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What is Energy? • Energy is the ability to do work or produce heat. • Potential energy is energy due to composition or position. • Kinetic energy is energy of motion. Law of Conservation of Energy •The total energy of the universe is constant and can neither be created nor destroyed; it can only be transformed. •Energy can’t be destroyed, but it can: • Move or transfer to other locations • Transform into a different type Heat (represented by q) • Form of energy • Flows from a warmer object to a cooler object • Heat is transferred until both objects are at the same temperature Measuring Heat • A calorie is defined as the amount of energy required to raise the temperature of one gram of water one degree Celsius. • The energy content of food is measured in Calories, or 1000 calories. Also known as a kilocalorie. • A joule is the SI unit of heat and energy, equivalent to 0.2390 calories. Relationships among Energy Units 1 J = 0.2390 cal 1 cal = 4.184 J 1 Cal = 1 kcal 1 Cal = 4.184 kJ Heat Conversions • A breakfast of cereal, orange juice, and milk might contain 230 Calories. Convert this energy to joules. • A fruit and oatmeal bar contains 142 nutritional Calories. Convert this energy to calories. • An exothermic reaction releases 86.5 kJ of energy. How many calories of energy are released? Specific Heat • The specific heat of any substance is the amount of heat required to raise one gram of that substance one degree Celsius. Specific Heats at 298 K (25⁰C) Substance c in J/g°C Aluminum 0.897 Bismuth 0.123 Copper 0.386 Brass 0.380 Gold 0.126 Lead 0.128 Silver 0.233 Tungsten 0.134 Zinc 0.387 Mercury 0.140 Alcohol(ethyl) 2.44 Water 4.186 Ice (-10 C) 2.05 Granite 0.803 Concrete 0.840 Q=mc∆T • Tells us the heat that is gained or lost by a system • Q = heat gained or lost • Measured in Joules • M = mass of sample • C = specific heat of sample • ∆T = change in temperature (final temp-initial temp) • Must be measured in Celsius! Example 1 A 500.g piece of Iron increases its temperature 7oC when heat energy is added. How much heat energy produced this change in temperature? Example 2 When 300. calories of energy is lost from a 125g object, the temperature decreases from 45 oC to 40oC. What is the specific heat of this object? What was the most likely identity of the object? Example 3 A scientist wants to raise the temperature of a 10.0 g sample of aluminum from -45oC to 15oC. How much heat energy is required to produce this change in temperature? System and Surroundings • study of heat changes that accompany chemical reactions and phase changes. • Between the system and the surroundings • The system is the specific part of the universe that contains the reaction or process you wish to study. • The surroundings are everything else other than the system in the universe Heat Transfer • Heat is neither created nor destroyed; it is transferred • • Endothermic – system absorbs heat from the surroundings Exothermic – system releases heat to the surroundings Endo (absorb heat?) or Exo (release heat?) Endothermic or exothermic process? evaporating alcohol leaves burning boiling water water cooling melting ice freezing water Endo vs. Exo • Endothermic • Heat is a reactant • Q is positve • Exothermic • Heat is a product • Q is negative Enthalpy • Total heat content of a system • Heat of fusion, is the amount of energy required to melt one gram of a substance • Heat of vaporization, is the amount of energy required to boil one gram of a substance Calorimeter • insulated device used for measuring the amount of heat absorbed or released in a chemical reaction or physical process. Phase Change Diagrams Heat of Fusion Heat of Vaporization Increasing Temperature Condensation Evaporation Freezing Melting Increasing Energy Heat of fusionH2O = 334 J/g Heat of vaporizationH2O = 2260 J/g Ex. How much heat, in joules, is needed to heat 150.0 g liquid water from 0.0°C to 20.0OC? • Happens in one phase • Q=mcT • Q=(150.0g)(4.184 J/g*°C)(20°C) Ex. How much heat, in joules, is needed to melt 150.0 g of ice at 0oC? • Happening at one temp • Phase change • Q=mHf • Q=(150.0g)(334 J/g) Ex. How many joules are needed to convert 5.0 g of ice at -15oC to steam at 130oC? • Changing temps • Changing phases • Q=mcT • Q=(5.0 g)(2.09 J/g°C)(15°C) – why 15°C??? • Q=(5.0 g)(334 J/g) • Q=(5.0 g)(4.184 J/g°C)(100°C) – why 100°C??? • Q=(5.0 g)(2260 J/g) • Q=(5.0 g)(2.01 J/g°C)(35°C) Add these together to get total Q Exothermic vs. Endothermic Recap • Which type of reaction needs heat/energy? • Endothermic • Which type of reaction releases heat/energy? • Exothermic • Which reaction would you see the reactants have a lower total energy than the products? • Endothermic • Which reaction would you see the reactants have a higher total energy than the products? • Exothermic Energy Diagrams • Activation Energy - the minimum quantity of energy that the reacting species must possess in order to undergo a specified reaction • Activated complex - the maximum energy point along the reaction path • H – difference in energy between the products and the reactants (H = Hproducts – Hreactants) Endothermic and Exothermic Endothermic Exothermic Energy Diagram H H is positive H is negative CH4(g) + 2O2(g) CO2(g) + Sample 2H2O + 241.8 kJ 2H2(g) + O2(g) 2H2O(g) + 890 kJ Reaction ΔH = 241.8 kJ (energy is required) ΔH = -890 kJ (energy is released) Where’s the energy? Energy is a reactant Energy is a product Using Heat in Stoichiometry • Can be used in reaction as reactant or product (kJ/ mole) CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + 890 kJ • In this equation, 890 kJ of heat are produced per the amount of each reactant listed. How much energy is produced if 5.0 L of oxygen were consumed? 1 𝑚𝑜𝑙𝑒 𝑂2 890 𝑘𝐽 5.0 L O2 x x = 99 kJ 22.4 𝐿 𝑂2 2 𝑚𝑜𝑙𝑒 𝑂2 Example 2 1049 kJ of energy are released when aluminum reacts with hydrochloric acid 2 Al (s) + 6 HCl (aq) 2 AlCl3 (aq) + 3 H2 (g) + 1049 kJ How many grams of aluminum are needed when 616 kJ of energy are produced? 2 𝑚𝑜𝑙𝑒𝑠 𝐴𝑙 26.98 𝑔 𝐴𝑙 616 kJ x x = 31.7 g Al 1049 𝑘𝐽 1 𝑚𝑜𝑙𝑒 𝐴𝑙 Example 3 In the combustion of C2H6, 3120 kJ of energy are produced. How much energy is produced in the combustion of 21.50 grams of C 2H6? 2 C2H6 (g) + 7 O2 (g) 4 CO2 (g) + 6 H2O (l) + 3120 kJ 1 𝑚𝑜𝑙𝑒 𝐶2𝐻6 3120 𝑘𝐽 21.50 g C2H6 x x = 1115 kJ 30.08 𝑔 𝐶2𝐻6 2 𝑚𝑜𝑙𝑒 𝐶2𝐻6 Calculating H H = Hproducts – Hreactants Look up the enthalpies of each reactant and product Sum of the products – sum of reactants Ex. What is the H in the combustion of methane? CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g) H = (-394 kJ + (2 x -286 kJ)) – (-75 kJ + (2 x 0 kJ)) = -891 kJ Is the reaction endothermic or exothermic? Example 2 How much energy is involved in the reaction between magnesium and oxygen? 2 Mg (s) + O2 (g) 2 MgO (s) H = (2 x -601.24 kJ) – ((2 x 0 kJ) + 0 kJ) = -1202.48 kJ Example 3 How much energy is involved in the decomposition of solid magnesium hydroxide? Mg(OH)2 (s) H2O + MgO (s) H = (-286 kJ + -601.24 kJ) - (-924.54 kJ) = 37.1 kJ Hess’s Law Hess’s law states that if you can add two or more thermochemical equations to produce a final equation for a reaction, then the sum of the enthalpy changes for the individual reactions is the enthalpy change for the final reaction. Example 1 How much energy is involved in the decomposition of hydrogen peroxide? 2 H 2O 2 2 H 2O + O 2 a b 2 H 2 + O 2 2 H 2O H2 + O 2 H2O 2 ΔH = -572 kJ ΔH = -188 kJ Since H2O2 is a reactant, we need to flip reaction b and change the sign. What else do we have to do with equation b? Example 1 2 H 2O 2 2 H 2O + O 2 a b 2 H 2 + O 2 2 H 2O H2 + O 2 H2O 2 ΔH = -572 kJ ΔH = -188 kJ B becomes 2 H2O2 2 H2 + 2 O2 and ΔH = 188 kJ (x2) = 376 kJ Add b to a to get total ΔH -572 kJ + 376 kJ = -196 kJ Example 2 Calculate the enthalpy of the following chemical reaction: CS2 (l) + 3 O2(g) CO2 (g) + 2 SO2 (g) a b c C (s) + O2 CO2 (g) S (s) + O2 SO2 (g) C (s) + 2 S (s) CS2 (l) ΔH = -393.5 kJ ΔH = -296.8 kJ ΔH = 87.9 kJ Will we have to do anything to any of the reactions? Which ones and what do we have to do? Example 2 Calculate the enthalpy of the following chemical reaction: CS2 (l) + 3 O2(g) CO2 (g) + 2 SO2 (g) a b c C (s) + O2 CO2 (g) 2 S (s) + 2 O2 2 SO2 (g) CS2 (l) C (s) + 2 S (s) Now add them all together! ΔH = -393.5 kJ ΔH = -593.6 kJ ΔH = -87.9 kJ Entropy (S) • A measure of randomness or disorder • Associated with probability (There are more ways for something to be disorganized than organized.) • Entropy increases going from a solid to a liquid to a gas. • Entropy increases when solutions are formed. • Entropy increases in a reaction when more atoms or molecules are formed. • The entropy of a substance increases with temperature. Entropy Law of Thermodynamics In any spontaneous process there is always an increase in the entropy of the universe. The energy of the universe is constant but the entropy of the universe is increasing. Spontaneous Process • occurs without outside intervention • may be fast or slow When the weather warms, the melting of ice is spontaneous. The 3rd Law of Thermodynamics • The entropy of a perfect crystal at 0 K is zero. • Soreaction = Soprod- Soreact • They are given in joules/K.mol.