Transcript Introductory Chemistry: A Foundation Introductory Chemistry Basic
Introductory Chemistry: A Foundation, 6
th
Ed. Introductory Chemistry, 6
th
Ed. Basic Chemistry, 6
th
Ed.
by Steven S. Zumdahl & Donald J. DeCoste
University of Illinois
Chapter 10
Energy
Energy and Energy Changes
•
Energy
: ability to do work or produce heat – Chemical, mechanical, thermal, electrical, radiant, sound, nuclear – Potential and kinetic • Energy may affect matter.
– e.g. Raise its temperature, eventually causing a state change, or cause a chemical change such as decomposition •
All physical changes and chemical changes involve energy changes.
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10 | 3
Energy and Energy Changes
•
Potential Energy
: energy due to composition or position •
Kinetic Energy
: energy due to motion – - ½ mv 2 Copyright © Houghton Mifflin Company. All rights reserved.
10 | 4
Energy and Energy Changes
(cont.) •
Law of Conservation of Energy
: energy can be
converted
from one form to another, but cannot be created or destroyed Copyright © Houghton Mifflin Company. All rights reserved.
10 | 5
Work and Energy
•
Work
: force acting over a distance – w = f • d – Work done on a system will increase the energy of the system, whereas work done by the system will decrease the energy of the system •
State function
: a property that changes independent of pathway Copyright © Houghton Mifflin Company. All rights reserved.
10 | 6
Temperature and Heat
•
Heat
: a flow of energy due to a temperature difference •
Temperature
: a measure of the random motions of the components of a substance Copyright © Houghton Mifflin Company. All rights reserved.
10 | 7
Temperature and Heat
(cont.) Copyright © Houghton Mifflin Company. All rights reserved.
10 | 8
Exothermic vs. Endothermic
• • • •
System
: that part of the universe that we wish to study
Surroundings
: everything else in the universe
Exothermic process:
is any process that gives off heat – transfers thermal energy from the system to the surroundings. Example: when a match is struck, it is an exothermic process because energy is produced as heat. and 2H 2 O (
l
) + energy 2H 2 (
g
) + O 2 (
g
)
Endothermic process:
absorbs heat Example: melting ice to form liquid water is an endothermic process because the ice absorbs heat in order to melt energy + H 2 O (
s
) H 2 O (
l
) Copyright © Houghton Mifflin Company. All rights reserved.
10 | 9
Exothermic Process
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10 | 10
Enthalpy (H)
is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure.
D
H
=
H
(products) –
H
(reactants) D
H
= heat given off or absorbed during a reaction
at constant pressure
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10 | 11
Thermochemical Equations Is D
H
negative or positive?
System absorbs heat Endothermic D
H
> 0 6.01 kJ are absorbed for every 1 mole of ice that melts at 0 0 C and 1 atm.
H 2 O (
s
) H Copyright © Houghton Mifflin Company. All rights reserved.
2 O (
l
) D
H
= 6.01 kJ
Thermochemical Equations • • • The stoichiometric coefficients always refer to the number of moles of a substance H 2 O (
s
) H 2 O (
l
) D
H
= 6.01 kJ If you reverse a reaction, the sign of D
H
changes H 2 O (
l
) H 2 O (
s
) D
H
= 6.01
kJ If you multiply both sides of the equation by a factor
n
, then D
H
must change by the same factor
n
.
2H 2 O (
s
) 2H 2 O (
l
) D
H
= 2 x 6.01 = 12.0 kJ Copyright © Houghton Mifflin Company. All rights reserved.
10 | 13
Thermochemical Equations • The physical states of all reactants and products must be specified in thermochemical equations.
H 2 O (
s
) H 2 O (
l
) D
H
= 6.01 kJ H 2 O (
l
) H 2 O (
g
) D
H
= 44.0 kJ How much heat is evolved when 266 g of white phosphorus (P 4 ) burn in air?
P 4 (
s
) + 5O 2 (
g
) P 4 O 10 (
s
) D
H
= -3013 kJ 266 g P 4 x 1 mol P 4 123.9 g P 4 x 3013 kJ 1 mol P 4 = 6470 kJ Copyright © Houghton Mifflin Company. All rights reserved.
10 | 14
Different enthalpies
• Heat of reaction ( D H r or D H rxn )- heat energy absorbed or released during a reaction.
• Heat of formation ( D H f )- heat energy absorbed or released during synthesis of one mole of a compound from its elements at 298 K and 1 atm pressure (STP standard temp and pressure).
• Heat of solution ( D H sol )- heat energy absorbed or released when a substance dissolves in a solvent.
• Heat of combustion ( D H comb )- heat energy released when a substance reacts with oxygen to form CO 2 and H 2 O.
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10 | 15
• Heat of fusion ( D H fus )= Energy needed to melt one mole (solid to liquid) • Heat of vaporization ( D H vap )=Energy needed to boil one mole (liquid to gas) • In a phase change graph, it is possible to calculate the total energy involved as well as the energy consumed in each step.
• Note that water has different values for sp. heat depending upon its physical state.
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10 | 16
Thermodynamics
• The Law of Conservation of Energy is also known as The First Law of Thermodynamics. It can be stated as “the energy of the universe is constant.” • Internal Energy (E) = kinetic energy + potential energy • ΔE = q + w =
change
in internal energy q = heat absorbed by the system w = work done on the system Copyright © Houghton Mifflin Company. All rights reserved.
10 | 17
Units of Energy
•
One calorie
= amount of energy needed to raise the temperature of one gram of water by 1°C – kcal = energy needed to raise the temperature of 1000 g of water 1°C •
joule
– 4.184 J = 1 cal • In nutrition, calories are capitalized.
– 1 Cal = 1 kcal Copyright © Houghton Mifflin Company. All rights reserved.
10 | 18
Example - Converting Calories to Joules
Convert 60.1 cal to joules.
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10 | 19
Energy & Temperature of Matter
• The amount the temperature of an object increases depends on the amount of heat added (q).
– If you double the added heat energy the temperature will increase twice as much.
• The amount the temperature of an object increases when heat is added depends on its mass – If you double the mass it will take twice as much heat energy to raise the temperature the same amount.
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10 | 20
Specific Heat Capacity
•
Specific heat
(s): the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius
Amount of Heat = Specific Heat x Mass x Temperature Change Q = s x m x
D
T
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10 | 21
Specific Heat Capacity
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10 | 22
Example #1:
Calculate the amount of heat energy (in joules) needed to raise the temperature of 7.40 g of water from 29.0°C to 46.0°C.
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10 | 23
Example #1
(cont.) Specific heat of water = 4.184 g J C Mass = 7.40 g Temperature change = 46.0°C – 29.0°C = 17.0°C
Q = s • m •
D
T
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10 | 24
Example #2
A 1.6 g sample of metal that appears to be gold requires 5.8 J to raise the temperature from 23°C to 41°C. Is the metal pure gold?
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10 | 25
Example #2
Table 10.1 lists the specific heat of gold as 0.13
Therefore the metal cannot be pure gold.
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10 | 26
Enthalpy
•
Change in enthalpy
(ΔH p = q p ): the amount of heat exchanged when heat exchange occurs under conditions of constant pressure • Enthalpy is a
state function
• ΔH is independent of the path taken Copyright © Houghton Mifflin Company. All rights reserved.
10 | 27
Hess’s Law
•
Hess’s Law
: in going from a set of reactants to a set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.
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10 | 28
Hess’ Law
(cont.) • ΔH reaction = ∑Δh steps • If the direction of a reaction is reversed, the sign of ΔH is reversed.
• ΔH forward = -Δh reverse • Magnitude of ΔH α quantities of reactants and products Copyright © Houghton Mifflin Company. All rights reserved.
10 | 29
Hess’s Law
(cont.) • Overall reaction: N 2 + 2O 2 2NO 2 ΔH = 68 kJ • This reaction can be carried out in 2 steps: N 2 + O 2 2NO + O 2 2NO ΔH = 180 kJ 2NO 2 ΔH = -112 kJ ------------------------------------------------------- N 2 + 2O 2 2NO 2 ΔH = 68 Kj Note: the sum of the two reactions gives the overall reaction and the same is true for the sum of the enthalpy change values.
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10 | 30
Calorimetry
• The amount of heat flow transferred during a reaction is determined from temperature measurements made in a calorimeter.
• Heat loss is minimized by having insulation. A simple calorimeter can be made in the lab by stacking 2 styrofoam cups. A calorimeter minimizes heat exchange between the system and the surroundings.
• • Amount of heat produced is calculated by measuring the temp change in the surrounding water.
D H= D t (H20) X m H2O XCp H2O Copyright © Houghton Mifflin Company. All rights reserved.
10 | 31
Calorimetry
(cont.) Copyright © Houghton Mifflin Company. All rights reserved.
10 | 32
Energy Quality & Quantity
• While the total amount or quantity of energy in the universe is constant (1st Law) the quality of energy is degraded as it is used.
Burning of petroleum: High grade concentrated energy Low grade energy (heat) Copyright © Houghton Mifflin Company. All rights reserved.
10 | 33
Fuels • Petroleum
– A fossil fuel composed mainly of hydrocarbons
• Natural gas
– Consists largely of methane – Also contains ethane, propane, and butane Copyright © Houghton Mifflin Company. All rights reserved.
10 | 34
Fuels
(cont.) Copyright © Houghton Mifflin Company. All rights reserved.
10 | 35
Fuels
(cont.) Copyright © Houghton Mifflin Company. All rights reserved.
10 | 36
Fuels
(cont.) • Coal – Matures geologically through stages Copyright © Houghton Mifflin Company. All rights reserved.
10 | 37
Global Warming
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10 | 38
Global Warming
(cont.) Copyright © Houghton Mifflin Company. All rights reserved.
10 | 39
Energy Use and Sources
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10 | 40
Energy as a Driving Force
• Most processes that occur spontaneously involve an “energy spread.” – Heat flows from high to low temperature and “spreads” …or a “matter spread” – Salt dissolves or “spreads” in water Copyright © Houghton Mifflin Company. All rights reserved.
10 | 41
Entropy
• Entropy (S) is a measure of disorder or randomness.
– As a system becomes more disordered, ΔS >0 •
Second Law of Thermodynamics
: the entropy of the universe is always increasing.
• Tendency in nature is to increase disorder (unless external forces counteract). Ex—messy room, throwing a puzzle • -Entropy in solids< liquids 10 | 42 • Entropy=S (absolute Entropy for a substance=Entropy • at absolute zero temp -273K). Units J/K.mol • - change in Entropy= D S D S=S products -S reactants – - Entropy can be the driving force behind reactions. Ex- reactions leading to the formation of gases (from solids) are favored. – - Higher temp=higher Entropy (due to more KE of particles) – - Lower temp=lower Entropy (less KE) – - Ideal conditions for a spontaneous reaction • Increase Entropy (disorder) • Decrease enthalpy of products Copyright © Houghton Mifflin Company. All rights reserved. 10 | 43 • Reactions can thus occur even if not entropically favored. Ex • 2H 2 (g)+O 2 (g) 2H 2 O (g) • 3 molecules 2 molecules (decrease in entropy same physical state) • But reaction is highly exothermic (ie. Much lower enthalpy of products) Copyright © Houghton Mifflin Company. All rights reserved. 10 | 44