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Heat Transfer

10.4k plays, 6th -  10th  , magnetic field, semiconductors, specific heat and calorimeter.

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First Law of Thermodynamics

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20 questions

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  • 1. Multiple Choice 1 minute 1 pt Why is water often used as a coolant in automobiles, other than the fact that it is abundant? Water expands very little as it is heated. The freezing temperature of water has a relatively large value. The specific heat of water is relatively small and its temperature can be rapidly decreased. The specific heat of water is relatively large and it can store a great deal of thermal energy. Water does not easily change into a gas.
  • 2. Multiple Choice 1 minute 1 pt Heat is added to a substance, but its temperature does not increase. Which one of the following statements provides the best explanation for this observation? The substance has unusual thermal properties. The substance must be cooler than its environment. The substance must be a gas. The substance must be an imperfect solid. The substance undergoes a change of phase.
  • 3. Multiple Choice 1 minute 1 pt Complete the following statement: The first law of thermodynamics states that… the entropy of the universe is increasing. entropy is a function of the state of a system. heat is a form of energy. the change in the internal energy of a system is given by Q − W. two systems in thermal equilibrium with a third system are in equilibrium with each other.
  • 4. Multiple Choice 1 minute 1 pt When applying the first law of thermodynamics to a system, when is heat a positive quantity? when the system does work when the system has work done on it when the system absorbs heat when the system loses heat when no work is done either on the system or by the system
  • 5. Multiple Choice 1 minute 1 pt Which one of the following phrases correctly describes an adiabatic process? no loss of energy occurs no transfer of energy as heat no change in temperature occurs no change in system volume occurs no change in system pressure occurs
  • 6. Multiple Choice 1 minute 1 pt The product of the pressure and volume of a system  has the same SI units as which one of the following choices? force work acceleration momentum impulse

The first law of thermodynamics states that the change in the internal energy of a system is equal to the difference in energy transferred to or from the system as heat and

The first law of thermodynamics is a restatement of the

Zeroth law of thermodynamics

Law of heat addition

Principle of entropy

Conservation of energy

The internal energy of an ideal gas depends on

Volume and temperature

Pressure and temperature

Temperature only

The internal energy change in a system that has absorbed 2Kcal of heat and done 500J of work is

When a gas undergoes isothermal expansion, its internal energy

remains constant

The conditions for the thermodynamic process to be isothermal,

Temperature remains constant

Pressure remains constant

The process should be very slow and the walls of cylinder should be perfectly conducting

The process should be very fast and the walls of cylinder should be perfectly non-conducting

  • 13. Multiple Choice 30 seconds 1 pt What is the equation for the first law of thermodynamics? U=Q+W U=Q-W U=-Q+W U+Q+W=0
  • 14. Multiple Choice 30 seconds 1 pt The first law of thermodynamics is basically the same as which law from physics 1? Newton's first law The law of conservation of energy The law of conservation of momentum Newton's second law
  • 15. Multiple Choice 30 seconds 1 pt The SI unit for measuring heat is the  pascal newton Joule calorie
  • 16. Multiple Choice 30 seconds 1 pt Thermodynamics is the study of what? The transfer of energy The creation of energy The speed of reactions  The affect of heat on the speed of a reaction 
  • 17. Multiple Choice 30 seconds 1 pt The amount of heat needed to melt one mole of a substance is called: Specific Heat Latent Heat Heat of vaporization Heat of fusion
  • 19. Multiple Choice 30 seconds 1 pt Heat is measured in  joules grams degrees celcius
  • 20. Multiple Choice 30 seconds 1 pt The Law of _______________ states that energy cannot be created nor destroyed only transferred  Specific Heat Conservation of Energy Exothermic Reaction Potential Energy

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  • 15.1 The First Law of Thermodynamics
  • Introduction to Science and the Realm of Physics, Physical Quantities, and Units
  • 1.1 Physics: An Introduction
  • 1.2 Physical Quantities and Units
  • 1.3 Accuracy, Precision, and Significant Figures
  • 1.4 Approximation
  • Section Summary
  • Conceptual Questions
  • Problems & Exercises
  • Introduction to One-Dimensional Kinematics
  • 2.1 Displacement
  • 2.2 Vectors, Scalars, and Coordinate Systems
  • 2.3 Time, Velocity, and Speed
  • 2.4 Acceleration
  • 2.5 Motion Equations for Constant Acceleration in One Dimension
  • 2.6 Problem-Solving Basics for One-Dimensional Kinematics
  • 2.7 Falling Objects
  • 2.8 Graphical Analysis of One-Dimensional Motion
  • Introduction to Two-Dimensional Kinematics
  • 3.1 Kinematics in Two Dimensions: An Introduction
  • 3.2 Vector Addition and Subtraction: Graphical Methods
  • 3.3 Vector Addition and Subtraction: Analytical Methods
  • 3.4 Projectile Motion
  • 3.5 Addition of Velocities
  • Introduction to Dynamics: Newton’s Laws of Motion
  • 4.1 Development of Force Concept
  • 4.2 Newton’s First Law of Motion: Inertia
  • 4.3 Newton’s Second Law of Motion: Concept of a System
  • 4.4 Newton’s Third Law of Motion: Symmetry in Forces
  • 4.5 Normal, Tension, and Other Examples of Forces
  • 4.6 Problem-Solving Strategies
  • 4.7 Further Applications of Newton’s Laws of Motion
  • 4.8 Extended Topic: The Four Basic Forces—An Introduction
  • Introduction: Further Applications of Newton’s Laws
  • 5.1 Friction
  • 5.2 Drag Forces
  • 5.3 Elasticity: Stress and Strain
  • Introduction to Uniform Circular Motion and Gravitation
  • 6.1 Rotation Angle and Angular Velocity
  • 6.2 Centripetal Acceleration
  • 6.3 Centripetal Force
  • 6.4 Fictitious Forces and Non-inertial Frames: The Coriolis Force
  • 6.5 Newton’s Universal Law of Gravitation
  • 6.6 Satellites and Kepler’s Laws: An Argument for Simplicity
  • Introduction to Work, Energy, and Energy Resources
  • 7.1 Work: The Scientific Definition
  • 7.2 Kinetic Energy and the Work-Energy Theorem
  • 7.3 Gravitational Potential Energy
  • 7.4 Conservative Forces and Potential Energy
  • 7.5 Nonconservative Forces
  • 7.6 Conservation of Energy
  • 7.8 Work, Energy, and Power in Humans
  • 7.9 World Energy Use
  • Introduction to Linear Momentum and Collisions
  • 8.1 Linear Momentum and Force
  • 8.2 Impulse
  • 8.3 Conservation of Momentum
  • 8.4 Elastic Collisions in One Dimension
  • 8.5 Inelastic Collisions in One Dimension
  • 8.6 Collisions of Point Masses in Two Dimensions
  • 8.7 Introduction to Rocket Propulsion
  • Introduction to Statics and Torque
  • 9.1 The First Condition for Equilibrium
  • 9.2 The Second Condition for Equilibrium
  • 9.3 Stability
  • 9.4 Applications of Statics, Including Problem-Solving Strategies
  • 9.5 Simple Machines
  • 9.6 Forces and Torques in Muscles and Joints
  • Introduction to Rotational Motion and Angular Momentum
  • 10.1 Angular Acceleration
  • 10.2 Kinematics of Rotational Motion
  • 10.3 Dynamics of Rotational Motion: Rotational Inertia
  • 10.4 Rotational Kinetic Energy: Work and Energy Revisited
  • 10.5 Angular Momentum and Its Conservation
  • 10.6 Collisions of Extended Bodies in Two Dimensions
  • 10.7 Gyroscopic Effects: Vector Aspects of Angular Momentum
  • Introduction to Fluid Statics
  • 11.1 What Is a Fluid?
  • 11.2 Density
  • 11.3 Pressure
  • 11.4 Variation of Pressure with Depth in a Fluid
  • 11.5 Pascal’s Principle
  • 11.6 Gauge Pressure, Absolute Pressure, and Pressure Measurement
  • 11.7 Archimedes’ Principle
  • 11.8 Cohesion and Adhesion in Liquids: Surface Tension and Capillary Action
  • 11.9 Pressures in the Body
  • Introduction to Fluid Dynamics and Its Biological and Medical Applications
  • 12.1 Flow Rate and Its Relation to Velocity
  • 12.2 Bernoulli’s Equation
  • 12.3 The Most General Applications of Bernoulli’s Equation
  • 12.4 Viscosity and Laminar Flow; Poiseuille’s Law
  • 12.5 The Onset of Turbulence
  • 12.6 Motion of an Object in a Viscous Fluid
  • 12.7 Molecular Transport Phenomena: Diffusion, Osmosis, and Related Processes
  • Introduction to Temperature, Kinetic Theory, and the Gas Laws
  • 13.1 Temperature
  • 13.2 Thermal Expansion of Solids and Liquids
  • 13.3 The Ideal Gas Law
  • 13.4 Kinetic Theory: Atomic and Molecular Explanation of Pressure and Temperature
  • 13.5 Phase Changes
  • 13.6 Humidity, Evaporation, and Boiling
  • Introduction to Heat and Heat Transfer Methods
  • 14.2 Temperature Change and Heat Capacity
  • 14.3 Phase Change and Latent Heat
  • 14.4 Heat Transfer Methods
  • 14.5 Conduction
  • 14.6 Convection
  • 14.7 Radiation
  • Introduction to Thermodynamics
  • 15.2 The First Law of Thermodynamics and Some Simple Processes
  • 15.3 Introduction to the Second Law of Thermodynamics: Heat Engines and Their Efficiency
  • 15.4 Carnot’s Perfect Heat Engine: The Second Law of Thermodynamics Restated
  • 15.5 Applications of Thermodynamics: Heat Pumps and Refrigerators
  • 15.6 Entropy and the Second Law of Thermodynamics: Disorder and the Unavailability of Energy
  • 15.7 Statistical Interpretation of Entropy and the Second Law of Thermodynamics: The Underlying Explanation
  • Introduction to Oscillatory Motion and Waves
  • 16.1 Hooke’s Law: Stress and Strain Revisited
  • 16.2 Period and Frequency in Oscillations
  • 16.3 Simple Harmonic Motion: A Special Periodic Motion
  • 16.4 The Simple Pendulum
  • 16.5 Energy and the Simple Harmonic Oscillator
  • 16.6 Uniform Circular Motion and Simple Harmonic Motion
  • 16.7 Damped Harmonic Motion
  • 16.8 Forced Oscillations and Resonance
  • 16.10 Superposition and Interference
  • 16.11 Energy in Waves: Intensity
  • Introduction to the Physics of Hearing
  • 17.2 Speed of Sound, Frequency, and Wavelength
  • 17.3 Sound Intensity and Sound Level
  • 17.4 Doppler Effect and Sonic Booms
  • 17.5 Sound Interference and Resonance: Standing Waves in Air Columns
  • 17.6 Hearing
  • 17.7 Ultrasound
  • Introduction to Electric Charge and Electric Field
  • 18.1 Static Electricity and Charge: Conservation of Charge
  • 18.2 Conductors and Insulators
  • 18.3 Coulomb’s Law
  • 18.4 Electric Field: Concept of a Field Revisited
  • 18.5 Electric Field Lines: Multiple Charges
  • 18.6 Electric Forces in Biology
  • 18.7 Conductors and Electric Fields in Static Equilibrium
  • 18.8 Applications of Electrostatics
  • Introduction to Electric Potential and Electric Energy
  • 19.1 Electric Potential Energy: Potential Difference
  • 19.2 Electric Potential in a Uniform Electric Field
  • 19.3 Electrical Potential Due to a Point Charge
  • 19.4 Equipotential Lines
  • 19.5 Capacitors and Dielectrics
  • 19.6 Capacitors in Series and Parallel
  • 19.7 Energy Stored in Capacitors
  • Introduction to Electric Current, Resistance, and Ohm's Law
  • 20.1 Current
  • 20.2 Ohm’s Law: Resistance and Simple Circuits
  • 20.3 Resistance and Resistivity
  • 20.4 Electric Power and Energy
  • 20.5 Alternating Current versus Direct Current
  • 20.6 Electric Hazards and the Human Body
  • 20.7 Nerve Conduction–Electrocardiograms
  • Introduction to Circuits and DC Instruments
  • 21.1 Resistors in Series and Parallel
  • 21.2 Electromotive Force: Terminal Voltage
  • 21.3 Kirchhoff’s Rules
  • 21.4 DC Voltmeters and Ammeters
  • 21.5 Null Measurements
  • 21.6 DC Circuits Containing Resistors and Capacitors
  • Introduction to Magnetism
  • 22.1 Magnets
  • 22.2 Ferromagnets and Electromagnets
  • 22.3 Magnetic Fields and Magnetic Field Lines
  • 22.4 Magnetic Field Strength: Force on a Moving Charge in a Magnetic Field
  • 22.5 Force on a Moving Charge in a Magnetic Field: Examples and Applications
  • 22.6 The Hall Effect
  • 22.7 Magnetic Force on a Current-Carrying Conductor
  • 22.8 Torque on a Current Loop: Motors and Meters
  • 22.9 Magnetic Fields Produced by Currents: Ampere’s Law
  • 22.10 Magnetic Force between Two Parallel Conductors
  • 22.11 More Applications of Magnetism
  • Introduction to Electromagnetic Induction, AC Circuits and Electrical Technologies
  • 23.1 Induced Emf and Magnetic Flux
  • 23.2 Faraday’s Law of Induction: Lenz’s Law
  • 23.3 Motional Emf
  • 23.4 Eddy Currents and Magnetic Damping
  • 23.5 Electric Generators
  • 23.6 Back Emf
  • 23.7 Transformers
  • 23.8 Electrical Safety: Systems and Devices
  • 23.9 Inductance
  • 23.10 RL Circuits
  • 23.11 Reactance, Inductive and Capacitive
  • 23.12 RLC Series AC Circuits
  • Introduction to Electromagnetic Waves
  • 24.1 Maxwell’s Equations: Electromagnetic Waves Predicted and Observed
  • 24.2 Production of Electromagnetic Waves
  • 24.3 The Electromagnetic Spectrum
  • 24.4 Energy in Electromagnetic Waves
  • Introduction to Geometric Optics
  • 25.1 The Ray Aspect of Light
  • 25.2 The Law of Reflection
  • 25.3 The Law of Refraction
  • 25.4 Total Internal Reflection
  • 25.5 Dispersion: The Rainbow and Prisms
  • 25.6 Image Formation by Lenses
  • 25.7 Image Formation by Mirrors
  • Introduction to Vision and Optical Instruments
  • 26.1 Physics of the Eye
  • 26.2 Vision Correction
  • 26.3 Color and Color Vision
  • 26.4 Microscopes
  • 26.5 Telescopes
  • 26.6 Aberrations
  • Introduction to Wave Optics
  • 27.1 The Wave Aspect of Light: Interference
  • 27.2 Huygens's Principle: Diffraction
  • 27.3 Young’s Double Slit Experiment
  • 27.4 Multiple Slit Diffraction
  • 27.5 Single Slit Diffraction
  • 27.6 Limits of Resolution: The Rayleigh Criterion
  • 27.7 Thin Film Interference
  • 27.8 Polarization
  • 27.9 *Extended Topic* Microscopy Enhanced by the Wave Characteristics of Light
  • Introduction to Special Relativity
  • 28.1 Einstein’s Postulates
  • 28.2 Simultaneity And Time Dilation
  • 28.3 Length Contraction
  • 28.4 Relativistic Addition of Velocities
  • 28.5 Relativistic Momentum
  • 28.6 Relativistic Energy
  • Introduction to Quantum Physics
  • 29.1 Quantization of Energy
  • 29.2 The Photoelectric Effect
  • 29.3 Photon Energies and the Electromagnetic Spectrum
  • 29.4 Photon Momentum
  • 29.5 The Particle-Wave Duality
  • 29.6 The Wave Nature of Matter
  • 29.7 Probability: The Heisenberg Uncertainty Principle
  • 29.8 The Particle-Wave Duality Reviewed
  • Introduction to Atomic Physics
  • 30.1 Discovery of the Atom
  • 30.2 Discovery of the Parts of the Atom: Electrons and Nuclei
  • 30.3 Bohr’s Theory of the Hydrogen Atom
  • 30.4 X Rays: Atomic Origins and Applications
  • 30.5 Applications of Atomic Excitations and De-Excitations
  • 30.6 The Wave Nature of Matter Causes Quantization
  • 30.7 Patterns in Spectra Reveal More Quantization
  • 30.8 Quantum Numbers and Rules
  • 30.9 The Pauli Exclusion Principle
  • Introduction to Radioactivity and Nuclear Physics
  • 31.1 Nuclear Radioactivity
  • 31.2 Radiation Detection and Detectors
  • 31.3 Substructure of the Nucleus
  • 31.4 Nuclear Decay and Conservation Laws
  • 31.5 Half-Life and Activity
  • 31.6 Binding Energy
  • 31.7 Tunneling
  • Introduction to Applications of Nuclear Physics
  • 32.1 Medical Imaging and Diagnostics
  • 32.2 Biological Effects of Ionizing Radiation
  • 32.3 Therapeutic Uses of Ionizing Radiation
  • 32.4 Food Irradiation
  • 32.5 Fusion
  • 32.6 Fission
  • 32.7 Nuclear Weapons
  • Introduction to Particle Physics
  • 33.1 The Yukawa Particle and the Heisenberg Uncertainty Principle Revisited
  • 33.2 The Four Basic Forces
  • 33.3 Accelerators Create Matter from Energy
  • 33.4 Particles, Patterns, and Conservation Laws
  • 33.5 Quarks: Is That All There Is?
  • 33.6 GUTs: The Unification of Forces
  • Introduction to Frontiers of Physics
  • 34.1 Cosmology and Particle Physics
  • 34.2 General Relativity and Quantum Gravity
  • 34.3 Superstrings
  • 34.4 Dark Matter and Closure
  • 34.5 Complexity and Chaos
  • 34.6 High-temperature Superconductors
  • 34.7 Some Questions We Know to Ask
  • A | Atomic Masses
  • B | Selected Radioactive Isotopes
  • C | Useful Information
  • D | Glossary of Key Symbols and Notation

If we are interested in how heat transfer is converted into doing work, then the conservation of energy principle is important. The first law of thermodynamics applies the conservation of energy principle to systems where heat transfer and doing work are the methods of transferring energy into and out of the system. The first law of thermodynamics states that the change in internal energy of a system equals the net heat transfer into the system minus the net work done by the system. In equation form, the first law of thermodynamics is

Here Δ U Δ U size 12{ΔU} {} is the change in internal energy U U size 12{U} {} of the system. Q Q size 12{Q} {} is the net heat transferred into the system —that is, Q Q size 12{Q} {} is the sum of all heat transfer into and out of the system. W W size 12{W} {} is the net work done by the system —that is, W W size 12{W} {} is the sum of all work done on or by the system. We use the following sign conventions: if Q Q size 12{Q} {} is positive, then there is a net heat transfer into the system; if W W size 12{W} {} is positive, then there is net work done by the system. So positive Q Q size 12{Q} {} adds energy to the system and positive W W size 12{W} {} takes energy from the system. Thus Δ U = Q − W Δ U = Q − W size 12{ΔU=Q - W} {} . Note also that if more heat transfer into the system occurs than work done, the difference is stored as internal energy. Heat engines are a good example of this—heat transfer into them takes place so that they can do work. (See Figure 15.3 .) We will now examine Q Q size 12{Q} {} , W W size 12{W} {} , and Δ U Δ U size 12{ΔU} {} further.

Making Connections: Law of Thermodynamics and Law of Conservation of Energy

The first law of thermodynamics is actually the law of conservation of energy stated in a form most useful in thermodynamics. The first law gives the relationship between heat transfer, work done, and the change in internal energy of a system.

Heat Q and Work W

Heat transfer ( Q Q size 12{Q} {} ) and doing work ( W W size 12{W} {} ) are the two everyday means of bringing energy into or taking energy out of a system. The processes are quite different. Heat transfer, a less organized process, is driven by temperature differences. Work, a quite organized process, involves a macroscopic force exerted through a distance. Nevertheless, heat and work can produce identical results. For example, both can cause a temperature increase. Heat transfer into a system, such as when the Sun warms the air in a bicycle tire, can increase its temperature, and so can work done on the system, as when the bicyclist pumps air into the tire. Once the temperature increase has occurred, it is impossible to tell whether it was caused by heat transfer or by doing work. This uncertainty is an important point. Heat transfer and work are both energy in transit—neither is stored as such in a system. However, both can change the internal energy U U size 12{U} {} of a system. Internal energy is a form of energy completely different from either heat or work.

Internal Energy U

We can think about the internal energy of a system in two different but consistent ways. The first is the atomic and molecular view, which examines the system on the atomic and molecular scale. The internal energy U U size 12{U} {} of a system is the sum of the kinetic and potential energies of its atoms and molecules. Recall that kinetic plus potential energy is called mechanical energy. Thus internal energy is the sum of atomic and molecular mechanical energy. Because it is impossible to keep track of all individual atoms and molecules, we must deal with averages and distributions. A second way to view the internal energy of a system is in terms of its macroscopic characteristics, which are very similar to atomic and molecular average values.

Macroscopically, we define the change in internal energy Δ U Δ U size 12{ΔU} {} to be that given by the first law of thermodynamics:

Many detailed experiments have verified that Δ U = Q − W Δ U = Q − W size 12{ΔU=Q - W} {} , where Δ U Δ U size 12{ΔU} {} is the change in total kinetic and potential energy of all atoms and molecules in a system. It has also been determined experimentally that the internal energy U U size 12{U} {} of a system depends only on the state of the system and not how it reached that state . More specifically, U U size 12{U} {} is found to be a function of a few macroscopic quantities (pressure, volume, and temperature, for example), independent of past history such as whether there has been heat transfer or work done. This independence means that if we know the state of a system, we can calculate changes in its internal energy U U size 12{U} {} from a few macroscopic variables.

Making Connections: Macroscopic and Microscopic

In thermodynamics, we often use the macroscopic picture when making calculations of how a system behaves, while the atomic and molecular picture gives underlying explanations in terms of averages and distributions. We shall see this again in later sections of this chapter. For example, in the topic of entropy, calculations will be made using the atomic and molecular view.

To get a better idea of how to think about the internal energy of a system, let us examine a system going from State 1 to State 2. The system has internal energy U 1 U 1 size 12{U rSub { size 8{1} } } {} in State 1, and it has internal energy U 2 U 2 size 12{U rSub { size 8{2} } } {} in State 2, no matter how it got to either state. So the change in internal energy Δ U = U 2 − U 1 Δ U = U 2 − U 1 size 12{ΔU=U rSub { size 8{2} } - U rSub { size 8{1} } } {} is independent of what caused the change. In other words, Δ U Δ U size 12{ΔU} {} is independent of path . By path, we mean the method of getting from the starting point to the ending point. Why is this independence important? Note that Δ U = Q − W Δ U = Q − W size 12{ΔU=Q - W} {} . Both Q Q size 12{Q} {} and W W size 12{W} {} depend on path , but Δ U Δ U size 12{ΔU} {} does not. This path independence means that internal energy U U size 12{U} {} is easier to consider than either heat transfer or work done.

Example 15.1

Calculating change in internal energy: the same change in u u size 12{u} {} is produced by two different processes.

(a) Suppose there is heat transfer of 40.00 J to a system, while the system does 10.00 J of work. Later, there is heat transfer of 25.00 J out of the system while 4.00 J of work is done on the system. What is the net change in internal energy of the system?

(b) What is the change in internal energy of a system when a total of 150.00 J of heat transfer occurs out of (from) the system and 159.00 J of work is done on the system? (See Figure 15.4 ).

In part (a), we must first find the net heat transfer and net work done from the given information. Then the first law of thermodynamics ( Δ U = Q − W ) ( Δ U = Q − W size 12{ΔU=Q - W} {} ) can be used to find the change in internal energy. In part (b), the net heat transfer and work done are given, so the equation can be used directly.

Solution for (a)

The net heat transfer is the heat transfer into the system minus the heat transfer out of the system, or

Similarly, the total work is the work done by the system minus the work done on the system, or

Thus the change in internal energy is given by the first law of thermodynamics:

We can also find the change in internal energy for each of the two steps. First, consider 40.00 J of heat transfer in and 10.00 J of work out, or

Now consider 25.00 J of heat transfer out and 4.00 J of work in, or

The total change is the sum of these two steps, or

Discussion on (a)

No matter whether you look at the overall process or break it into steps, the change in internal energy is the same.

Solution for (b)

Here the net heat transfer and total work are given directly to be Q = – 150 . 00 J Q = – 150 . 00 J size 12{Q"=-""150" "." "00"" J"} {} and W = – 159 . 00 J W = – 159 . 00 J size 12{W"=-""159" "." "00"" J"} {} , so that

Discussion on (b)

A very different process in part (b) produces the same 9.00-J change in internal energy as in part (a). Note that the change in the system in both parts is related to Δ U Δ U size 12{ΔU} {} and not to the individual Q Q size 12{Q} {} s or W W size 12{W} {} s involved. The system ends up in the same state in both (a) and (b). Parts (a) and (b) present two different paths for the system to follow between the same starting and ending points, and the change in internal energy for each is the same—it is independent of path.

Human Metabolism and the First Law of Thermodynamics

Human metabolism is the conversion of food into heat transfer, work, and stored fat. Metabolism is an interesting example of the first law of thermodynamics in action. We now take another look at these topics via the first law of thermodynamics. Considering the body as the system of interest, we can use the first law to examine heat transfer, doing work, and internal energy in activities ranging from sleep to heavy exercise. What are some of the major characteristics of heat transfer, doing work, and energy in the body? For one, body temperature is normally kept constant by heat transfer to the surroundings. This means Q Q size 12{Q} {} is negative. Another fact is that the body usually does work on the outside world. This means W W size 12{W} {} is positive. In such situations, then, the body loses internal energy, since Δ U = Q − W Δ U = Q − W size 12{ΔU=Q - W} {} is negative.

Now consider the effects of eating. Eating increases the internal energy of the body by adding chemical potential energy (this is an unromantic view of a good steak). The body metabolizes all the food we consume. Basically, metabolism is an oxidation process in which the chemical potential energy of food is released. This implies that food input is in the form of work. Food energy is reported in a special unit, known as the Calorie. This energy is measured by burning food in a calorimeter, which is how the units are determined.

In chemistry and biochemistry, one calorie (spelled with a lowercase c) is defined as the energy (or heat transfer) required to raise the temperature of one gram of pure water by one degree Celsius. Nutritionists and weight-watchers tend to use the dietary calorie, which is frequently called a Calorie (spelled with a capital C). One food Calorie is the energy needed to raise the temperature of one kilogram of water by one degree Celsius. This means that one dietary Calorie is equal to one kilocalorie for the chemist, and one must be careful to avoid confusion between the two.

Again, consider the internal energy the body has lost. There are three places this internal energy can go—to heat transfer, to doing work, and to stored fat (a tiny fraction also goes to cell repair and growth). Heat transfer and doing work take internal energy out of the body, and food puts it back. If you eat just the right amount of food, then your average internal energy remains constant. Whatever you lose to heat transfer and doing work is replaced by food, so that, in the long run, Δ U = 0 Δ U = 0 size 12{ΔU=0} {} . If you overeat repeatedly, then Δ U Δ U size 12{ΔU} {} is always positive, and your body stores this extra internal energy as fat. The reverse is true if you eat too little. If Δ U Δ U size 12{ΔU} {} is negative for a few days, then the body metabolizes its own fat to maintain body temperature and do work that takes energy from the body. This process is how dieting produces weight loss.

Life is not always this simple, as any dieter knows. The body stores fat or metabolizes it only if energy intake changes for a period of several days. Once you have been on a major diet, the next one is less successful because your body alters the way it responds to low energy intake. Your basal metabolic rate (BMR) is the rate at which food is converted into heat transfer and work done while the body is at complete rest. The body adjusts its basal metabolic rate to partially compensate for over-eating or under-eating. The body will decrease the metabolic rate rather than eliminate its own fat to replace lost food intake. You will chill more easily and feel less energetic as a result of the lower metabolic rate, and you will not lose weight as fast as before. Exercise helps to lose weight, because it produces both heat transfer from your body and work, and raises your metabolic rate even when you are at rest. Weight loss is also aided by the quite low efficiency of the body in converting internal energy to work, so that the loss of internal energy resulting from doing work is much greater than the work done. It should be noted, however, that living systems are not in thermal equilibrium.

The body provides us with an excellent indication that many thermodynamic processes are irreversible . An irreversible process can go in one direction but not the reverse, under a given set of conditions. For example, although body fat can be converted to do work and produce heat transfer, work done on the body and heat transfer into it cannot be converted to body fat. Otherwise, we could skip lunch by sunning ourselves or by walking down stairs. Another example of an irreversible thermodynamic process is photosynthesis. This process is the intake of one form of energy—light—by plants and its conversion to chemical potential energy. Both applications of the first law of thermodynamics are illustrated in Figure 15.5 . One great advantage of conservation laws such as the first law of thermodynamics is that they accurately describe the beginning and ending points of complex processes, such as metabolism and photosynthesis, without regard to the complications in between. Table 15.1 presents a summary of terms relevant to the first law of thermodynamics.

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3.A: The First Law of Thermodynamics (Answer)

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Check Your Understanding

3.1. \(\displaystyle p_2(V_2−V_1)\)

3.2. Line 1, \(\displaystyle ΔE_{int}=40J\);

line 2, \(\displaystyle W=50J\) and \(\displaystyle ΔE_{int}=40J\);

line 3, \(\displaystyle Q=80J\) and \(\displaystyle ΔE_{int}=40J\); and

line 4, \(\displaystyle Q=0\) and \(\displaystyle ΔE_{int}=40J\)

3.3. So that the process is represented by the curve \(\displaystyle p=nRT/V\) on the pV plot for the evaluation of work.

3.4. \(\displaystyle 1.26×10-^3J\).

Conceptual Questions

1. a. SE; b. ES; c. ES

3. Some of the energy goes into changing the phase of the liquid to gas.

5. Yes, as long as the work done equals the heat added there will be no change in internal energy and thereby no change in temperature. When water freezes or when ice melts while removing or adding heat, respectively, the temperature remains constant.

7. If more work is done on the system than heat added, the internal energy of the system will actually decrease.

9. The system must be in contact with a heat source that allows heat to flow into the system.

11. Isothermal processes must be slow to make sure that as heat is transferred, the temperature does not change. Even for isobaric and isochoric processes, the system must be in thermal equilibrium with slow changes of thermodynamic variables.

13. Typically \(\displaystyle C_p\) is greater than \(\displaystyle C_V\) because when expansion occurs under constant pressure, it does work on the surroundings. Therefore, heat can go into internal energy and work. Under constant volume, all heat goes into internal energy. In this example, water contracts upon heating, so if we add heat at constant pressure, work is done on the water by surroundings and therefore, \(\displaystyle C_p\) is less than \(\displaystyle C_V\).

15. No, it is always greater than 1.

17. An adiabatic process has a change in temperature but no heat flow. The isothermal process has no change in temperature but has heat flow.

19. \(\displaystyle p(V−b)=−c_T\) is the temperature scale desired and mirrors the ideal gas if under constant volume.

21. \(\displaystyle V−bpT+cT^2=0\)

25. 1.4 times

27. pVln(4)

29. a. 160 J; b. –160 J

31. \(\displaystyle W=900J\)

The figure is a plot of pressure, p, in atmospheres on the vertical axis as a function of volume, V, in Liters on the horizontal axis. The horizontal volume scale runs from 0 to 10 Liters, and the vertical pressure scale runs from 0 to 2 atmospheres. Four segments, A, B, C, and D are labeled. Segment A is a horizontal line with an arrow to the right, extending from 4 L to 10 L at a constant pressure of 2 atmospheres. Segment B is a vertical line with an arrow downward, extending from 2 atmospheres to 0.5 atmospheres at a constant 10 L.  Segment C is a horizontal line with an arrow to the left, extending from 10 L to 4 L at a constant pressure of 0.5 atmospheres. Segment D is a vertical line with an arrow upward, extending from 0.5 atmospheres to 2 atmospheres at a constant 4 L.

33. \(\displaystyle 3.53×10^4J\)

35. a. 1:1;

37. a. 600 J;

41. a. 600 J;

c. –160 J

45. a. 150 J;

47. No work is done and they reach the same common temperature.

49. 54,500 J

51. a. \(\displaystyle (p_1+3V^2_1)(V_2−V_1)−3V_1(V^2_2−V^2_1)+(V^3_2−V^3_1)\);

b. \(\displaystyle \frac{3}{2}(p_2V_2−p_1V_1)\);

c. the sum of parts (a) and (b); d. \(\displaystyle T_1=\frac{p_1V_1}{nR}\) and \(\displaystyle T_2=\frac{p_2V_2}{nR}\)

The figure is a plot of pressure, p in MegaPascals, on the vertical axis as a function of volume, V in Liters, on the horizontal axis. The horizontal volume scale runs from 0 to 6. The vertical pressure scale runs from 0 to 3. Two points, A at 2 Liters, 3 MegaPascals, and B at 6 Liters, and an unlabeled pressure, are shown and are connected by a curve. The curve is monotonically decreasing and concave up.

b. \(\displaystyle W=4.39kJ,ΔE_{int}=−4.39kJ\)

55. a. 1660 J;

b. −2730 J;

c. It does not depend on the process.

57. a. 700 J;

59. a. −3 400 J;

b. 3400 J enters the gas

63. a. 370 J;

67. pressure decreased by 0.31 times the original pressure

69. \(\displaystyle γ=0.713\)

The figure is a plot of pressure, p, in atmospheres on the vertical axis as a function of volume, V, in liters on the horizontal axis. The horizontal volume scale runs from 0 to 20, and the vertical pressure scale runs from 0 to 9. The data from the previous table is plotted as points and the equation y equals 8.4372 x to the minus 0.713 power is plotted as a curve. The points all lie on or very close to the curve.

73. An adiabatic expansion has less work done and no heat flow, thereby a lower internal energy comparing to an isothermal expansion which has both heat flow and work done. Temperature decreases during adiabatic expansion.

75. Isothermal has a greater final pressure and does not depend on the type of gas.

The figure is a plot of pressure, p, in atmospheres on the vertical axis as a function of volume, V, in liters on the horizontal axis. The horizontal, V, axis runs from 1.0 to 2.0. The vertical, p, axis runs from 0 to about 40. Two isotherms are shown. One isotherm is for T equal to 500 K, with the pressure starting at about 40 atmospheres when the volume is 1.0 Liter and decreasing with volume to about 25 atmospheres at 2.0 liters. The second isotherm is for T equal to 300 K, with the pressure starting at about 25 atmospheres when the volume is 1.0 Liter and decreasing with volume to a little over 10 atmospheres at 2.0 liters. A third plot, labeled “Adiabatic” starts with the 500 K isotherm, at 1.0 L and about 40 atmospheres, and ends with the 300 K isotherm, at 2.0 L and just over 10 atmospheres.

Additional Problems

79. a. \(\displaystyle W_{AB}=0,W_{BC}=2026J,W_{AD}=810.4J,W_{DC}=0;\)

b. \(\displaystyle ΔE_{AB}=3600J,ΔE_{BC}=374J\);

c. \(\displaystyle ΔE_{AC}=3974J\);

d. \(\displaystyle Q_{ADC}=4784J\);

e. No, because heat was added for both parts AD and DC . There is not enough information to figure out how much is from each segment of the path.

83. a. 59.5 J;

85. \(\displaystyle 2.4×10^3J\)

87. a. 15,000 J;

b. 10,000 J;

c. 25,000 J

91. A cylinder containing three moles of nitrogen gas is heated at a constant pressure of 2 atm. a. −1220 J; b. +1220 J

93. a. 7.6 L, 61.6 K;

c. \(\displaystyle 3.63L⋅atm=367J\);

d. −367 J

Challenge Problems

95. a. 1700 J; b. 1200 J; c. 2400 J

97. a. 2.2 mol;

b. \(\displaystyle V_A=6.7×10^{−2}m^3, V_B=3.3×10^{−2}m^3\);

c. \(\displaystyle T_A=2400K,T_B=397K\); d. 26,000 J

Contributors and Attributions

Samuel J. Ling (Truman State University), Jeff Sanny (Loyola Marymount University), and Bill Moebs with many contributing authors. This work is licensed by OpenStax University Physics under a  Creative Commons Attribution License (by 4.0) .

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1st Law of Thermodynamics

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To understand and perform any sort of thermodynamic calculation, we must first understand the fundamental laws and concepts of thermodynamics. For example, work and heat are interrelated concepts. Heat is the transfer of thermal energy between two bodies that are at different temperatures and is not equal to thermal energy. Work is the force used to transfer energy between a system and its surroundings and is needed to create heat and the transfer of thermal energy. Both work and heat together allow systems to exchange energy. The relationship between the two concepts can be analyzed through the topic of Thermodynamics, which is the scientific study of the interaction of heat and other types of energy.

Introduction

To understand the relationship between work and heat, we need to understand a third, linking factor: the change in internal energy. Energy cannot be created nor destroyed, but it can be converted or transferred. Internal energy refers to all the energy within a given system, including the kinetic energy of molecules and the energy stored in all of the chemical bonds between molecules. With the interactions of heat, work and internal energy, there are energy transfers and conversions every time a change is made upon a system. However, no net energy is created or lost during these transfers.

Law of Thermodynamics

The First Law of Thermodynamics states that energy can be converted from one form to another with the interaction of heat, work and internal energy, but it cannot be created nor destroyed, under any circumstances. Mathematically, this is represented as

\[ \Delta U=q + w \label{1}\]

  • \(ΔU\) is the total change in internal energy of a system,
  • \(q\) is the heat exchanged between a system and its surroundings, and
  • \(w\) is the work done by or on the system.

Work is also equal to the negative external pressure on the system multiplied by the change in volume:

\[ w=-p \Delta V \label{2}\]

where \(P\) is the external pressure on the system, and \(ΔV\) is the change in volume. This is specifically called "pressure-volume" work.

The internal energy of a system would decrease if the system gives off heat or does work. Therefore, internal energy of a system increases when the heat increases (this would be done by adding heat into a system). The internal energy would also increase if work were done onto a system. Any work or heat that goes into or out of a system changes the internal energy. However, since energy is never created nor destroyed (thus, the first law of thermodynamics), the change in internal energy always equals zero. If energy is lost by the system, then it is absorbed by the surroundings. If energy is absorbed into a system, then that energy was released by the surroundings:

\[\Delta U_{system} = -\Delta U_{surroundings} \]

where ΔU system is the total internal energy in a system, and ΔU surroundings is the total energy of the surroundings.

The above figure is a visual example of the First Law of Thermodynamics. The blue cubes represent the system and the yellow circles represent the surroundings around the system. If energy is lost by the cube system then it is gained by the surroundings. Energy is never created nor destroyed. Since the area of the clue cube decreased the visual area of the yellow circle increased. This symbolizes how energy lost by a system is gained by the surroundings. The affects of different surroundings and changes on a system help determine the increase or decrease of internal energy, heat and work.

Example \(\PageIndex{1}\)

A gas in a system has constant pressure. The surroundings around the system lose 62 J of heat and does 474 J of work onto the system. What is the internal energy of the system?

To find internal energy, ΔU, we must consider the relationship between the system and the surroundings. Since the First Law of Thermodynamics states that energy is not created nor destroyed we know that anything lost by the surroundings is gained by the system. The surrounding area loses heat and does work onto the system. Therefore, q and w are positive in the equation ΔU=q+w because the system gains heat and gets work done on itself.

\[\begin{align} ΔU &= (62\,J) + (474\,J) \\[4pt] &= 536\,J \end{align}\]

Example \(\PageIndex{2}\)

A system has constant volume (ΔV=0) and the heat around the system increases by 45 J.

  • What is the sign for heat (q) for the system?
  • What is ΔU equal to?
  • What is the value of internal energy of the system in Joules?

Since the system has constant volume (ΔV=0) the term -PΔV=0 and work is equal to zero. Thus, in the equation ΔU=q+w w=0 and ΔU=q. The internal energy is equal to the heat of the system. The surrounding heat increases, so the heat of the system decreases because heat is not created nor destroyed. Therefore, heat is taken away from the system making it exothermic and negative. The value of Internal Energy will be the negative value of the heat absorbed by the surroundings.

  • negative (q<0)
  • ΔU=q + (-PΔV) = q+ 0 = q
  • ΔU = -45J

Outside Links

  • Hamby, Marcy. "Understanding the language: Problem solving and the first law of thermodynamics." J. Chem. Educ. 1990 : 67, 923.
  • Chang, Raymond. Chemistry: Ninth Edition
  • Petrucci, Harwood, Herring, Madura. General Chemistry: Ninth Edition

Contributors and Attributions

  • Lauren Boyle (UCD)

First Law Of Thermodynamics

The First Law of Thermodynamics states that heat is a form of energy, and thermodynamic processes are therefore subject to the principle of conservation of energy. This means that heat energy cannot be created or destroyed. It can, however, be transferred from one location to another and converted to and from other forms of energy.

Introducing State Variables

Thermodynamic state variables are the macroscopic quantities that determine a system’s thermodynamic equilibrium state. A system not in equilibrium cannot be described by state variables. State variables can further be classified as intensive or extensive variables. Intensive variables are independent of the dimensions of the system like pressure and temperature, while extensive variables depend on dimensions of the system like volume, mass, internal energy etc.

Explaining the first law of thermodynamics

The first law of thermodynamics relates to heat, internal energy, and work.

The first law of thermodynamics, also known as the law of conservation of energy, states that energy can neither be created nor destroyed, but it can be changed from one form to another.

First Law of Thermodynamics

It can be represented mathematically as

  • ΔQ is the heat given or lost
  • ΔU is the change in internal energy
  • W is the work done

We can also represent the above equation as follows,

So we can infer from the above equation that the quantity (ΔQ – W) is independent of the path taken to change the state. Further, we can say that internal energy increases when the heat is given to the system and vice versa.

Sign Conventions

The table below shows the appropriate sign conventions for all three quantities under different conditions:

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First Law of Thermodynamics Solved Examples

1. Calculate the change in the system’s internal energy if 3000 J of heat is added to a system and a work of 2500 J is done.

Solution: The following sign conventions are followed in the numerical: Solution: The following sign conventions are followed in the numerical:

  • Q is positive as heat is added to the system
  • W is positive if work is done on the system

Hence, the change in internal energy is given as: \(\begin{array}{l}\Delta U=3000-2500\end{array} \) \(\begin{array}{l}\Delta U=500\end{array} \) The internal energy of the system is 500 J.

2. What is the change in the internal energy of the system if 2000 J of heat leaves the system and 2500 J of work is done on the system? Solution: The change in the internal energy of the system can be identified using the formula:

Substituting the values in the following equation, we get

ΔU = -2000-(-3000)

ΔU = -2000+3000

ΔU = 1000 Joule

Internal energy increases by 4500 Joules.

first law of thermodynamics assignment quizlet

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STUDYPLAYTerms in this set (30)Related questionsSets found in the same folderSets with similar termsOther sets by this creatorOther Quizlet setsVerified questionsVOCABULARYFrom the list below, supply the words needed to complete the paragraph. Some words will not be used. genial, novice, enraptured, juggernaut, nocturnal, marital, obstreperous, levity. Despite her position as regional manager for Tyndall Systems, Shawna felt like a(n) _____ every time she attended the monthly sales meeting at Tyndall corporate headquarters. Perhaps she was just getting old, she reasoned, but she knew that few could endure her ____ schedule six days out of the week. Tyndall was a(n) _____ in the information technology arena, buying and consolidating other corporations and firing dissenters with impunity. Shawna told her husband that she would retire in two years; she hoped in time to mitigate their rapidly multiplying _____ problems. She was no longer the hard worker that Tyndall wanted for managing a regional hub, and the stress from trying to meet the demand had caused her once _____ manner to reverse-not that she needed it any more at the office.

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Upgrade to remove adsOnly RUB 2,325/yearSciencePhysicsThermodynamicsSTUDYFlashcardsLearnWriteSpellTestPLAYMatchGravityTerms in this set (17)What is the first law of thermodynamics and how does it relate to energy use? The first law of thermodynamics states that energy is conserved in a chemical reaction. It is neither created nor destroyed. Explain the difference between the spontaneity of a reaction and the speed at which the reaction occurs. Can a catalyst make a nonspontaneous reaction spontaneous? Spontaneity of a reaction is the direction in which and the extent to which a chemical reaction proceeds. Kinetics is the speed of the reaction (how fast). A reaction may be thermodynamically favorable, but may be very slow at a given temperature. While the rate of a spontaneous process can be increased by the use of a catalyst, a nonspontaneous process cannot be made to be spontaneous by the use of a catalyst. Provide the definition of the second law of thermodynamics. How does the second law explain why heat travels from a substance at higher temperature to one at lower temperature?

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What does the first law of thermodynamics state?

The laws of thermodynamics are deceptively simple to state, but they are far-reaching in their consequences. The first law asserts that if heat is recognized as a form of energy, then the total energy of a system plus its surroundings is conserved ; in other words, the total energy of the universe remains constant.

What does the 1st law of thermodynamics state quizlet?

The first law of thermodynamics, also known as Law of Conservation of Energy, states that energy can neither be created nor destroyed ; energy can only be transferred or changed from one form to another.

Which law of thermodynamics states that quizlet?

states that energy can not be created or destroyed ; it can only be redistributed or changed from one form to another. states that the disorder in the universe always increases. You just studied 5 terms!

What is first law of thermodynamics in physics?

The First Law of Thermodynamics states that heat is a form of energy , and thermodynamic processes are therefore subject to the principle of conservation of energy. This means that heat energy cannot be created or destroyed. ... "So, it's a restatement of conservation of energy."

What is first law of thermodynamics in chemistry?

The first law of thermodynamics, also known as Law of Conservation of Energy, states that energy can neither be created nor destroyed; energy can only be transferred or changed from one form to another . For example, turning on a light would seem to produce energy; however, it is electrical energy that is converted.

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    1 / 21 Flashcards Learn Test Match Q-Chat Created by Redonaj Taken from the MCAT Primer Question Book, provided by John Wetzel, an author at www.wikipremed.com. Students also viewed chapter 6 smartbook 56 terms Rachie6996 Preview physics ; second law of thermodynamics 11 terms lemiru-x thermodynamics First Law Of Thermodynamics 9 terms babyirl_

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    The first law of thermodynamics states that the change in internal energy of a system Δ U equals the net heat transfer into the system Q , plus the net work done on the system W . In equation form, the first law of thermodynamics is, Δ U = Q + W [Wait, why did my book/professor use a negative sign in this equation?]

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  8. 3.4: First Law of Thermodynamics

    Solution. From the first law, the change in the system's internal energy is. ΔEintAB = QAB −WAB = 400J − 100J = 300J. Δ E i n t A B = Q A B − W A B = 400 J − 100 J = 300 J. Consider a closed path that passes through the states A and B. Internal energy is a state function, so ΔEint Δ E i n t is zero for a closed path.

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    In equation form, the first law of thermodynamics is. size 12 {ΔU=Q - W} {} 15.1. Here size 12 {ΔU} {} is the change in internal energy size 12 {U} {} of the system. size 12 {Q} {} is the net heat transferred into the system —that is, size 12 {Q} {} is the sum of all heat transfer into and out of the system. size 12 {W} {} is the net work ...

  11. 3.A: The First Law of Thermodynamics (Answer)

    1. a. SE; b. ES; c. ES. 3. Some of the energy goes into changing the phase of the liquid to gas. 5. Yes, as long as the work done equals the heat added there will be no change in internal energy and thereby no change in temperature. When water freezes or when ice melts while removing or adding heat, respectively, the temperature remains constant.

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