TOPIC 5: ENERGETICS & THERMOCHEMISTRY 5.1 MEASURING ENERGY CHANGES Mrs. Page

UNDERSTANDINGS  Heat is a form of energy.  Temperature is a measure of the average kinetic energy of the particles.  Total energy is conserved in a chemical reaction.  Chemical reactions that involve the transfer of heat between the system and the surroundings are described as endothermic or exothermic.  The enthalpy change H (∆H) for a chemical reaction is indicated in kJmol -1  ∆H values are usually expressed under standard conditions, known as ∆H°, including standard states

APPLICATION & SKILLS  Calculation of the heat change when the temperature of a pure substance is changed using q=mc∆H  A calorimetry experiment for an enthalpy of reaction should be covered and the results evaluated. NOS  Fundamental Principal – conservation of energy  Making careful observation – measurable energy transfers between systems and surroundings

Thermodynamics  The study of energy and how it is converted between different forms.  First Law of Thermodynamics (Law of conservation of energy): Energy cannot be created or destroyed, it can only be converted to different forms.  The total amount of energy in a given system is conserved  In a chemistry: potential energy is stored in the bonds & the temperature of the reaction is an indication of the kinetic energy of the particles

Thermodynamics  Heat(q)moves from warmer to cooler objects  There are 3 types of heat transfers:  Conduction – particles in direct contact  Convection – through a fluid (liquids or gases)  Radiation – through space by electromagnetic waves  Absolute zero (0K) – all motion of particles stops and entropy (the amount of disorder in the system) is minimal.  As temperature increases, kinetic energy increases

SYSTEMS  A system is the part of the universe that we are studying.  Everything else is called the surroundings.  In chemistry the system is usually the reaction mixture taking place in a test tube/beaker/flask. While the surroundings include the test tube/flask/beaker, measuring devices, and air surrounding the system.  There are 3 types of systems:  Open Systems: transfer matter and energy across the boundary  Closed Systems: transfer energy but not matter across the boundary  Isolated Systems: do not transfer matter or energy across the boundary (matter can only move around within the system)

Exothermic Reactions  Give heat to surroundings (releasing energy)  Temp. inside vessel goes UP  Examples: Making chemical bonds, combustion of fuels, mixing alkalis and acids

Endothermic Reactions  Takes heat from surroundings (energy is required)  Temp. inside vessel goes DOWN  Examples: Breaking chemical bonds, photosynthesis, melting ice I-Nq3Os

Enthalpy Change (of a system)  The amount of heat energy taken in/given out in a chemical reaction (at constant pressure)  You cannot measure the enthalpy (H) of a system, only the change in enthalpy (  H) from initial to final state  Exothermic:  H is negative  Endothermic:  H is positive

Enthalpy Level Diagrams  Show change in enthalpy between reactants and products  No scale on y axis because you can’t measure initial and final H values

Exothermic Reaction Example  Product always lower energy (enthalpy)  Products more stable  Negative  H

Endothermic Reaction Example  Product always higher energy (enthalpy)  Heat energy must be put into the system  Products less stable  Positive  H

Your Turn

Two Types of Stability  Thermodynamic Stability: lower energy in relation to enthalpy change  Kinetic Stability: the speed a reaction occurs (Topic 6)  You can have thermodynamic stability but it may occur incredibly slowly (ex: a diamond changing into graphite)

Activation Energy  The minimum amount of energy colliding particles must have before the collision results in a reaction.  This barrier must be overcome for the reaction to occur.  Higher Activation Energy = slower reaction

Important Notes   H does not tell anything about the speed of a reaction  Exothermic reaction tend to occur spontaneously  Endothermic reactions require adding energy (must overcome the activation energy)

Calorimeters  An apparatus used to measure the amount of heat being exchanged between a system and the surroundings.  Systemic Error:  Heat loss to surroundings

Specific Heat Capacity (c)  The energy required to raise the temperature of 1g of substance by 1K Note: ∆1K = ∆1  C  Units: kJkg -1 K -1  For example: Aluminum has c = 0.90 kJkg -1 K -1 This means that if 0.90 kJ of heat energy is added to 1 kilogram of Al the temperature will raise 1K  Substances with higher c are more difficult to heat up.  Specific heat capacity (c) is not affected by the size of they system.  The specific heat capacity of water is 4.18 kJkg -1 K -1

FORMULA

You Try How much heat is released when 10.0kg of copper with a specific heat capacity of kJkg -1 K -1 is cooled from 85.0K to 25.0K?

Applying Specific Heat 40.0 grams of an unknown metal is heated to 91.3  C and then plunged into 100.0g of water at 21.3  C. The water and the metal reach a common temperature of 28.4  C. Given that the specific heat capacity of water is 4.18Jg -1  C -1, calculate the specific heat capacity of the metal. Step 1: Find q for water Step 2: Find q for metal Step 3: The energy to heat the water comes from the cooling of the metal so q water = q metal c= 1.18Jg -1  C -1

CANDLE LAB  Mass of Candle Before & After  Mass of water (D = 1 g/ml)  Temperature of water Before & After  What factors will affect heat transfer?  Distance between candle and beaker?  What assumptions must we make?  All heat transferred from candle to water  q water = q candle  Calculate the q/gram candle  J/mass  Discuss sources of error