Thermochemistry

1. Thermochemistry

2. Heat & Temperature • Joseph Black – explained heat in terms of a fluid (Lavoisier had called this fluid “caloric” from Latin word for heat. • Count Rumford – friction could convert mechanical energy into heat (motion as cause) • John Dalton – idea of atoms

3. Heat History • James Prescott Joule – tried to find the mechanical equivalent of heat (where a given amount of energy produces the same amount of heat) • James Clerk Maxwell – developed a solid explanation showing relationship between motion of atoms and heat.

4. Heat • Heat flows from hot to colder areas due to a temperature difference only (till thermal equilibrium is established). • Heat is a form of internal energy which is transferred from one object to another due to a difference intemperature between the objects.

5. Heat Content • The heat content of a substance is the total energy of all the particles of that substance. • The total energy combines both kinetic and potential energies.

6. Temperature • The temperature of a body of matter is a measure of the average kineticenergy of the random motion of its particles. • Temperature is that property of a substance which determines whether it is in thermal equilibrium with another object.

7. Thermal Equilibrium • Thermal equilibrium is the situation in which no heat moves from one object to another (they have the same temperature).

8. Thermometers • Thermometers work on idea of thermal expansion = the amount of expansion or contraction is always the same for the same increase or decrease in temperature. • 3 types: gas (air), liquid (Hg & alcohol), solid (bimetallic) • Know creation and calibration ideas

9. Temperature Conversions • K = °C + 273.15 • °C = K – 273.15 • °F = 9/5 °C + 32 • °C = 5/9 (°F - 32)

10. Scale Comparisons Fahrenheit Celsius Kelvin Boiling Pt. H2O 212 100 373 Body temp 98.6 37 310 Freezing 32 0 273 Coincidence -40 -40 233 Absolute zero -460 -273 0

11. Heat Units • 15° calories = the amount of heat needed to raise 1 gram of water from 14.5° to 15.5° C at 1 atmosphere of pressure • kilocalorie = kcal or Calorie = 1000 cal • 1 calorie = 4.185 Joule • 1 kcal = 4185 Joule

12. Specific Heat • Specific heat is the amount of heat needed to raise 1 gram of water 1° C at 1 atmosphere of pressure • What is the degree change if 1 calorie of heat is added to 1 gram samples of: water helium ice gold

13. Heat Flow (Q) • Q = m x ∆t x cp where Q = heat flowm = mass ∆t = change in temperature cp = specific heat

14. Principle of Heat Exchange • The amount of heat lost by a substance is equal to the amount of heat gained by the substance to which it is transferred. • m x ∆t x cp = m x ∆t x cp heat lost heat gained

15. Specific Heat Notes • Specific heat – how well a substance resist changing its temperature when it absorbs or releases heat • Water has high cp – results in coastal areas having milder climate than inland areas (coastal water temp. is quite stable which is favorable for marine life).

16. More Specific Heat • Organisms are primarily water – thus are able to resist more changes in their own temperature than if they were made of a liquid with a lower cp

17. Water and Heat • When calories of heat are added to water there is a small change in temperature because most of the heat energy is used to disrupt hydrogen bonds before water molecules can begin to move faster. • Temp. of water drops – many additional hydrogen bonds form, releasing a considerable amount of heat energy.

18. Absolute Zero

19. Material Data

20. Heat Transfer Mechanisms • Conduction – faster vibrating particles collide with less energetic neighbor and transfer energy to it • Convection – motion of hot fluid, displacing cold fluid in path setting up convection current • Radiation – energy transmitted by electromagnetic waves

21. Thermal Expansion of Water • From 0° to 4° the volume of water in a sample decreases (the greatest density is at 4° c) • Ice floats: body of water freezes from top down allowing life underneath to continue

22. Ice – Open Structure • Water mlcl can participate in 4 bonds with other water mlcl (solid mlcl can have as many as a dozen bonds with surrounding mlcl resulting in a more compact substance). • The spaces between mlcl in ice are greater than the same spaces in liquids.

23. Density of Ice • Density of ice increases from 0 to 4° as large clusters of mlcl break into smaller clusters that takes up less space in the aggregate. Above 4° normal thermal expansion is seen with a decrease in density.

24. Latent Heat • Heat of Fusion – amount of heat needed to change solid to liquid at its melting point • Heat of Vaporization – heat needed to change liquid to gas at boiling point • Heat of Sublimation – heat to change a solid to gas • Heat of Condensation – heat released when gas condenses to a liquid

25. Matter • Matter is defined as any material that has mass, occupies volume, and exhibits inertia (resistance to movement).

26. States of Matter • Solids – definite shape and volume, resist deformation • Very close spacing of particles • Particles appear to vibrate around fixed points • Particles vibrate faster at higher temp.

27. Types of Solids • Crystalline – particles arranged in regular, repeated patterns (long-rangeorder) example: NaCl (s) • Amorphous – solids that lack the definite arrangement of crystalline solids (have short-range order) • Examples: pitch, glass, plastics

28. Liquids • Definite volume, resist compression, take shape of container • Greater spacing between particles, particles appear to travel in straight line paths between collisions but appear to rotate and/or vibrate about moving points

29. Gases • Have no definite shape or volume, take shape and volume of container • Can be compressed or dispersed, particles vibrate very rapidly, relatively far apart • There are no intermolecular forces holding particles together

30. Plasma • Very high temperature ionized gas • No fixed volume or shape • Most are mixtures that are not easily containable • Particles are electrically charged and of low density • Example: the Milky Way

31. Energy • Energy – having the ability to do work • Work – a push or pull over a distance • Force – a push or pull • Momentum – mass x velocity • Linear momentum of a moving body is a measure of its tendency to continue in motion at a constant velocity

32. Potential Energy • Potential Energy – the energy a body possesses by virtue of its position, composition, and/or condition • P.E. is the stored energy • P.E. = mass x gravity x height

33. Kinetic Energy • K.E. = the energy of motion • K.E. is conserved in all elastic collisions • K.E. = ½ m v2 (m = mass, v = velocity) • Heat energy flows from hot objects to cooler ones by transfer of K.E. when particles collide (conduction).

34. Intermolecular Forces • P.E. forces that hold mlcl together and in correct position in solids. • P.E. forces that hold mlcl together in liquids. • These forces are between mlcl. • Gases have enough K.E. to prevent formation of these forces.

35. Kinetic Molecular Theory of Gases • Gases are mlcl in continuous motion. • An increase in temp. increases speed thus increasing K.E. • All gases are compressible • Gases display diffusion • Gases can be liquified (called liquifaction)

36. Closed System Info  Pressure, Volume, Temp. • Nothing escapes or enters system • All mlcl in motion (have K.E.) • Mlcl exert uniform pressure against walls of container • Mlcl exert pressure on other mlcl as they collide, push, bounce off other mlcl

37. Pressure • Pressure = Force / Area • Atmospheric Pressure = cumulative net force per area generated by weight of our atmosphere • Values = 14.7 lb/in2, 101.3 kPa, 760 mm of Hg, 1 atm, 1033 g/cm2

38. Gas Pressure • The pressure a gas exerts on the walls of its container is the sum of the forces acting (= the frequency of collisions plus the force of each collision) due to the random collisions of near limitless numbers of moving molecules.

39. Collisions • Inelastic collisions – the normal type in which objects lose energy and slow down • Elastic collisions – particles bounce off, exchange energies but there is no loss of energy (energy is conserved but may be redistributed)

40. Conservation in Collisions • Energy is conserved only in elastic collisions • Momentum is conserved in every collision in which there is no friction.

41. Gas Laws • Gay-Lussac P ≈ T • Holding volume constant, the pressure is proportional to the absolute temp. • P1 / T1 = P2 / T2

42. Gas Laws • Boyle’s Law V ≈ 1/P • If the temp. is held constant, the volume of a gas varies inversely with the pressure • P1V1 = P2V2

43. Gas Laws • Charles’ Law V ≈ T • If the pressure is held constant, the volume of a gas is proportional to its absolute temp. • V1 / T1 = V2 / T2 • For every degree increase in temp. the volume increases by 1/273 of its original volume

44. Better Gas Law Equations • Combined Gas Law: P1V1/T1 = P2V2/T2 • Ideal Gas Law: PV = n R T (where n = # moles, and R = gas constant)

45. Chemical Properties Chemical properties are those properties of a substance that can be determined by a chemical test. They are seen by the material’s tendency to change, either alone or by interaction, and in doing so form different materials.

46. Chemical Properties • Does the substance support combustion? Burn itself? • How does it react with acids? With oxygen? With electricity? • Examples: alcohol burns, wood decays, sodium explodes and burns in water

47. Physical Properties • Physical properties are those properties used in identifying substances when we use our senses. These do not require chemical analysis.

48. Physical Properties • Color, hardness, density, texture, magnetic attraction, solubility, taste, light transmission, viscosity, refractive index, specific heat, boiling point, melting-freezing point, odor, expansion-contraction coefficients

49. Physical Changes in State • This is a change in the physical properties of a substance without a change in the chemical composition. • The arrangement of molecules may be changed but the molecular makeup remains the same. • These changes involve intermolecular forces which increase or decrease during the change.

50. Physical Changes • Ice (0° C) + heat  steam (100° C) 36 g 25 920 cal 36 g • Steam (100° C)  ice (0° C) + heat 36 g 36 g 25 920 cal