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Temperature Definition Instrument Scales

Temperature Definition Instrument Scales. TEMPERATURE. Is associated with heat but it is NOT HEAT. IT IS NOT A FORM OF ENERGY!!!! ( Heat is) Review: What is KINETIC ENERGY?. KINETIC ENERGY (KE). Is associated with movement. If an object is moving fast has high KE

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Temperature Definition Instrument Scales

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  1. Temperature • Definition • Instrument • Scales

  2. TEMPERATURE • Is associated with heat but it is NOT HEAT. IT IS NOT A FORM OF ENERGY!!!! • ( Heat is) • Review: What is KINETIC ENERGY?

  3. KINETIC ENERGY (KE) • Is associated with movement. • If an object is moving fast has high KE • If an object is moving slowly it has low KE

  4. Temperature • In scientific measurements, the Celsius and Kelvin scales are most often used. • The Celsius scale is based on the properties of water. • 0C is the freezing point of water. • 100C is the boiling point of water.

  5. Temperature: A measure of the average kinetic energy of the particles in a sample. If an object is at HIGH temperature its particles are moving FAST At LOW temperature particles move SLOWLY

  6. Instrument to measure temperatureTHERMOMETER

  7. FIXED POINTS OF A THERMOMETER • BOILING POINT OF WATER • FREEZING POINT OF WATER

  8. Temperature • The Kelvin is the SI unit of temperature. • It is based on the properties of gases. • There are no negative Kelvin temperatures. • K = C + 273.15

  9. Temperature • The Fahrenheit scale is not used in scientific measurements. • F = 9/5(C) + 32 • C = 5/9(F − 32)

  10. 2. SI unit for temp. is the Kelvin a. K = C + 273 (10C = 283K) b. C = K – 273 (10K = -263C)

  11. DO NOW • What is 35ºC in Kelvin? • What is 10 K in ºC?

  12. What is the temperature of the beakers? • Which one needs more heat to boil? • What is the difference between heat and temperature?

  13. What is the relationship between heat and temperature?

  14. 1. Temperature is the average kinetic energy of the particles in a substance.

  15. Heat and Temperature • The temperature of an object tells us how HOT the object is. It is measured in degrees Celsius - °C, with a thermometer. • Temperature is NOT a form of energy. • HEAT IS A FORM OF ENERGY!!!

  16. So how temperature and heat are related? • Temperature is NOTthe same as HEAT ENERGYalthough the two quantities are related.

  17. They are both a the same temperature. Which contains more heat?

  18. Types of energy • POTENTIAL ENERGY : STORED ENERGY. The energy inside the substance. • KINETIC ENERGY : Associated with motion. Average KE = TEMPERATURE

  19. HEAT • The form of ENERGY that flows between two samples of matter due to their difference in temperature. • It is measured in Joules. Other units are calories and kJ. • It cannot be measured directly!

  20. Cup gets cooler while hand gets warmer Heat Heat always flows from warmer (HIGH T) to cooler(LOW T) objects. Ice gets warmer while hand gets cooler

  21. Energy • The ability to do work. • Energy is measured in Joules (J) • Any change re quires energy. Changes can be Exothermic or Endothermic • Exothermic changes- release or give off heat while they occur. (condensation, freezing) • Endothermic changes Absorb heat as they occur (melting, boiling)

  22. ENERGY AND CHEMICAL REACTIONS • EXOTHERMIC REACTIONS: reactions that RELEASE heat as they occur. Example: any combustion. • ENDOTHERMIC REACTIONS: reactions that ABSORB heat energy as they occur.

  23. Heating and Cooling • If an object has become hotter, it means that it has gained heat energy. • If an object cools down, it means it has lost energy

  24. 6. Specific Heat a. Some things heat up or cool down faster than others. Land heats up and cools down faster than water

  25. HEAT CAPACITY • The amount of heat needed to increase the temperature of a material by 1oC. It depends on the MASS and the CHEMICAL COMPOSITION of the material.

  26. SPECIFIC HEAT CAPACITY • The amount of heat needed to increase the temperature of 1 g of substance by 1 C. • Depends only on the chemical composition.

  27. SPECIFIC HEAT CAPACITY

  28. Why does water have such a high specific heat? water metal Water molecules form strong bonds with each other; therefore it takes more heat energy to break them. Metals have weak bonds and do not need as much energy to break them.

  29. How to calculate changes in HEAT The heat absorbed or released in a chemical reaction Q = m x T x Cp Q = HEAT m = mass of substance T= change in temperature (Tf – Ti) Cp = specific heat of substance

  30. Calorimeter

  31. Bomb Calorimetry A more sophisticated model is the bomb calorimeter, it has a chamber where a chemical reaction takes place and a device to start the reaction.

  32. HW answers • 1) 42000 J • 2) 6300J • 3) 5 C • 4) 20 C • 5) 336 g • 6) 14700 J =14.7 kJ • 7) 14.24 % • 8) b • 9)b 10 7J (d) • 1 • 4 • 3 • 4 • 3 • 3 • 2 • 2 • 4

  33. Aim: How to determine the melting point of a substance? Challenge: To determine the melting point of water. Your write up should include the following components: I. TITLE II. PURPOSE MATERIALS- IV. PROCEDURE V. DATA VI. CALCULATIONS/GRAPH VII.CONCLUSION (including a discussion of sources of error.)

  34. What is the freezing point of water? • How would the mass of the sample affect the melting point? Explain. • Describe the steps needed to determine the boiling point of water. • Describe the phase changes that take place as water melts and as water boils.

  35. Vocabulary • condensation • deposition • freezing • fusion • gaseous phase • heat • heat of fusion • heat of vaporization • kinetic molecular theory • sublimation • temperature • vaporization • solidification • melt • solid phase • liquid phase

  36. October 17 • Objective: What are the characteristics of each state of matter? • How to determine the amount of heat needed to change state?

  37. Phases of Matter • The structure and arrangement of particles and their interactions determine the physical state of a substance at a given temperature and pressure. • The three phases of matter (solids, liquids, and gases) have different properties.

  38. States of Matter

  39. States of Matter .

  40. Particle diagrams

  41. Phase Change • DURING A PHASE CHANGE TEMPERATURE REMAINS CONSTANT • Phase changes • Melting or Fusion: Phase change from solid to liquid. Endothermic. • Vaporization – Phase change from liquid to gas. It occurs during boiling. Endothermic.

  42. Phase Changes • Evaporation: from liquid to gas. It occurs at all temperatures and ONLY at the surface of the liquid. Endothermic. • Condensation: From gas to liquid. Exothermic. • Freezing, solidification or crystallization: from liquid to gas. Exothermic.

  43. Special Phase Changes • Sublimation: from solid directly to gas without changing to liquid first. • Dry Ice CO2 (s) --> CO2 (g) • Iodine I2 (s) --> I2 (g)

  44. Special Phase Changes • Deposition: From gas to solid without going to liquid state.

  45. Special Phase Changes

  46. HEAT OF FUSION FOR WATER (TABLE B) • Amount of heat needed to completely melt 1g of water (ice!). • 334 J/g • 334 Joules of heat are necessary to completely melt 1 g of water. • HOW MUCH HEAT IS NEEDED TO MELT 10 g OF WATER?

  47. HEAT OF VAPORIZATION FOR WATER (TABLE T) • The amount of heat needed to completely vaporize one g of water at its boiling point. • 2260 J/g • Water needs 2260 J of heat per gram to convert to gas!

  48. Summarizing When there is a change in temperature use Q = m x C x T While the substance is melting/freezing use Q = m x H f While the substance is boiling/condensing use Q = m x H vap

  49. Do now! • How much heat is needed to completely melt 10 g of ice at 0 0 C ? • How much heat is needed to vaporize 10 g of water at 100 0 C ?

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