Chapter 10 - PowerPoint PPT Presentation

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Chapter 10. Thermodynamics. section 1. What is Energy?. energy (def)- the capacity of an object to do work or produce heat Ch. 10 is concerned with 2 types of energy: kinetic and potential. potential. kinetic. kinetic energy (def)- energy due to the motion of an object.

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Chapter 10

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Chapter 10

Thermodynamics

section 1

What is Energy?

• energy (def)-the capacity of an object to do work or produce heat

• Ch. 10 is concerned with 2 types of energy: kinetic and potential

potential

kinetic

kinetic energy (def)-energy due to the motion of an object

• is a measure of how fast something is moving

• depends on the mass of an object and its velocity

• can be calculated by:

KE = ½ mv2

KE is kinetic energy

m is mass

v is velocity

potential energy (def)- energy due to the position or composition of an object

• is a measure of stored energy

• “positional” example: water held behind a dam

potential energy (cont.)

• “compositional” example: the energy stored in unburned gasoline

• here the potential energy is stored as chemical energy within the atoms

• during combustion this energy is released as bonds are broken and new bonds are made

gasoline + oxygen  carbon dioxide + water + energy

(gravitational) potential energy can be converted to kinetic energy

example: a falling ball---initial position:

• at first, the ball has high potential energy (due to gravity) because its height is high

• it has low kinetic energy because it is not moving

example: a falling ball---as it falls:

• the potential energy decreases as its height (distance from the ground) decreases

• the kinetic energy increases as the velocity picks up

example: before-during-after:

• at any time during the fall the total energy remains the same

Etotal = EKE + Epot

• (as the potential energy decreases the kinetic energy of the falling ball increases)

All Energy is Measured in Joules

• energy can take many forms; all can be measured in joules (J)

• energy transfers are also measured in joules

• the amount of energy transferred from one sample must be equal to the energy received by a second sample

James Joule

1818-1899

The amount of energy transferred from one sample must be equal to the energy received by a second sample

• the total combined energy of the 2 samples remains constant

• energy is not created or destroyed in a transfer

Law of Conservation of Energy (1st Law of Thermodynamics)

law of conservation of energy(def)- energy can’t be created or destroyed, but it can be converted from one form to another

the total amount of energy in the universe is constant

Temperature and Heat are not the Same Thing

• temperature (def)-a measure of the average kinetic energy of an object

• the temperature of a sample depends on the average kinetic energy of its particles

• the higher the temperature, the faster the particles are moving

temperature is an intensive property

• intensive property(def)-a property that does not depend on the amount (# of particles) of a sample

• example: a small cup of 96ºC water is at the same temperature as a huge Cauldron of 96ºC of water (the particles in each have the same average speed)

Heat

• heat (def)-the energy transferred between objects that are at different temperatures

• heat is always transferred from the warmer object to the cooler object

• the study of heat is called thermodynamics

• heat is an extensive property

Heat is an extensive property

• extensive property(def)-a property that depends on the amount of the sample

• the amount of energy that can be transferred as heat depends on the amount of the sample

• example: there is more heat in a cauldron of 96ºC water than in one cup of 96ºC water

• if two samples of the same substance are at the same temperature, the larger sample can transfer more heat because it has more particles

Describing Energy Changes

• to study the energy changes during a process you must first determine the system and surroundings

• the “system” is what you are studying, it is the part of the universe you are focusing attention on

• the “surroundings” are everything else; everything outside the system

Describing Energy Changes (cont.)

• the change in energy (E) for a process is equal to the difference between the energy of the system after the process and before the process

E = Efinal - Einitial

Exothermic vs. Endothermic

exothermic process

• in an exothermic process heat “exits” the system and is released into the surroundings

• an exothermic process or reaction will “feel” hot

exothermic process (cont)

• the change in energy for an exothermic process is negative

• the final energy of the system is less than the initial energy of the system because heat is released

E is negative; Efinal0

E = Efinal - Einitial

(where EfinalEinitial )

Exothermic vs. Endothermic (cont.)

endothermic process

• in an endothermic process or reaction heat “enters” the system as it absorbs heat from the surroundings

• an endothermic process will “feel” cold

endothermic process (cont)

• the final energy of the system is more than the initial energy of the system because heat is gained

E is positive; Efinal0

E = Efinal - Einitial

(where EfinalEinitial )

Calorie vs. Joule

• energy can be measured in calories (cal) or joules (J)

• calorie (def)-the amount of energy (heat) required to raise the temperature of one gram of water by one degree Celsius

conversions:

1 calorie = 4.184 joules

1 cal = 4.184 J

calorie (cal) vs. Calorie (Cal)

food energy is often measured in Calories (Cal)

1 Calorie (food) = 1 kilocalorie

1 Cal = 1000 cal

Specific Heat

• different substances respond differently to being heated

• specific heat(def)-the amount of energy needed to raise the temperature of one gram of a substance by 1ºC

• the unit for specific heat is J/gºC (joule per gram degree Celsius)

• the symbol for specific heat is cp

Specific Heat (cont.)

• a high specific heat means a substance can absorb a lot of energy before its temperature is raised

• water has a high specific heat

• metals have low specific heats

Heat Calculations

• heat changes are calculated differently depending on whether the substance remains “in state” during the change

• when the substance remains IN STATE, only a temperaturechange occurs ( the substance does not change state)

…when a substance remains “in state” Q (heat) is calculated using this formula:

Q = m Tcp

(Q means heat, m means mass, cp means specific heat and T means the change in temperature) ….or…

Q = mass x T x specific heat

(T = Tfinal – Tinitial)

• if Q is positive the process is endothermic (heat was absorbed)

• if Q is negative the process is exothermic (heat was released)

…during a phase change (change of state) a different formula is needed:

Q = mass x heat of fusion (or heat of vaporization)

Molar Heat Capacity

• molar heat capacity(def)-the energy (heat) needed to increase the temperature of 1 mole of a pure substance by 1 K

• the symbol “C” stands for molar heat capacity

• the unit for molar heat capacity is J/mol K (joule per mol Kelvin)

molar heat capacity can be calculated using the following:

q = n C T

heat = (amt. in mol)(mol. heat cap.)(change in temp.)

• the equation shows the amount of heat (q) needed to increase the temperature of n moles of a substance by T

Read p. 343-344. List at least 2 important facts for each heading.

K. Molar Heat Capacity Depends on the Number of Atoms

L. Molar Heat Capacity is Related to Specific Heat

M. Heat Results in Disorderly Particle Motion