Unit 4 Thermodynamics What is Energy? With your table group come up with a definition and examples of energy. Record on your group’s whiteboard

Unit 4 Thermodynamics: PS 3-3 Design, build, and refine a device that works within given constraints to convert one form of energy into another form of energy. a. Students know that energy can be characterized as potential or kinetic. b. Students know the law of conservation of energy and that energy can be converted from one form to another.

ENERGY: The ability to do work or produce heat Two basic forms: Potential: energy of composition or position of an object. Ex. water behind a dam or energy in bonds of a molecule 2. Kinetic: energy of motion Ex. water moving out of a dam or the movement of molecules

PE: Potential Energy KE: Kinetic Energy

Six Forms of Energy Chemical: potential energy stored in bonds For example coal, glucose, and petroleum Mechanical (work): energy of a moving object For example wind, and a car crash Thermal (heat): energy of the motion of molecules (friction) Example a cup of hot water

Six Forms of Energy Radiation (light): energy related to movement of light, electromagnetic waves, or particles For example visible light, x-rays, and sunshine. Electrical: energy of the movement of electrons For example a circuit, and lightning. Nuclear: energy stored inside the nucleus of an atom. For example fusion occurring in the sun

Law of conservation of energy: Energy can not be created or destroyed Energy can be converted.

Conversion of Energy

Energy used in moving a car Only 25% of potential energy becomes mechanical (kinetic) energy the rest is lost as heat (thermal energy).

Computer Activities Glencoe Activity PhET Activity

Unit 4 Thermodynamics PS 1-4: Develop a model to illustrate that the release or absorption of energy from a chemical reaction system depends upon the changes in total bond energy. a. Students know how to describe temperature and heat flow in terms of the motion of molecules (or atoms). b. Students know chemical processes can either release (exothermic) or absorb (endothermic) thermal energy.

Unit 4 Temperature and Heat

d. They are all sitting at the same ambient temperature so they will have the same average kinetic energy.

Unit 4 Energy Initial model Come up with a model to show what would happen if you take two equal amounts of water, one at 00C and one at 1000C and mix together. Remember models both show what happens and explain WHY! Show what happens at the macroscopic (visual) and particle levels

Heat: A form of energy that flows from a warmer object to a cooler object. The warmer object loses heat and becomes colder (temperature lowers) The colder object absorbs heat and become warmer (temperature rises)

Measuring Heat energy: Three major energy units: calorie: The amount of energy needed to raise the temperature of 1 gram of water 10C Joule (J), calorie (cal), Food Calorie (Cal) 1 calorie = 4.184 joules Food Calories are equivalent to kilocalories (Kcal) 1 Calorie = 1 Kcal= 1000 calories

Example #1 Convert 60.1 cal to joules 19

Example #1 Convert 60.1 cal to joules 1 cal = 4.184 joules 60.1 cal x 4.184 J = 1 cal = 251 J 19

Example #2 A reaction released 8650 joules, how many Food Calories of energy are released? 19

Example #2 A reaction released 8650 joules, how many Food Calories of energy are released? 1 cal = 4.184 joules, 1000 cal = 1 Cal 8650 J x _1 cal x 1 Cal = 4.184 J 1000 cal = 2.07 Calories 19

Temperature: A measure of the average kinetic energy of the particles in a sample of matter. Determines the DIRECTION of heat flow - We measure in K or 0C

PhET Energy Forms and Changes Do the PhET activity with your partners.

Virtually every reaction either releases or absorbs heat. Thermochemistry: Study of heat changes in reactions Law of conservation of energy: Energy is neither created nor destroyed within the universe The universe consists of the system and the surroundings. System: the reacting particles Surroundings: everything else including all non-reacting particles.

Chemical Energy: Is the energy within bonds in chemical compounds (form of potential energy) When bonds break they absorb energy (endothermic) When bonds form they release energy (exothermic)

Exothermic and Endothermic Reactions: All Chemical reactions require breaking bonds which absorbs energy and making bonds which releases energy. The overall change in energy results in a reaction that is either Exothermic or Endothermic. - always from the view of the system

Potential Energy of Reactants vs. Products

Products have less potential energy so Exothermic

Examples…. *It feels HOT because surroundings heat up Exothermic: A process that has an overall release of energy. Heat flows out of the system and Δ H is - *It feels HOT because surroundings heat up Reactants  Products + ENERGY Examples…. O2 + H2  H2O + energy SnCl2(s) + Cl2 (g)  SnCl4 (s) + 186KJ Burning Formation of compound from its elements Water condensing /Freezing Most Synthesis Reactions

Examples…. *It feels COLD because surroundings cool down Endothermic: A process that has an overall absorption of energy. Heat flows into the system and Δ H is + *It feels COLD because surroundings cool down Reactants + ENERGY  Products Examples…. Sunlight + CO2 + H2O  C6H12O6 + O2 Producing sugar by photosynthesis 2 H2O (l)  2 H2 (g) + O2 (g) Electrolysis Boiling Melting Most Decomposition reactions

Unit 4 Final model Go back to your energy model #1 paper and complete your final model. Remember the model is meant to show what would happen if you take two equal amounts of water, one at 00C and one at 1000C and mix together. Remember models both show what happens and explain WHY! Once you are done explain your model to your group partners!

Unit 4: Energy and Thermodyamics PS 1-4 Develop a model to illustrate that the release or absorption of energy from a chemical reaction system depends upon the changes in total bond energy

At the beginning of the reaction there are reactants with certain energy

In order for the reaction to occur Activation Energy needs to be reached. Activation energy (Ea) = minimum amount of energy to start a reaction. The more energy needed to start a reaction, the higher the activation energy.

At the end of the reaction there are products with certain energy

ΔH = Hproducts - Hreactants ∆H = difference in energy of reactants & products ΔH = Hproducts - Hreactants Time vs. energy Energy } Reactants DH Products Time - If ΔH is positive the reaction is Endothermic - If ΔH is negative the reaction is Exothermic

Exothermic vs. Endothermic

For example… The evaporation of water H2O (l)  H2O (g) Does this reaction require a little or a lot of energy to start? Does this reaction have a positive or negative ΔH? Is this reaction exothermic or endothermic?

For example… The formation of NH3 N2 + 3H2  2NH3 Does this reaction require a little or a lot of energy to start? Does this reaction have a positive or negative ΔH? Is this reaction exothermic or endothermic?

Quick Review What is energy? How is it measured? 2. What is it called when a reaction releases energy?____________ How does it feel?___________ 3. What does it mean to “heat something up”?

PS 3-1 Create a computation model to calculate the change in the energy of one component in a system when the change in energy of the other component(s) and energy flows in and out of the system are known.

Group Talk… If you want to make water boil, what do you have to do? What do you think is happening at the molecular level? When we get out of the shower we feel cold. Why is this? IT TAKES ENERGY TO VAPORIZE!

States of matter review… The phase of a material depends on the energy of its particles

Phase changes

Phase Change Diagram How phases of matter change with the addition of heat energy Note vaporization takes a lot more energy than melting.

From a solid to liquid (melting): From a liquid to gas (vaporization): How does energy play a role in phase changes? Label each process as exothermic or endothermic From a solid to liquid (melting): From a liquid to gas (vaporization):

From a gas to a liquid(condensation): From a liquid to a solid(freezing/fusion):

Checking for understanding: answer after changes of state movie In terms of heat energy, what happens when a solid changes to a liquid and then to a gas? 2. Would this be considered exothermic or endothermic? Why? 3. When does temperature increase? Why? 4. When does temperature remain constant? Why? https://blausen.com/en/video/changes-of-state/

Enthalpy (H) - Is the heat content of a system at constant pressure You can measure the enthalpy of reaction (Hrxn) This is negative if heat is released (exo) This is positive if heat is absorbed (endo)

Fusion: (melting)/freezing 00C Latent heat of fusion= (Hfus): There is a tool to calculate the amount of energy gained or released during phase changes: Fusion: (melting)/freezing 00C Latent heat of fusion= (Hfus): Energy required to melt a substance, or energy given off when a substance freezes Equation… Q=mHfus m=mass Hfus (H2O)= 334 j/g = 80 cal/g

Example 1: how much energy is released (in joules) when 50 grams of water freezes? Example 2: How many grams of water can you melt with 3,555 calories of energy?

vaporization: (boiling)/condensing 1000C Latent heat of vaporization= (Hvap) Energy required to boil a substance, energy released when a substance condenses Equation… Q=mHvap m=mass Hvap (H2O)= 2260 j/g =540 cal/g

Example 1: how much energy is needed (in calories) to boil 250 grams of water? Example 2: How much energy is given off when 5.0 moles of water (H2O) condenses? (hint, first convert mol  g)

Group Talk… If you want to make water boil, what do you have to do? What do you think is happening at the molecular level? When we get out of the shower we feel cold. Why is this? IT TAKES ENERGY TO VAPORIZE!

Partner Talk… It is a warm day and you go to the park. Is everything going to feel the same? Why/Why not?

PS 3-1 Create a computation model to calculate the change in the energy of one component in a system when the change in energy of the other component(s) and energy flows in and out of the system are known.

Specific Heat (Heat Capacity) Specific heat (c): the amount of heat to increase 1g of substance by 10C The larger the heat capacity the better at resisting changes in temperature! (the slower it will heat up

Specific Heat of some common substances: Iron (0.46 J/g0C) Plastic (~1 J/g0C) Paper (1.4 J/g0C) Water (4.18 J/g0C) Which of the above substances would heat up the fastest? WHY?  Which of the above substances would heat up the slowest (most resistant to heat change)? WHY?

Specific heat demonstration- soda bottle/blowtorch 1. What do you expect to happen to the water filled soda bottle when a blow torch is used on it? 2. What happened? 3. What do you expect to happen when the water is taken out of the bottle and the blowtorch is used? 4. Explain the different results using the term specific heat.

Heat Energy (Q) can be found using the equation… Q=mcT Q = energy change (J or cal) m= mass of substance being heated (grams or kg) C = specific heat capacity of substance (J/g0C or (cal/g0C) T= change in temperature (0C) cH2O= 1 cal/g0C or 1 kcal/kg 0C PAY ATTENTION TO UNITS!!!

Mass = 7.40 g Q = x Q= m Cp ∆T Q= (7.40) (4.184) (17.0) Q= 526J Example 1– Calculate the amount of heat energy (in joules) needed to raise the temperature of 7.40 g of water from 29.0°C to 46.0°C. Specific Heat of Water = 4.184 Mass = 7.40 g Q = x Temperature Change = 46.0°C – 29.0°C = 17.0°C Q= m Cp ∆T Q= (7.40) (4.184) (17.0) Q= 526J 22

Example 2: Calculate the amount of heat energy (in joules) needed to raise the temperature of 50.0 g of Iron from - 10.0°C to 5.0°C (heat capacity of iron is 0.46 j/g0C).

CALORIMETRY 1. A reaction takes place in a calorimeter during which 40.0 g of water is heated from 24.0oC to 50.0oC. Find the heat of reaction (∆H).

CALORIMETRY 2. The temperature of a piece of copper with a mass of 95.4 g increases from 25.0oC to 48.0oC when the metal absorbs 205.3 calories of heat. What is the specific heat of the copper?

HESS’ LAW If Equation (C) is the sum of equations (A) and (B), then: ∆H for (C) = ∆H for (A) + ∆H for (B) Ex. 1: Given: C(s) + ½ O2(g) → CO(g) ∆H = -26.4 kcal CO(g) + ½ O2(g) → CO2(g) ∆H = -67.7 kcal Find: ∆H for: C(s) + O2(g) → CO2(g)

HESS’ LAW Ex. 2 Calculate ∆H for: NO(g) + ½ O2(g)  NO2(g)

HESS’ LAW Ex. 3 Given: 1) Sn(s) + ½ O2(g)  SnO(s) ∆H = -68 kcal 2) SnO2(s)  SnO(s) + ½ O2(g) ∆H = 70 kcal   Calculate the heat of formation of SnO2.  

HESS’ LAW Ex. 4: Find ∆H for: CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) See Table 4.3 and H2O(l)  H2O(g) ∆H = 9.72 kcal

HESS’ LAW Ex. 5 Calculate ∆H for: 2 C(s) + H2(g)  C2H2(g) Given: C2H2(g) + 5/2 O2(g)  2 CO2(g) + H2O(l) ∆H = -310.6 kcal C(s) + O2 (g)  CO2(g) ∆H = -94.1 kcal H2(g) + ½ O2(g)  H2O(l) ∆H = -68.3 kcal  Calculate ∆H for: 2 C(s) + H2(g)  C2H2(g)  

EXO vs. ENDO Hold up “X” sign for exo, hold up “E” sign for endo Unit 4 Review EXO vs. ENDO Hold up “X” sign for exo, hold up “E” sign for endo

Are we exothermic or endothermic? Feels cold Endothermic

Are we exothermic or endothermic?

Are we exothermic or endothermic? Freezing (fusion) Exothermic

Are we exothermic or endothermic? RELEASES ENERGY Exothermic

Are we exothermic or endothermic?

Are we exothermic or endothermic?  H = -345 J Exothermic

Are we exothermic or endothermic? 2H2 + O2  2H2O + Heat Exothermic

Are we exothermic or endothermic? Feels hot Exothermic

Are we exothermic or endothermic?

Are we exothermic or endothermic?

Are we exothermic or endothermic? Condensation Exothermic

Are we exothermic or endothermic?  H = + 405 kJ Endothermic

Are we exothermic or endothermic? Absorbs energy Endothermic

Are we exothermic or endothermic? Melting Endothermic

Are we exothermic or endothermic? Lighting a match Exothermic

Are we exothermic or endothermic? CH4(g) + 2 O2 (g)  CO2(g) + 2 H2O(g) +heat Exothermic

Are we exothermic or endothermic? 2KClO3 + heat  2KCl + 3 O2 Endothermic

Are we exothermic or endothermic? 2HBr (g)  H2 (g) + Br2 (g) ΔH = + 103kJ Endothermic