Thermodynamics

The study of heat and energy

So far…

• We’ve shown that matter (solids, liquids, etc) can be changed and rearranged in chemical reaction.

• However, we must also consider that energy is exchanged during all chemical and physical changes.

Energy

• Energy: The capacity (ability) to do work and/or supply heat. Exists in many different forms.– Chemical potential energy is the type of

energy stored in chemicals and released/absorbed in reactions.

– Gasoline, nitroglycerine, etc all have high chemical PEs.

Movement of energy:

• When describing the flow of energy, it is useful to set a few boundaries:

• First, we will be examining what is known as a system, which will almost always be described by a chemical equation.

System vs Surroundings

• A system is the part of the universe that we are studying while EVERYTHING else is known as the surroundings.– In your notebooks, give me at least 2

examples of a system and its surroundings.• For example: a medical study would take a part of

your body as its “system” while the rest of your body and universe would be the surroundings.

Types of Systems.

• There are three different systems:– Open: open to the surroundings and able to

trade both matter and energy.– Closed: trades energy but not matter. – Isolated: trades absolutely nothing.

• In your notebooks, give me one example of each type. Closed and open systems are easy. For isolated, think of ways you keep food warm or hot.

Energy flow

• Energy, like most things, likes to spread out to its surroundings. This is why hot objects cool off and vice versa.

• Keep in mind, energy is not destroyed nor is it created- just moved around.

(E = q + w)

• Energy flow comes in two types: endothermic and exothermic.

Endo vs. Exo

• Endothermic processes (systems) absorb heat from the surroundings. Think cold.– Reactions that build molecules, such as

combination and synthesis.– Ice absorbing heat in order to melt.

Endo vs. Exo

• Exothermic reactions/systems release heat into the surroundings. Think hot.– Destruction of chemicals by combustion and

decomposition are good examples.– Fire, explosions, digesting food, etc.

• PRACTICE: 5.15

Measuring heat and energy.

• The two units most often used to measure heat and energy are the joule (J) and calorie (c).

• Calories in food are actually designated with a C, not c, and stand for kilocalories, 1000 calories each.

Measuring heat

• A joule is the amount of energy needed to lift 1 kg (about 2.2 lbs) one meter off the floor against the force of gravity.

– We use this to describe work done by chemical reactions.

• A calorie is the amount of energy needed to heat up one gram of water exactly 1˚C.

– This unit is usually used to describe heating and cooling processes.

Heat

• Heat (q) is the energy that is transferred from one object to another.

• (+) heat means energy goes into the system. (ENDOthermic)

• (-) heat means energy goes out of the system. (EXOthermic)

Heats of Reaction

• Enthalpy (ΔH) is a measure of the energy absorbed or released by a chemical reaction. Same as q, but for chemical rxns only.

• Enthalpy is calculated under constant pressure, so work is 0. (in other words, q = ΔH).

• Found by subtracting the energy of the reactants from that of the products.

• PRACTICE: 5.27

Specific Heat and Heat Capacity

• Heat capacity (C) is the energy needed to raise the temperature of an ENTIRE OBJECT by 1˚C.

• Specific heat (c) is the energy needed to raise the temperature of 1 GRAM of an object by 1˚C.

• Objects with a LOW heat capacity heat up more quickly than objects with HIGH heat capacity.

• Metals, being better conductors, heat up more quickly than water, styrofoam, and wood.

• Thus, heat capacity and specific heat are a reflection of how fast something heats up or cools off.

• High c/C: slow to heat/slow to cool (water)

• Low c/C: fast to heat/fast to cool (metals)

Why the difference?

• Temperature is NOT energy, but is the result of energy causing atoms to vibrate.

• All the different atoms and compounds react when heated, but in varying amounts.

• More vibration = more TEMPERATURE, but NOT necessarily more energy.

– Glass and water molecules vs. metal atoms. Note that glasses and metals form connections between each other.

Calorimetry Methods

• Constant Pressure (coffee-cup calorimetry)– Solution rxns and others where expansion is

not an issue.– q = mcΔT

• Constant Volume (bomb calorimetry)– Used when the system expands non-

negligibly.

Practical use in the lab

• The study of heat changes in chemical and physical processes is called calorimetry. (literally, measurement of calories.)

• The formula is:• q = mcΔT

q = heat, or enthalpy

m = mass

c = specific heat

ΔT = change in temperature

Example

• How much heat would it take to boil a full pot of water (about 4000 g) if you start at room temperature (25˚C)?

Solution

• q = mcΔTq = ?

m = 4000 g

c = for water, 4.184 J/g ˚C

ΔT = Water boils at 100˚C so, 100-25 = 75˚C

• Put it all together: q = (4000)(4.184)(75)– q = 1,255,200 J (alot of energy)

Practice

• 5.39

• 5.44

Conductors and Insulators

• Conductors HELP energy flow while insulators SLOW energy down.

• Heat travels through many objects through the vibrations of atoms. The more tightly packed together atoms are (more dense), the more easily energy can flow.

What it all means

• The more dense something is, the easier it can conduct energy. Metals, for example, do this very well.

• Insulators use less dense materials such as air to keep you warm, such as a coat (keep your heat from escaping.)

Question 1

• A gas is a good _______, and has a _____ heat capacity.– A. Conductor, low– B. Conductor, high– C. Insulator, low– D. Insulator, high

Question 2

• Given a block of metal and a pan of water weighing the same amount.– Which would heat up most quickly and which

would cool off the most slowly?• A. Water/Water• B. Metal/Metal• C. Metal/Water• D. Water/Metal

Question 3

• What best describes this object?– A. High heat capacity, insulator– B. Low heat capacity, conductor– C. High heat capacity, conductor– D. Low heat capacity, insulator

Heat flow in chemical reactions

• Heat flow is given by “q”, but for chemical reactions we refer to the term enthalpy, ΔH (pronounced as “delta H.”

• A closely related value, known as entropy, is given by ΔS and stands for the amount of energy lost to disorder and chaos.

Phase changes

• All substances must gain or lose energy in order to undergo phase changes.– Gas Liquid Solid (lose energy, exo)– Solid Liquid Gas (gain energy, endo)– Dissolving solids into liquids (usually endo)

Heats of Fusion/Solidification

• The heat of fusion (ΔHfus) is the energy gained by a solid to go to a liquid, energy lost to return to a solid is known as heat of solidification (ΔHsolid) .

Heats of Vaporization/Condensation

• The heat of vaporization (ΔHvap) is the energy gained to change from a liquid to a gas and the reverse is the heat of condensation (ΔHcond) .

Dissolving

• Any liquid containing a dissolved substance is known as a homogeneous mixture, or solution.

• The energy needed to dissolve something is called the heat of solution, (ΔHsoln) and is always endothermic, +.

Phase Changes and Temperature- Water

Phase changes

• Every substance has its own unique heats for each type of phase change. The higher the heat capacity, the higher these values will be.

• Phase changes are isothermal, meaning that the temperature stays the same until the change is complete.

Question 4

• The study of heat flows in physical and chemical processes is known as…?– A. Chemistry– B. Biothermistry– C. Thermometry– D. Calorimetry

Question 5

• Can liquid water exist at 0˚C?• A. Yes• B. No• C. Only if you add chemicals to it• D. Not enough information

Question 6

• Which enthalpy would be used here? – A. ΔHfus

– B. ΔHcond

– C. ΔHsoln

– D. ΔHsolid

(Rain)

Homework!

• Read 5.1-5.3 (Pages 145-154)

• Key Terms and a brief summary (3-4 sentences) for each section.

• 5.8, 5.16, 5.19, 5.20, 5.23

• Due to be handed in Tomorrow!

Hess’s Law

• Hess’s Law states that reactions that occur in a series of steps have a total enthalpy equal to the sum of the individual steps.

• Practice- 5.53

Heats of Formation

• Also known as enthalpies of formation (ΔHf): measure of the energy change when a substance is formed from raw elements.

• Values are noted as standard ( ex. ΔH˚, pronounced “delta H not”) when elements are in their standard states.