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 temperature THERMOMETER

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

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 (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 NOT the same as HEAT ENERGY although 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. Heat always flows from warmer (HIGH T) to cooler(LOW T) objects.
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. a. Some things heat up or cool down faster than others.
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
MATERIAL
SPECIFIC HEAT J/g C
Water
4.18
Alcohol
2.43
Aluminum
0.90
Iron
0.45
Lead
0.13
Sand
0.83

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. Q = m x T x Cp Q = HEAT m = mass of substance
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) J
2) 6300J
3) 5 C
4) 20 C
5) 336 g
6) J =14.7 kJ
7) %
8) b
9)b J (d) 1 4 3 2

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 sublimation deposition temperature freezing
vaporization
fusion
solidification
gaseous phase
melt
heat
solid phase
heat of fusion
liquid phase
heat of vaporization
kinetic molecular theory

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 C ?