Thermodynamics

New section in table of contents

Energy

 Energy Defined: Matter is stuff (atoms), Energy is the stuff that makes matter do stuff.

 What are some different types of energy you know of?

First Law of Thermodynamics

 KC 1: 1 st Law of Thermodynamics - energy can be neither created nor destroyed (conservation of energy).

Mechanical Energy

 Energy due to a object’s motion (kinetic) or position (potential).

Thermal Energy

 Heat Energy  The heat energy of an object determines how active its atoms are.  A hot object is one whose atoms and molecules are excited and show rapid movement.

 A cooler object's molecules and atoms will show less movement.

Electromagnetic Energy

 Light energy  Includes energy from gamma rays, x-rays, ultraviolet rays, visible light, infrared rays, microwave and radio bands

Electrical Energy

 Energy caused by the movement of electrons.

 Easily transported through power lines and converted into other forms of energy.

Nuclear Energy

 The energy released during nuclear fission or fusion. Fusion: Stars Fission: Nuclear Power Plants; Earth Core

Chemical Energy

 Energy that is available for release from chemical reactions.

Energy



Matter

Anything that has mass and takes up space  Anything made of atoms or subatomic particles.



Energy

Energy is of central importance in physics.  It is impossible to give a comprehensive definition of energy because of the many forms it may take.  KC 2: Energy is the ability to do work or produce heat.

Energy

 KC 3: Chemical potential energy is energy stored in a substance because of its composition  Chemical potential energy is important in chemical reactions

Chemical Energy

Temperature

 KC 4: Temperature is a quantity proportional to molecular kinetic energy.

 Temperature is the measure of the movement of atoms.  Average kinetic energy of molecules is measured in terms of temperature.

Rank the following pictures according to their average kinetic energy (low to high) A B C

Heat

 KC 5: Heat is a form of energy; that is in the process of flowing from a warmer object to a cooler object.

Chemical Energy Demos

2 more things to do

 Ticket out the door- Post-it Note with name  Write one complete sentence that shows how energy, temperature, and heat are related.

 Closure- Exit Pass- The system with the most heat will burn you the worst (200mL boiling water). Write the question below and think!

Why can dry ice, solid CO 2 burn you, but CO 2 in the gas phase cannot? Gases have more kinetic energy then liquids or solids.

Cold and Hot Packs

 How do you think instant hot and cold packs work? Keep in mind chemical reactions and energy.

Cold or Hot Pack?

 Experiment A:  A spoon full of sodium bicarbonate (NaHCO 3 ) with 2 squirts of water in a sandwich bag.  Experiment B:  A spoon full of calcium chloride (CaCl 2 ) with 2 squirts of water in a sandwich bag.

KC 6: Exothermic Reaction chemical reaction that releases energy in the form of light or heat.

is a

Activation Energy Energy of reactants Energy of products Reactants

-

D

H

Products

•

Reaction Progress

KC 7: Endothermic Reaction absorption of heat.

chemical reaction accompanied by the

Activation Energy Energy of reactants Energy of products Products

+

D

H

Reactants Reaction progress

Complete the sentences below for both experiment using the correct word.

1.

2.

3.

4.

5.

Experiment ____ is similar to ____(hot /cold pack) The salt which dissolves in water is_____.

The heat energy is _____(given off/ taken in).

The temperature _____ (falls/rises). The reactions is _____ (endothermic/exothermic).

Chemical Energy and the Universe

 Chemists are interested in changes in energy during reactions.

 Enthalpy is a measure of the total energy in a chemicals bonds.

 KC 8: Enthalpy= heat of reaction= ΔH is the change in enthalpy during a reaction symbolized as ΔH rxn .

ΔHrxn = H final – H initial ΔHrxn = H products – H reactants

KC 9: Endothermic Reaction = positive ∆H

KC 10: Exothermic Reaction = negative ∆H

Enthalpy of a Reaction

 KC 11: ΔHrxn = ∑n∆H products – ∑n∆H reactants

Example

 KC 12: Use the table to find the ∆H rxn for the following reaction: SO 2 (g) + NO 2 (g) → SO 3 (g) + NO(g) Substance SO 2 (g) NO 2 (g) SO 3 (g) NO(g) ∆H f (kJ/mol) -296.8

33.1

-395.8

90.3

Enthalpy of Reaction

 There will NOT be values given for substances that are in their “natural” state  Example – oxygen: It does not take energy to make O 2 (g) carbon: It does not take energy to make C(s)  These heats of formation (∆H f ) will be zero  You MUST know this!!

Specific Heat

 KC 13: The specific heat of any substance is the amount of heat required to raise one gram of that substance one degree Celsius.

 Some objects require more heat than others to raise their temperature.

Specific Heat

 Intrinsic Property  It is independent of how much of the material is present.

Measuring Heat

Measuring Heat

 KC 14: A calorie is defined as the amount of energy required to raise the temperature of one gram of water one degree Celsius.

 Food is measured in Calories, or 1000 calories (kilocalorie).

 KC 15: A joule is the SI unit of heat and energy, equivalent to 0.2390 calories.

Specific Heat of Water

 Life on earth evolved in water, and all life still depends on water.  KC 16: Water has a specific heat capacity of 4.2 J/ g o C, which means that it takes 4.2 joules of energy to heat 1g of water by 1 o C.  KC 17: This is unusually high and it means that water does not change temperature easily.

KC 18: Calculating Heat

q = c × m × ΔT  q = heat absorbed or released  c = specific heat of substance  m = mass of substance in grams  ΔT = change in temperature in Celsius (Final – Initial)

Calculating Heat

 KC 19: A 12.5g sample of aluminum was originally at 20°C. It was heated to a temperature of 34.3°C. How much heat was absorbed by the sample? (specific heat of aluminum = 0.9J/g°C)

Calculating Heat

 KC 20: A sample of zinc was heated from an initial temperature of 234°C to 439°C. There was 198.5J of heat absorbed and the specific heat of zinc is 0.388J/g°C. What was the mass of the sample?

Calculating Heat

 KC 21: Calculate the amount of heat needed in calories to raise 23g of water from 49°C to 236°C.

Closure – Exit Pass

 Why did one ice cube melt faster than the other?

 You must use the correct vocabulary in your response.

Calorimetry

 How do scientists measure the heat of reactions?

 Scientists use a calorimeter to measure the temperature change of a given reaction  They can then use q=mc∆T to calculate the heat of that reaction

Calorimetry

 KC 22: Calorimetry – using temperature change to calculate the heat of a reaction as it moves from reactants to products

Calorimetry

 This process of calorimetry is how scientists measure the amount of calories in food  You will be doing a similar process to rank different snack foods from smallest kcal/g to largest amount of kcal/g  Video

Calorimetry Example

 Calculate the number of calories absorbed by the water, based on the information below.

Mass of empty beaker: 300g Temperature of water before reaction: 27.0 ˚C Mass of beaker with H 2 O: 465g Temperature of water after reaction: 65.0 ˚C

Lab – Burning Food

      Measure approximately 15mL of water into the can. Record the mass of the water in data table (NOTE: 1mL of water = 1g) Measure and record the initial temperature of the water in the can then remove the thermometer.

Place the food sample on the paper clip and record its mass in the data table.

Quickly ignite the food sample with a Bunsen burner and place the platform under the test tube. Place the thermometer in the can. Allow the food to burn completely.

Measure and record the highest temperature achieved by the water in the can during the burning of the food.

Record the mass of the food sample after burning and record in data table.

Data Table

Food Sample: Mass of water Initial temp of water Initial mass of food sample (this also includes the paper clip and cork) Final temp of water Final mass of food sample (this also includes the paper clip and cork) Change in temp of water Change in mass of food sample q of water (from calculation) Kcal/g of food sample (from calculation) Peanuts Doritos Cheetos Pretzels

Hess’s Law

 KC 23: Hess’s law – the amount of heat released or absorbed in a chemical reaction does not depend on the number of steps in the reaction  The ∆H rxn reactions can be determined from using other  This is like systems of equations: 2x + 3y = 23 6y + x = -45

Hess’s Law

 RULES:  Flip an equation, change the sign on the ∆H  Multiply by a factor, multiply the ∆H by the same factor (can be a whole number or a fraction, ½, ¼ )  Add equations down to get the goal reaction, add ∆H’s down to get the ∆H for the goal reaction

KC 24: Example Hess’s Law

Reaction Reaction ∆H 2 C(s) + O 2 (g) → 2 CO(g) -221kJ C(s) + O 2 (g) → CO 2 (g) -393kJ C(s) + CO 2 (g) → 2CO(g) ?

∆H

BRN – find ∆H

C 2 H 5 OH(l) + 2 O 2 (g) --> 2 CO 2 (g) + 2 H 2 O(l) C (s) + O 2 (g) --> CO 2 (g) -875. kJ -394.51 kJ H 2 (g) + ½ O 2 (g) --> H 2 O (l) -285.8 kJ 2H 2 (g) + 2C(s) + O 2 (g) --> C 2 H 5 OH(l) ?

Lab Reminders

 Keep the water boiling while you are doing the experiment  Be careful when placing the metal in the test tube (I’m going to give you 2 samples)  Leave test tube in boiling water for 5-10min  Use test tube holder to pour metal into calorimeter  Don’t forget to record initial temperature of water  Use about 15-20mL of water in the calorimeter  Record highest temperature of water reached for final temp  Assume initial temp of metal is 100°C

Enthalpy

 KC 25: enthalpy is the measure of the energy that is released or absorbed by the substance when bonds are broken and formed during a reaction  When bonds are formed, energy is released  When bonds are broken, energy is absorbed  Exothermic reaction: -∆H  Endothermic reaction: +∆H

Entropy

 KC 26: entropy (∆S) is a measure of the randomness or disorder of the system; the greater the disorder of a system, the greater its entropy  As molecules or ions become more dispersed, their disorder increases and their entropy increases Reaction NaCl(s) → Na + (aq) + Cl (aq) 2Na(s) + Cl 2 (g) → 2NaCl(s) Entropy change ∆S +43 -181

Gibbs Free Energy

 KC 27: Gibbs free energy (∆G) is a measure of the spontaneity of the process  KC 28: A spontaneous reaction is one that does occur or is likely to occur without continuous outside assistance, such as input of energy  A reaction is spontaneous if the Gibbs energy change is negative  If a reaction has a ∆G greater than 0, the reaction is nonspontaneous

Gibbs Free Energy

 Reactions that have large negative ∆G values often release energy (-∆H) and increase disorder (+∆S)  ∆G = ∆H - T∆S  Example: 2 K(s) + 2 H 2 O(l) → 2K + (aq) + 2 OH (aq) + H 2 (g) ∆H = -392kJ, ∆S = 0.047kJ/K, T = 298K

KC 29: Spontaneity (thermodynamically favored)

∆H + _ + + ∆S + T Low High Low High Low High Low High ∆G + + + + Spontaneous?

Yes No No Yes Yes No

Gibbs Free Energy

 KC 30: Given that the changes in enthalpy and entropy are -139kJ and 277J/K respectively for the reaction given below, calculate the change in Gibbs energy. Then, state whether the reaction is spontaneous at 25°C.

C 6 H 12 O 6 (aq) → 2 C 2 H 5 OH(aq) + 2 CO 2 (g)

Practice

 Which of the following will be true when a pure substance in liquid phase freezes spontaneously?

A.

∆G, ∆H, and ∆S are all positive B.

∆G, ∆H, and ∆S are all negative C.

∆G and ∆H are negative, but ∆S is positive D.

∆G and ∆S are negative, but, ∆H is positive

Practice

 Which of the following pairs of conditions will favor a spontaneous reaction?

A.

A decrease in entropy and a decrease in enthalpy B.

A decrease in entropy and an increase in enthalpy C.

An increase in entropy and a decrease in enthalpy D.

An increase in entropy and an increase in enthalpy

Review Video - closure

 Entropy  3, 2, 1  3 things you learned  2 things you found interesting  1 question you have

Bellringer

 What three phases can H 2 O exist?

 Can liquid water that is boiling ever get hotter than 100 C?

 What are the bubbles made of when water boils?

Heating Curve

Heating Curves

•

liquid

•

solid

•

added energy

•

gas

Heating Curves

boiling/ cond. pt melting/ freezing pt

•

liquid

•

solid

•

added energy

•

gas

Heating Curves

boiling/ cond. pt melting/ freezing pt

•melting/freezing •occurring here •boiling/condensing •occurring here •

liquid

•

solid

•

added energy

•

gas

 C ice = 2.09J/g°C  C water = 4.184J/g°C  C steam = 2.01J/g°C ∆H fus = 334J/g ∆H vap = 2260J/g

q=mc∆T q = m∆H

A 12g sample of ice initially at -5°C is heated to a temperature of 95°C. Calculate the amount of heat that is absorbed during this process.

C ice = 2.09J/g°C ∆H fus = 334J/g C water = 4.184J/g°C ∆H vap = 2260J/g C steam = 2.01J/g°C

A 3.5g sample of water at 13.6°C is heated to a temperature of 112°C. Calculate the amount of heat absorbed during this process.

C ice = 2.09J/g°C ∆H fus = 334J/g C water = 4.184J/g°C ∆H vap = 2260J/g C steam = 2.01J/g°C

Cooling Curves

Removal of energy

Heating and Cooling Curve Assignment

 You will be given 2 element cards to create a heating and cooling curve. The melting points and boiling points are on the card for each element.

 You will need to draw a heating curve for one and a cooling curve for the other.

Closure: Jeopardy

d D e a A b B c C

 Answer- This letter Indicated when the matter is a solid.

 Answer- 95°C  Answer- 70°C