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THERMODYNAMICS

THERMODYNAMICS . Thermodynamics (from Greek therme = heat, dynamis = strength, power) = branch of physic dealing with energy transformations from and into thermal energy; mechanics: mechanical (external) energies of systems, governed by Newton's laws;

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THERMODYNAMICS

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  1. THERMODYNAMICS • Thermodynamics • (from Greek therme = heat, dynamis = strength, power) = branch of physic dealing with energy transformations from and into thermal energy; • mechanics: mechanical (external) energies of systems, governed by Newton's laws; • thermodynamics: internal energy of systems and its relation to work; • keywords of thermodynamics: temperature, heat, internal/thermal energy, entropy • four laws of thermodynamics: • heat transfer, thermal equilibrium • energy conservation • not all thermal energy is useful; • impossibility to reach absolute zero temperature • topics to be discussed: • thermal energy, temperature, heat • 0th law • temperature scales • thermal expansion • heat capacity, specific heat • heat transfer: conduction, convection, radiation • 1st law • heat engines, efficiency • 2nd law, entropy

  2. thermal energy, temperature, heat • Brownian motion: • Robert Brown observed burlap seeds dancing in water (1827); explained by A. Einstein (1905); calculated mea net distance travelled by random motion; • experimental verification by Jean Perrin (1908). • thermal motion: • disorganized random motion of constituent atoms and molecules within body of matter; • thermal energy: • kinetic energy of thermal motion (translational, rotational, vibrational) associated with ensemble of particles • temperature: • is measure of average value of thermal energy of atoms and molecules (not total amount of thermal energy); • (temperature of a substance is independent of total number of atoms/molecules) • is a measure of the ability of randomly moving particles to impart thermal energy to a thermometer; • heat • = thermal energy transferred from a region of high temperature to region of lower temperature; • body stores thermal energy (internal energy); • heat = thermal energy “in transit”

  3. 0th law of thermodynamics • between bodies of different temperature (i.e. of different average internal thermal energy), heat will flow from the body of higher temperature to the body of lower temperature until the temperatures of the two bodies are the same; • then the bodies are in “thermal equilibrium” • two bodies are in thermal equilibrium (at same temperature) if there is no heat flow between them; • corollary: if two bodies are in thermal equilibrium with a third body, then they are in thermal equilibrium with each other. •  can use thermometer to compare temperature • note: • observation only shows that temperatures equalize - heat flow is hypothesis

  4. TEMPERATURE SCALES • Temperature: • was measured long before it was understood; • Galilei (around 1592): “device to measure degree of hotness”; inverted narrow-necked flask,warmed inhand, put upside down into liquid; liquid level indicates temperature; OK, but not calibrated. • Hooke, Huygens, Boyle (1665): “fixed points” -freezing or boiling point of water; • C. Renaldini (1694): use both freezing and boiling point. • Fahrenheit scale: • Gabriel Daniel Fahrenheit (Danzig, 1686-1736), glassblower and physicist; • reproducible thermometer using mercury (liquid throughout range) (around 1715) • 0 point: lowest temperature of winter of 1709, (using mix of water, ice, salt) • 96o= body temperature (96 divisible by 12, 8), • water freezes at 32oF, boils at 212oF • Celsius scale: • Anders Celsius (Swedish astronomer, 1701 - 1744) • 0o C = ice point (mixture of water and ice at 1 atm) • 100o C = boiling point of water at 1 atm. (1742) • relation between Fahrenheit and Celsius degrees: • TC = (5/9)(TF - 32 ) , TF = (9/5)TC + 32

  5. Temperature, cont’d • thermodynamic temperature scale • (absolute, Kelvin scale) • pressure vs temperature of gas at constant volume and volume vs temperature of gas at constant pressure extrapolate to zero at - 273.15o C • this is “absolute zero” • unit: Kelvin • Range of temperatures • highest temperature: in core of stars, 4109K; seems maximum; • hydrogen bomb ignites at , 4107K; • interior of Sun , 1.5106K; • plasma 105K; • 105K : clouds of atoms, ions, e, occasional molecule; • 5800 K: surface of the Sun; 5000 K: cool spots at surface of the Sun; evidence for some molecules; • 3000 K: water steam: about 1/4 of water molecules ruptured into atoms; • 2800 K: W light bulb filament; • 2000 K: molten lava; • 1520 oC: iron melts; 327 o C: lead melts; • 100oC (373 K): water boils; • 252 K: temp. of salt-ice mix; • 234 K: mercury freezes 194 K: dry ice freezes; • 77 K: nitrogen boils • 4 K: helium boils.

  6. THERMAL EXPANSION • solids, liquids and gases: • expand when heated • knowledge about this is old: e.g. red-hot iron rims put on wagon wheels; • thermometers are based on this; • heating  internal energy rises  vibrations have larger amplitude, equilibrium positions move farther apart. • typical metal expands by about 7% between 0 K and melting point. • L/L0 = T,  = coefficient of thermal expansion; • examples for values of  (in units of 10-6): • iron 10, • brass 19, • lead 30, • Pyrex glass 3, • ordinary glass 5 to 10, • concrete 10 to 14 • mercury 60, • ethanol 250 • have to account for this in construction, e.g. expansion joints at end of bridge, gaps in rails; also in dental fillings; • uses: thermostats, thermometers (bimetal strips) • anomaly of water: maximum of density at 4oC.

  7. HEAT CAPACITY • Heat capacity = • measure of ability of a substance to absorb thermal energy; • specific heat capacity = heat capacity per unit mass; • Q = c m T, Q = amount of thermal energy added, c = specific heat capacity, T = raise in temperature; • 1 calorie = 1 cal ( = 4.186 J) = thermal energy necessary to raise temperature of 1 gram of water by 1 degree Celsius; • 1 kcal = 1 Cal = thermal energy necessary to raise temperature of 1 kilogram of water by 1 degree Celsius; called “calorie” in nutrition; • water has high specific heat capacity moderating influence on climate • some values of specific heat capacity: • aluminum 0.21 • clay 0.22 • glass 0.20 • marble 0.21 • iron 0.11 • air 0.24 • water 1.00

  8. HEAT TRANSFER • Conduction: • = heat transfer by atomic/molecular collisions; • thermal conductivity = ability of substance to transmit heat, depends on atomic/molecular structure; • metals typically 400 times better than other solids; • most solids little better than liquids; • liquids about 10 times better than gases; • good heat conductor usually good electric conductor • Convection: • = heat transfer by motion of hot matter change of density of fluid (liquid or gas) due to heating; • flow of fluid up, away from heat source; • dominant mechanism for many heat loss processes in air; • examples: household radiator, hurricanes • purpose of fur, feathers, clothing, blankets:prevent convection • “chill-factor” • Radiation • = heat transfer by emission and absorption of electromagnetic radiation; e.g. Earth receives 1.4kW/m2 by radiation from the Sun.

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