ENERGY

Energy

¨  Energy (E) is the ability to do work. ¨  Many types, but we can say 3 main

types: n  Radiant n  Potential n  Kinetic

US Energy Consumption by Source

Radiant Energy ¨ Light Energy

¤ Visible and Invisible ¨ Travels in waves over

distances ¤ Electromagnetic waves

n Waves that spread out in all directions from the source

n Visible light, UV light, Infra Red Radiation, X-rays, microwaves, radio waves

Potential Energy (PE) ¨ Stored Energy

¤ Due to position n Gravitational PE n Elastic PE

¤ Chemical bonds n Chemical PE

n Nuclear energy n Fuels n Attractions between

molecules

Kinetic Energy (KE)

¨ Energy of motion ¤ Atomic vibrations ¤ Molecular

movement n Vibration n Rotation n Translation

¤ Movement of subatomic particles

Kinetic Energy

Can be calculated:

How are each type shown here?

How are each type shown here?

¨  Radiant ¤ Rainbow = visible light

¨  Kinetic ¤ Windmill moving

¨  Potential ¤ All molecules store

energy n Water in clouds n Air n Materials the windmill is

made from, the plants at the bottom

Where is the PE? KE?

Temperature Scales: Measuring that Thermal Energy

Boiling Freezing

Fahrenheit (oF) 212 32

Celsius (oC) 100 0

Kelvin 373 273

A Note on the Fahrenheit Scale

¨  NEVER use it in this class. Ever. Only Belize and the US use this scale.

¨  Gabriel Fahrenheit made great thermometers. His scale was replicated the world over because of this. But if you stop and think about it, does 32°F for freezing make sense, or 212°F for boiling? 180 degrees separates them.

¨  100 degrees, as in the Celcius scale (sometimes called the Centigrade scale) makes much more sense.

¨  Fahrenheit based 0°F on the freezing point of water mixed with NH4Cl, and 32°F for freezing water, and 96°F for human body temperature (he was off by 2.6°). Why? Because he felt like it and it was easy to draw lines at those intervals.

¨  (According to a letter Fahrenheit wrote to his friend Herman Boerhaave,[8] his scale was built on the work of Ole Rømer, whom he had met earlier. In Rømer’s scale, brine freezes at 0 degrees, ice melts at 7.5 degrees, body temperature is 22.5, and water boils at 60 degrees. Fahrenheit multiplied each value by four in order to eliminate fractions and increase the granularity of the scale. He then re-calibrated his scale using the melting point of ice and normal human body temperature (which were at 30 and 90 degrees); he adjusted the scale so that the melting point of ice would be 32 degrees and body temperature 96 degrees, so that 64 intervals would separate the two, allowing him to mark degree lines on his instruments by simply bisecting the interval six times (since 64 is 2 to the sixth power). I took this from Wikipedia.

Kelvin Temperatures

¨ Based on absolute zero (0 K, -273 oC) ¤ The temperature at which ALL KE stops ¤ NO molecular motion. ¤ Lowest temperature theoretically possible

n Can’t really get there in real life n (See 3rd Law of Thermodynamics in a few

slides)

¨ K = oC + 273 ¤ Technically 273.14, but we can stop at 3

significant digits

Why do we need the Kelvin scale?

¨  Two reasons ¤ We need a scale that is relative to molecular motion

for certain topics n You can’t use negative numbers to indicate motion when it

IS present n -20°C makes NO sense in light of indicating motion

n And 40°C ISN’T twice as much motion as 20°C, n  (40K IS twice the motion of 20K)

¤  Because when working with equations, can’t use zero n  We get undefined answers if we divide n  We get answers of 0 if we multiple n  And those answers would NOT make sense if compared to answers

calculated with a positive or negative number

¨  Energy (E): The ability to do work

¨  Exergy: The energy available to do work

¨  No symbol

¨  Entropy (S): The measure of the disorder of a system

¨  Enthalpy(H): The thermal energy (heat) content of a system

The 4 Es: Energy, Exergy, Entropy, & Enthalpy

There will be more on these!

Thermodynamics

¨  The study of energy flow ¤  inter-relation between heat, work, and energy of a

system

¨  Summary of the three laws: 1.  The energy in the universe is constant 2.  Things get more disorganized over time in a system

until everything is equal 3.  You can’t reach absolute zero

1st Law of Thermodynamics

¨  The energy in the universe is constant ¤ E=mc2

¤ Law of Conservation Matter n  Matter can not be created or destroyed

¤ Law of Conservation of Energy n Energy can not be created or destroyed

n However, matter and energy can both change forms in chemical reactions

n Can also interconvert between matter and energy in NUCLEAR reactions (more on this later this year.)

Summed up: You can not win. You can’t get something for nothing because

energy and matter are conserved.

Ener

gy

Time

Before the 2nd Law…

¨ Entropy (S) is a measure of DISORGANIZATION in a system (this simply put; there is a much more complicated description about the unavailable energy to do work) ¤ Anything disorganized has higher entropy than

something organized ¨ Exergy is the Energy available to do work

2nd Law of Thermodynamics

¨ Things get more disorganized over time in a system until everything equilibrium is reached (everything is equal) ¤ Heat flows from hot to cold, not the reverse ¤ Law of Entropy

n By nature, things get more disorganized to spread out energy and matter

¨ The quality of the energy (which is exergy) decreases over time

Summed up: You can not break even. You can not return to the same energy state because things get more disorganized (gain entropy)

Exergy and Energy ¨  The energy of the universe is constant, but exergy is constantly

consumed. This can be compared with a tooth-paste tube: When you squeeze the tube (= conduct any process) the paste (= exergy) comes out. You can never put the paste back in the tube again (try!), and in the end you have only the tube itself (= low-exergy) left.

¨  When you squeeze the tube, the depressions (= entropy) will increase. (The entropy of a system increases when exergy is lost) But you can never take the depressions in the tube and 'un-brush' your teeth. (I.e. entropy is not negative exergy.)

¨  When you buy energy from the electricity network, you actually buy exergy. You can find the energy as room temperature heat after some time, but you can not take that room temeperature energy back to the electricity company and ask for money back. They won't accept it.

Energy and Matter Gain Entropy Over Time

Exergy: The Energy available to do work

3rd Law of Thermodynamics

¨ You can’t reach absolute zero and expect things to happen ¤ At absolute zero, all kinetic motion

ceases. And that energy needs to go somewhere. It goes to something else. And gets transferred back until everything is at an equal temperature.

Summed up: You can not get out of the game, because absolute zero is unobtainable.

Law of Conservation of Energy

¨ Energy cannot be created or destroyed…but it CAN change forms. ¤ Example: Burning wood in a fire

¨ The energy in chemical bonds is released as heat (KE and PE), light (RE), sound (KE) ¤ These forms of energy are less useful

n have less exergy

Radiant Energy:

EM Waves

Potential Energy:

Stored

Kinetic Energy:

Motion

The CPE in these items could:

Rio Summer Olympics Proposed Solar Waterfall

http://www.snopes.com/photos/architecture/solartower.asp

Combinations of PE and KE are very common on a large scale

KE and PE animation

PE and KE

When E changes forms…

¨ The amount of energy one thing loses is gained somewhere else. ¤ E lost = E gained (Law of Conservation

of Energy) ¤ But the E gained is usually not all in one

place (2nd Law of thermodynamics) n It is spread out (more entropy)

n Often in the forms of heat and light n Which are less useful (less exergy)

Energy Transformations

•  What’s up with Temperature vs Heat?

•  Temperature is related to the average kinetic energy of the particles in a substance.

Thermal Energy: KE + PE on the small scale

àAs temperature increases, so does thermal energy (because the energy of the particles increased).

àIf the temperature stays the same, the thermal energy in a more massive substance is higher (because it is a total measure of energy).

Thermal energy relationships

Heat

The flow of thermal energy from one object to another.

Heat always flows from warmer to cooler objects.

Ice gets warmer while hand gets

cooler

Cup gets cooler while hand gets warmer

Heat and Temperature

¨ Heat: the measure of the flow of RANDOM kinetic energy

¨ Temperature: the measure of heat ¨ So…temperature is a measure of

kinetic energy of the particles of a substance

* Sometimes heat is radiated as IR (infra-red radiation, a form of radiant energy)

• PE from how the molecules are placed relative to each other (attractions)

• Farther = more PE, just like how something farther off the ground has higher gravitational PE

Thermal Energy • Thermal Energy is the total of all the (kinetic and potential) heat energy of all the particles in a substance.

¨  Energy is being gained/ absorbed by the object or substance (called the system) from the surroundings

¨  Have positive change in enthalpy values (+ΔH)

¨  Energy is lost/ released from the object or substance (called the system) to the surroundings

¨  Have negative change in enthalpy values (-ΔH)

Endothermic Exothermic

Exothermic and Endothermic Processes

“Exothermic Reactions? I studied them before they were cool.”

The big picture…

¨ How do we see this energy cycling in the real world, and not just as a part of Chemistry class?

¤ Around the house? ¤ In the environment? ¤ While thinking about a car?

• If the cup is the system, it is undergoing an exothermic process because it is losing heat to the surroundings (hand)

Ice gets warmer while hand gets

cooler

Cup gets cooler while hand gets warmer

• If the ice is the system, it is undergoing an endothermic process because it is absorbing heat from the surroundings (hand)

Which process is endothermic? Which is exothermic?

3°Consumers: Carnivores and Omnivores

2°Consumers: Carnivores and Omnimores

1°Consumers: Herbivores

Producers: Autotrophs

Trophic Levels and Energy

Energy Out; 90% per level

Consumers are all heterotrophs

Producers:

Trophic Levels and Energy

3°Consumers

2°Consumers

1°Consumers

Trophic Levels and Energy

While this shows only a 1° consumer, all animals “lose” E the same way; high levels lose more to motion than do lower levels

Class Question:

¨ How does a car exhibit all types of energy? ¤  Explain!

http://www.consumerenergycenter.org/transportation/consumer_tips/vehicle_energy_losses.html

Energy Loss in a Car

Can the world really run out of Energy?

World-Wide Energy Sources, (2007)

PHASE CHANGES & ENERGY

Phase Diagrams

¨  Tell what state of matter a material is in at a given temperature and pressure

¨  The triple point is the pressure and temperature when a solid, liquid, and a gas of the same substance exist at equilibrium ¤  Equilibrium: When there is no net change

n  Here referring to changes in state n  Can also refer to temperature and chemicals

¨  The critical point is the temperature above which a substance will always be a gas, regardless of pressure

¨  Fullerton Phase Diagram Explorer Link

Phase Diagrams

Phase Diagram for Water

A few terms

¨  Freezing Point - The temperature at which the solid and liquid phases of a substance are in equilibrium at atmospheric pressure. ¤ The same temperature as the melting point

¨  Boiling Point - The temperature at which the vapor pressure of a liquid is equal to the pressure on the liquid.

¨  Vapor Pressure- The pressure at which the vaporization rates are equal to condensation rates

Measuring Vapor Pressure

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/vaporv3.swf

Energy and Phase Changes

Phase Changes

Enthalpy(H): The heat (thermal energy) content of a system

States of Matter and Entropy

The states are NOT plateaus because entropy is NOT constant. This isn’t a phase change diagram.

Energy and Matter and Connected

¨ Any change in matter ALWAYS is accompanied by a change in energy

Phase Changes and Energy

Heating Curve Te

mpe

ratu

re,

C

Time, min

States on a heating curve

Temperature,  C

Time,  minutes

SOLID

States  of  Matter  on  a  Heating  Curve

LIQUID

GAS

Changes of state on a heating curve

Temperature,  C

Time,  minutes

Two  states  of  matter  exist  on  the  plateausChanges  of  State  on  a  Heating  Curve

SOLID  becoming  LIQUID

LIQUID  becoming  VAPOR

Remember: Phase Changes and Energy

Why does temperature remains constant when melting or boiling?

¨  During melting or boiling, energy is absorbed from the surroundings ¤ Due to the increase in the thermal energy of the particles

from the increase in PE of the particles n Molecules are

n moving apart n breaking attractions which n  Absorbs latent (hidden) heat

n  can not be measured on a thermometer

¤ Substance (system) gets warmer

The E’s and Heating

• Endothermic process • Energy is absorbed from surroundings

• Entropy increases • Enthalpy is positive (+ΔH) since heat added • Exergy decreases

Why does temperature remains constant when freezing or condensing?

¨  During freezing or condensing, energy is released to the surroundings ¤ Due to the decrease in the thermal energy from the

decrease in PE of the particles n Molecules are

n moving closer n  forming new attractions that are n  Releasing latent (hidden) heat

n  can not be measured on a thermometer

¤ Substance (system) gets colder

The E’s and Cooling

• Exothermic process • Energy is lost to surroundings

• Entropy decreases • Enthalpy is negative (-ΔH) since heat is lost • Exergy increases

Processes of a Heating Curve

Temperature,  C

Time,  minutes

SOLID

LIQUID

GAS

FUSION  (MELTING)

VAPORIZATION

What happens during each segment

Cooling Curve: The Reverse of a Heating Curve

Temperature,  C

Time,  minutes

Cooling curve

Temperature,  C

Cooling  Curve:  The  Reverse  of  a  Heating  Curve

gas

liquid

solid

Gas  and  liquid  present

Liquid  and  solid  present

condensation

Freezing  

Tem

p,

C

Time, min

Energy and Phase Change in Nature:

Energy and Phase Change in Nature:

Measuring the Energy of Phase Changes

¨ The math of thermal energy flow

REMEMBER: Energy and Matter and Connected

¨ Any change in matter ALWAYS is accompanied by a change in energy

¨  This includes changes in temperature and/ or phase

•  Things heat up or cool down at different rates.

Land heats up and cools down faster than water, and aren’t we lucky for that!?

Specific Heat : c

• Specific heat is the amount of heat required to raise the temperature of 1 kg of a material by one degree °C

• cwater = 4.184 J / g °C • the number is high; water “holds” its heat

• c sand= 0.664 J / g °C • less E than water to change it; it doesn’t hold heat as well as water does

This is why land heats up quickly during the day and cools quickly at night and why water takes longer.

Why does water have such a high specific heat?

Water molecules form strong attractions with other water molecules; it takes more heat energy to break those attractions than other

materials with weaker forces of attraction between them.

water metal

Specific Heat Capacities of Selected Substances

¤ cwater = 4.184 J / g °C ¤ cice = 2.09 J / g °C ¤ csteam = 1.99 J / g °C ¤ csand = 0.664 J / g °C ¤ cAl = 0.90 J / g °C ¤ cFe = 0.449 J / g °C

Heat can be Transferred even if there is No Change in State

q = mc∆T

Remember this? Which is process is endothermic? Which is exothermic?

Now we care about how much energy is being transferred, and are ready to calculate that change.

Calculating Changes in Energy: The Calorimetry Equation

q = mcrT • q = change in thermal energy

• (+) value means heat is absorbed • (-) value means heat is released

• m = mass of substance • rT = change in temperature (Tfinal – Tinitial) • c = specific heat of substance

• Each substance has a different c (see CRH, p__) • Different states of matter for the same substance may have a different c

Specific Heat Capacity Problems If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost

by the Al?

• heat gained or lost = q = mc∆T • where ∆T = Tfinal - Tinitial • q = (25.0g) (0.897 J/g•oC)(37 - 310)oC • q = - 6120 J

Notice that the negative sign on q signals heat “lost by” or transferred OUT of Al.

Was this an endothermic or exothermic process?

Specific Heat Capacity Problems If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost

by the Al?

• heat gained or lost = q = mc∆T • where ∆T = Tfinal - Tinitial • q = (25.0g) (0.897 J/g•K)(37 - 310)K • q = - 6120 J

Notice that the negative sign on q signals heat “lost by” or transferred OUT of Al.

Was this an endothermic or exothermic process? EXOTHERMIC

Heat was released from the Al

Or… Heat Transfer can cause a

Change of State

Changes of state involve energy changes at constant T

Ice + 333 J/g (heat of fusion) -----> Liquid water Is there an equation? Of course!

Or… Heat Transfer can cause a Change of State

Changes of state involve energy at constant T H20(s) +333J/g à H20(l)

Ice + 333 J/g (heat of fusion) à Liquid water

q = mΔHfusion • m= mass • ΔHfusion = the enthalpy of melting

•  the change in thermal energy associated with melting • Units are J/g or KJ/Kg

q = mΔHfusion

¨  WHY DO I NEED THIS WHEN I HAVE

q = mc∆T?

¨  Well, when a phase changes THERE IS NO change in temperature… but there is definitely a change in energy!

Sample Problem:

How much heat energy is required to melt 25.0g of ice,

(assuming constant temperature of O°C)?

q = mΔHfusion • m= 25.0g • ΔHfusion =333J/g for water

q= (25.0g)333J/g =8385J

Value is positive, which means heat is absorbed, which makes sense!

molecule

Latent heat* and the PE of particles

Regular arrangement breaks up

strong attraction weak attraction

*Latent means hidden. Latent heat is the thermal energy (potential energy) associated with the attractions between molecules, and can not be measured with a thermometer.

PE related to the forces of attraction between the particles

Energy has to be supplied to oppose the attractive force of the particles.

PE ê as molecules separate

solid è liquid or liquid è gas

average potential energy ê

Latent heat and the PE or particles

The transfer of energy does not change the KE.

Temperature does not change.

latent heat = change in PE between molecules during change of state

Latent heat and PE

Video and song: http://www.youtube.com/watch?v=jaaGqui9NVY

Remember…… •  Energy changes accompany changes in state; either:

•  Energy is added (endothermic)•  Gain thermal energy

•  Molecules •  Move more (gain KE)•  Separate (gain PE from broken attractions between

molecules)•  Have a higher entropy

•  Are more disorganizedOr

•  Energy is removed (exothermic)•  Molecules move less

•  Lose thermal energy•  Move less (lose KE)•  Move closer (lose PE from new attractions between

molecules)•  Have lower entropy

•  Get more organized

Latent Heats

¨  You have a certain energy change associated with changing state. These values are usually reported for fusion and vaporization as: ¤ ΔHfusion= (latent) Heat of fusion (melting) ¤ Δ  Hvaporization = (latent) Heat of vaporization ¤ Δ  Hsublimation =(latent) Heat of sublimation

¨  Different materials have different values for each

What about freezing and condensation?

¨  Values for freezing and condensation are not typically listed, but are the negative values of those for fusion and vaporization because the energy transferred is the same, but in the opposite direction

¤  (latent) Heat of freezing= -ΔHfusion

¤  (latent) Heat of condensation= -Δ  Hvaporization

Enthalpy changes with phase changes

Enthalpy values for H2O

¨ ∆Hfusion= 334 J/g ¨ ∆Hvaporization= 2259 J/g ¨ ∆Hsublimation = 25940 J/g

From: http://hyperphysics.phy-astr.gsu.edu/hbase/tables/phase.html#c1

¨  http://hyperphysics.phy-astr.gsu.edu/hbase/tables/phase.html#c2

Summing it all up: How do you know what to do to calculate energy changes?

•  Check to see if there is a temperature change. •  If yes, use q=mcΔT.

•  Also, check to see if there is a phase change. •  If yes, you need to use

•  q=  Δ  Hfusionmass •  or

•  q=  Δ  Hvaporizationmass •  depending on which one applies* •  or both if there are two phase changes

*If the material freezes or condenses. You can use the negative value Δ  Hfusion or Δ  Hvaporization

How much energy is required to change 0.5 kg of water at 0 °C to ice?

Things you know: • m = 0.5 kg= 500.g • There is no temperature change, and there is change of state (freezing) • The water is going to freeze So…..

this all tells you to use –ΔHfusion (negative of melting value) in

q= –ΔHfusionm q= (-334J/g)(500.g)= -1.67E5J

(The negative value makes sense since you are cooling the water, so energy leaves)

Total energy required

= latent heat (ice at 0 °C → water at 0 °C) + energy (water: 0 °C → 80 °C)

= mΔHf + mc ΔT

= (500g)(334 J/ g) + (500.g x 4.184J/g°C x 80°C)

= (167000 J) + (167360J) =

=334360J= 3.34E5 J (Value is positive, makes sense since you add E to heat water)

How much energy is required to melt 0.5 kg of ice at 0 °C

temperature raised to 80 °C?

Heat & Changes of State

What quantity of heat is required to melt 500. g of ice and heat the water to steam at 100. oC?

Heat of fusion of ice = 334 J/gSpecific heat of water = 4.184 J/g•°C

Heat of vaporization = 2259 J/g

+334 J/g

+2257 J/g

So… if I want the total heat to take ice and turn it to steam I need to add values from 3 steps… 1.  To melt the ice I need to multiply the heat of fusion with the mass

•  q = ∆Hfusionm 2.  Then, there is moving the temperature from 0°C to 100°C.

•  For this there is a change in temperature so we use •  q= mc∆T

3.  That just takes us to 100°C, what about vaporizing the molecules?

•  We need q=∆Hvaporizationm Add up all the values, and you have it. (However, if you are taking it from below the freezing point to above 100°C, you

need to add in the changes with q=mc ∆ T there, too!)

Putting it all together…

And now… More! Heat & Changes of State

How much heat is required to melt 500. g of ice and heat the water to steam at 100 oC?

1. To melt ice : q = m∆Hfusion q = (500.g)(334 J/g) = 1.67 x 105 J

2. To raise water from 0 oC to 100 oC : q = mc∆T q = (500. g)(4.184 J/g• oC)(100 - 0oC) = 2.1 x 105 J 3. To evaporate water at 100 oC: q = m∆Hvaporization q =(500.g)(2259 J/g) = 1.13 x 106 J 4. Total heat energy = 1.51 x 106 J = 1510 kJ

Maybe a picture can help….

Putting it all together:

¨ How are matter and energy related? What influences does energy have on matter? What does this tell us about the world as we know it?

Making Pizza: Changing Matter Describe the pizza making process in terms of: ¨  Matter

¤ States (s, l, g) ¤ Elements, compounds, mixtures

n Homogeneous and heterogeneous mixtures

¨  Properties and changes ¤ Both

n  chemical and physical n  Intrinsic (intensive) and extrinsic (extensive)

¨  Energy