Thermodynamics.

Thermodynamics

Intro to Thermodynamics
Almost all chemical reactions involve a change in energy (usually in the form of heat) between the reactants and its products Thermo – Dynamics –

Energy Energy- is the ability to do work or produce heat
Heat – is the transfer of energy between two objects of different temperatures

Thermodynamics Thermochemistry- is the study of heat change that accompany chemical reaction and phase changes. Heat absorbed: Endothermic Heat released: Exothermic

Thermodynamics Heat (q) flows from ___hot___ to __cold__*
Measured in joules (J) which is the SI unit of heat energy as well as all other forms of energy Other units: kJ, cal, Cal

Conversions 1000 J = 1 kJ 1 calorie = 4.184J
1 kcal 1000 cal = 1 Cal (food calorie)

Practice Problems A fruit and oatmeal bar contains 142 Calories. Convert this energy to calories. Convert 256 J to kcal. The breakfast I ate this morning contains 230 nutritional Calories. How much energy in joules will this breakfast supply?

Thermodynamics Temperature: A measure of the average kinetic energy of the atoms or molecules of substance Units: Celsius, Kelvin and Fahrenheit

Difference between Heat and Temperature
Heat is the transfer of thermal energy from a higher temperature to a lower temperature Thermal (heat) energy: the sum total of the kinetic energies of the particles in the sample of matter. Measured in Joules (J)

Temperature continued
Temperature measures the AVERAGE kinetic energy The faster the molecules are moving the higher the temperature The slower the molecules are moving lower the temperature

Temp vs. Heat Compare and contrast temperature and thermal energy of the Pacific Ocean to a boiling pot of water

Temp vs. Heat ANSWER Compare and contrast temperature and thermal energy of the Pacific Ocean to a boiling pot of water The Pacific Ocean would be at a lower temperature (average KE of the particles), but would have a lot more thermal energy (total KE of all the particles)

Enthalpy Heat of reaction (ΔH): The quantity of heat released or absorbed during a reaction ΔH = Hproducts – Hreactants Important to know the formula above, but you will NOT do any math with it!

ΔH Each reaction will have an enthalpy change which indicates whether the reaction is endothermic or exothermic.

The sign of the enthalpy of reaction:
If ΔH is negative then heat is being released and the reaction is exothermic Ex: H2(g) + Cl2(g)  2HCl(g) H = -185,000 J If ΔH is positive then heat is being absorbed and the reaction is endothermic Ex: CO2(g) + 2H2O(g)  2O2(g) + CH4(g) H = 890,000 J To determine if the reaction is exothermic or endothermic, you must know the value of ΔH

Endothermic vs. Exothermic

Example Rewrite the following equation including the ΔH value as a reactant or product and indicate if it is exothermic or endothermic H2 + ½ O2 H2O ΔH = kJ

You try it! Rewrite the following equation including the ΔH value as a reactant or product and indicate if it is exothermic or endothermic H2O  H2 + ½ O2 ΔH = kJ

Enthalpy ΔH is an extrinsic physical property because it depends on the amount present! Determine the ΔH for the following reactions: N2 + 3H2  2NH3 H= kJ ½ N2 + 3/2 H2  NH3 H = ______________ kJ 2NH3  N2 + 3H2 H = ______________ kJ

Phase Changes Phase change is the change of state of a substance (solid, liquid, or gas) To change phases, heat must be absorbed or released, but temperature remains constant.

Freezing (Solidification)
States of matter Energy Endo or Exo? H Sign Melting Freezing (Solidification) Evaporation Condensation

States of matter Energy Endo or Exo? H Sign Melting Solid to liquid Needs to be added (heat it up) Endo + Freezing (Solidification) Evaporating Condensation

States of matter Energy Endo or Exo? H Sign Melting Solid to liquid Needs to be added (heat it up) Endo + Freezing (Solidification) Liquid to solid Needs to be removed (cool it down) Exo - Evaporating Condensation

States of matter Energy Endo or Exo? H Sign Melting Solid to liquid Needs to be added (heat it up) Endo + Freezing (Solidification) Liquid to solid Needs to be removed (cool it down) Exo - Evaporating Liquid to gas Condensation

States of matter Energy Endo or Exo? H Sign Melting Solid to liquid Needs to be added (heat it up) Endo + Freezing (Solidification) Liquid to solid Needs to be removed (cool it down) Exo - Evaporating Liquid to gas Condensation Gas to liquid Exo

Let’s review Kinetic Energy – Energy of motion, directly relates to the temperature of the substance Potential energy – stored energy, the energy stored in the bonds

During a phase change Temperature stays the same
Because kinetic energy stays the same As potential energy increases or decreases depending on the phase change Solid  liquid or liquid  gas: PE increases Gas  liquid or liquid  solid: PE decreases

What does heat do? If a substance is heated, the heat can do ONE of these two things: Increase the temperature (Kinetic energy) OR Change the phase (Potential energy) Key idea: Temperature remains CONSTANT during phase changes!

Heating Curve

Definition Sign and energy movement
ΔHfus -Hfus ΔHvap -ΔH vap * Table 16-6 on page 502 in blue book for standard enthalpies of vaporization and fusion *

ΔHfus Heat required to melt one gram of a solid Positive since E is absorbed and PE increases -Hfus ΔHvap -ΔH vap * Table 16-6 on page 502 in blue book for standard enthalpies of vaporization and fusion *

ΔHfus Heat required to melt one gram of a solid Positive since E is absorbed and PE increases -Hfus Heat released to solidify one gram of a liquid Negative since E is released and PE decreases ΔHvap -ΔH vap * Table 16-6 on page 502 in blue book for standard enthalpies of vaporization and fusion *

ΔHfus Heat required to melt one gram of a solid Positive since E is absorbed and PE increases -Hfus Heat released to solidify one gram of a liquid Negative since E is released and PE decreases ΔHvap heat required to vaporize one gram of a liquid Positive since E is absorbed and PE increases -ΔH vap * Table 16-6 on page 502 in blue book for standard enthalpies of vaporization and fusion *

ΔHfus Heat required to melt one gram of a solid Positive since E is absorbed and PE increases -Hfus Heat released to solidify one gram of a liquid Negative since E is released and PE decreases ΔHvap heat required to vaporize one gram of a liquid Positive since E is absorbed and PE increases -ΔH vap heat released to condense one gram of a gas * Table 16-6 on page 502 in blue book for standard enthalpies of vaporization and fusion *

Summary

Phase Change Problems q = mΔH
Depending on the phase change you will either use ΔHvap or ΔHfus

Practice Problem Calculate the heat required to melt 25.7 g of solid water at its melting point.

Practice Problem How much heat is evolved when 275 go of water vapor condenses to a liquid at its boiling point?

Specific Heat amount of heat required to raise the temperature of one gram of a substance by one degree Celsius Units : J/g˚C You can find specific heat values in the reference table

Think about it Which has a higher specific heat? Water or sand
Metal pan or oven mitts

Think about it Which has a higher specific heat? Water or sand
Metal pan or oven mitts Why? The more difficult it is to heat something up, the higher the specific heat because more heat is required to raise the temperature of that substance

Specific Heat-the formula!
q= mcΔt Specific Heat J/g °C Change in Temperature (K or ° C) Tf - Ti Mass (g) Amount of Heat (J)

Specific Heat Substance Specific Heat (J/g°C) Water 4.184 Ice 2.03
Steam 2.01 Al (s) 0.897 Granite (s) 0.803 Iron (s) 0.449

Practice Problem 1 How much heat is lost when 4110 g of aluminum cools from 660˚C to 25.0˚C?

Practice Problem 2 Calculate the heat required to heat up 124 g of water from 17.5˚C to 25.0˚C.

Practice Problem 3 How much heat is required to heat 15.8 g of liquid water at 30.˚C to water vapor at 122˚C?

Practice Problem 4 How much heat is released when 21.2 g of water cools from 113.7˚C to -24.8˚C?

Calorimeter qlost=-qgained
Insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process qlost=-qgained mc∆Tlost = -mc∆Tgained

One kind of calorimeter works like this:
A known amount of reactants are sealed in a reaction chamber. The reaction chamber is immersed in a known quantity of water The water and reaction chamber are in an insulated vessel The energy given off (or absorbed) during the reaction is transferred to the water If the reaction released energy the water temperature increases If the reaction needs energy the water temperature decreases. Heat is not measured directly. Temperature is measured. Temperature is affected by the transfer of heat (energy).

Example A 20.0 g piece of metal at a temperature of 90.0˚C is dropped into an insulated container holding 125 g of water at 20.0˚C. If the final temperature is 23.0˚C, what is the specific heat of the metal?

Hess’s Law The overall enthalpy change in a reaction is equal to the sum of enthalpy changes for the individual steps in the process.

Key ideas The major utility of Hess’s Law is in calculating the enthalpy changes of reactions that would be difficult to measure. Note: according to the law of conservation of energy, energy can neither be created nor destroyed in an ordinary chemical reaction. Hess’s law tells us that we will never get more energy (or less energy for that matter) from a chemical reaction.

Example

Entropy Entropy (S) – is the measure of disorder or randomness of the particles that make up a system. The Law of Disorder –states that spontaneous processes always proceed in such a way that the entropy of the universe increases. (The second law of Thermodynamics)

Entropy From each group, select the part with the highest entropy:
Increasing temperature or decreasing temperature of a substance Reactants or products: 2NH3(g)  3H2(g) + N2(g) ice or water or steam

Entropy Practice Predict the sign of ∆Ssystem for each of the following changes. CIF(g) + F2(g)  CIF3(g) NH3(g)  NH3(aq) CH3OH(l)  CH3OH(aq) C10H8(l)  C10H8(s)

Entropy and the Phases of Matter
Place the phases of matter in order of increasing entropy: Solid  liquid  aqueous  gas

Quiz Topics Exothermic vs. Endothermic Enthalpy
Specific Heat (concepts and calculations) Phase changes (concepts and calculations) Heat curves (interpreting and calculations) Calorimetry problems Hess’s Law Entropy

Driving Forces of a Reaction
There are two factors that determine whether a reaction will occur spontaneously: Enthalpy – heat absorbed in a chemical reaction Tendency in nature  lower energy so exothermic reactions are favored (-H) Entropy – the measure of randomness of the particles in a system Tendency in nature  more disorder to positive changes in entropy are favored (+S)

Gibb’s Free Energy Measure of the spontaneity of a reaction
Negative ∆G means the reaction is spontaneous Positive G means the reaction is NOT spontaneous The tendency in nature  less free energy (-∆G), spontaneous reactions are favored ∆G = ∆H - T∆S T = Kelvin Temperature

∆G = ∆H - T∆S H S G Spontaneous? + -

Try it! NH4NO3(s) + heat  NO3-(aq) + NH4+(aq)

Kinetics and Equilibrium

Collision Theory Reacting substances (atoms/ions/molecules) must collide with: sufficient energy. (Activation Energy = Ea) the correct orientation.

Reaction Mechanism For: H2(g) + I2(g) ↔ 2 HI(g) Step 1: I2 ↔ 2I-
Many reactions actually take place through a series of steps involving two particle collisions. This step by step sequence is called a reaction mechanism For: H2(g) + I2(g) ↔ 2 HI(g) Step 1: I2 ↔ 2I- Step 2: I- + H2 ↔ 2HI

The forward reaction is endothermic
The reverse reaction is exothermic

Chemical Kinetics The area of chemistry that is concerned with reaction rates and reaction mechanisms Reaction rate: the change in concentration of reactants per unit time as a reaction proceeds

Factors that affect reaction rates
Nature of reactants More reactive the substance, the faster the reaction Surface area More surface area  quicker reaction Temperature Usually increasing temperature  increases rate Concentration When increased, the rate increases (more collisions) Catalysts Reaction takes a different path and lowers activation energy

Catalyst A substance that increases the rate of a chemical reaction without itself being consumed in the reaction. An example of a catalyst is an enzyme

Equilibrium Chemical equilibrium is a state in which the forward and reverse reactions balance each other because they take place at equal rates. Rateforward reaction = Ratereverse reaction Reversible Reaction is one that can occur in both forward and reverse directions.

Equilibrium Equilibrium is DYNAMIC, meaning that Equilibrium is a state of action. The reaction is constantly shifting and moving to maintain equilibrium even though we can not observe this.

Equilibrium Expression
aA + bB ↔ cC + dD Where A and B are the reactants, C and D are the products, and a, b, c, and d are the coefficients in a balanced equation. 𝐾𝑒𝑞= [𝐶] 𝑐 [ 𝐷] 𝑑 [𝐴] 𝑎 [ 𝐵] 𝑏

What is Keq A constant for each reversible reaction (at a given temperature) that relates the products present in a reaction at equilibrium to the reactants present

Formula Ignore all pure solids and pure liquids in the reaction. (Why?) Only the concentrations of substances that can actually change are included in K. This means that pure solids and liquids are omitted because their concentrations cannot change. Make a ratio of [products] / [reactants] The concentrations are all expressed as Molarity (mol/L) Make all coefficients in the balanced chemical equation into exponents in the ratio.

Try it! Find the Equilibrium Expression for the following:
N2O4(g) ↔ 2NO2(g) CO(g) + 3H2(g) ↔ CH4(g) + H2O(g) Ca(OH)2(s) + H2O(l)  Ca2+(aq) + 2OH-(aq) CaCO3(s) ↔ CaO(s) + CO2(g) 2H2(g) + O2(g)  2H2O (l)

Example 1: The following equilibrium concentrations were observed for the Haber process at 127°C. The Haber process is: N2(g) + 3H2(g)↔ 2NH3(g). Calculate ‘K’. [NH3] = 2.0 mol/L [N2] = 1.0 mol/L [H2] = 2.0 mol/L

Example 2 A different experiment was performed using the same process (Haber) as above and at the same temperature, but different concentrations. Calculate ‘K’. [NH3]=0.86 mol/L, [N2] = 0.75 mol/L, [H2] = 1.25 mol/L

Same K value between examples?
The only way to change K for a given reaction is to change the temperature!

Equilibrium Keq = 1: Amount reactants roughly equals products at equilibrium Keq > 1 : More products than reactants at equilibrium. Keq < 1 : More reactants than products at equilibrium.

Ionization of an Acid Ka
The larger this number, the stronger the acid! CH3COOH(aq) + H2O(l) ↔ H3O+(aq) + CH3COO-(aq)

Ionization of an Acid Ka
The larger this number, the stronger the acid! CH3COOH(aq) + H2O(l) ↔ H3O+(aq) + CH3COO-(aq) 𝐾 𝑎 = 𝐻 3 𝑂 + [𝐶 𝐻 3 𝐶𝑂 𝑂 − ] [𝐶 𝐻 3 𝐶𝑂𝑂𝐻]

Ionization of an Base Kb NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq))

Ionization of an Base Kb NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq)
𝐾 𝑏 = 𝑁𝐻 4 + [ 𝑂𝐻 − ] [𝑁 𝐻 3 ]

Solubility of a Slightly Soluble Salt
Ksp The larger the Ksp, the more it dissociates AgCl(s) ↔ Ag+(aq) + Cl-(aq)

Solubility of a Slightly Soluble Salt
Ksp The larger the Ksp, the more it dissociates (more soluble) AgCl(s) ↔ Ag+(aq) + Cl-(aq) 𝐾 𝑠𝑝 = 𝐴𝑔 + [ 𝐶𝑙 − ]

Le Chatelier

Le Chatelier’s Principle
A principle stating that if a stress is applied to a system in equilibrium, the equilibrium will shift to alleviate the effect of the stress.

A “stress” can be: Change in concentration of product or reactant
Change in temperature Change in pressure (just gases)

Changing Concentration
Equilibrium will shift AWAY from an increase in concentration and TOWARD a decrease in concentration.

This reaction is at equilibrium
heat NH3 N2 H2 What happens to the reaction if you increase the concentration of Nitrogen?

Increase nitrogen concentration
heat NH3 N2 H2 How does the reaction need to shift to make it go back to equilibrium? What will happen to all of the substances as a result? H2 will decrease, NH3 will increase Right

How have things changed?
heat NH3 N2 H2 How have things changed? heat NH3 H2 N2

Example 2 2SO3(g) ↔ 2SO2(g) + O2(g) Describe what happens when:
You decrease SO2 You increase O2

Changing Temperature If endothermic, heat is added to the product side
If exothermic, heat is added to the reactant side Keq (equilbrium constant): is only affected by a temperature change. Concentration and pressure do NOT change Keq Treat heat the same way you do concentration changes. If temperature increases, heat increases!

heat NH3 N2 H2 What happens when you increase the temperature? heat NH3 N2 H2 How does the reaction need to shift to make it go back to equilibrium? What will happen to all of the substances as a result? N2 and H2 increase, NH3 decreases Left

How have things changed?
heat NH3 N2 H2 How have things changed? heat NH3 N2 H2

Try it 2SO3(g) + heat ↔ 2SO2(g) + O2(g) Increase the temperature
Decrease the temperature

Changing Pressure Only affects equilibrium reactions with gaseous substances Think of your reaction in a closed container. When you increase pressure (decrease volume) you want to have less gas in your system When you decrease pressure (increase volume) you want to produce more gas in your system

Changing Pressure Rule Thumb
When you INCREASE pressure: always shift toward fewer moles of gas! When you DECREASE pressure: always shift toward more moles of gas!

Decrease Pressure Normal Increase Pressure

Try It! Ex 1: N2(g) + 3H2(g)  2NH3(g) Increase the pressure
Decrease the pressure Ex 2: 2SO3(g) ↔ 2SO2(g) + O2(g)

Stress Shift [CH4] [H2S] [CS2] [H2] Keq
CH4(g) + 2H2S(g) ↔ CS2(g) + 4H2(g) H = kJ/mol Stress Shift [CH4] [H2S] [CS2] [H2] Keq Increase CS2 Decrease H2 Increase Temp Increase Pressure

Stress Shift [CH4] [H2S] [CS2] [H2] Keq
CH4(g) + 2H2S(g) ↔ CS2(g) + 4H2(g) H = kJ/mol Stress Shift [CH4] [H2S] [CS2] [H2] Keq Increase CS2 Left Increase Increase ----- Decrease No Change Decrease H2 Right Increase Temp Right Changes! Increase Pressure Left

Test Topics Thermo Vocab Enthalpy (and Exo vs. Endo)
Specific Heat (concepts and calcs) Phase Changes (concepts and calcs) Heat Curves (labeling, concepts, and calcs) Entropy Gibb’s Free Energy (spontaneity) Energy Conversions Hess’s Law Collision Theory Potential Energy Diagrams Equilibrium concepts Equilibrium expressions Le Chatelier’s Principle Acid/Base Stuff

Thermodynamics.

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