar – Nuclear – Biomass – Wind – Synthetic fuels Section 10.1 Energy, Temperature, and Heat C. Energy as a Driving Force • Natural processes occur in the direction that leads to an increase in the disorder of the universe. • Example: – Consider a gas trapped as shown Section 10.1 Energy, Temperature, and Heat C. Energy as a Driving Force • What happens when the valve is opened? Section 10.1 Energy, Temperature, and Heat C. Energy as a Driving Force • Two driving forces – Energy spread – Matter spread Section 10.1 Energy, Temperature, and Heat C. Energy as a Driving Force • Energy spread – In a given process concentrated energy is dispersed widely. – This happens in every exothermic process. Section 10.1 Energy, Temperature, and Heat C. Energy as a Driving Force • Matter spread – Molecules of a substance spread out to occupy a larger volume. – Processes are favored if they involve energy and matter spread. Section 10.1 Energy, Temperature, and Heat C. Energy as a Driving Force • Entropy, S – function which keeps track of the tendency for the components of the universe to become disordered Section 10.1 Energy, Temperature, and Heat C. Energy as a Driving Force • What happens to the disorder in the universe as energy and matter spread? Section 10.1 Energy, Temperature, and Heat C. Energy as a Driving Force • Second law of thermodynamics – The entropy of the universe is always increasing.