1. Chapter 3
Matter and Energy

2. CHAPTER OUTLINE Energy & Heat Temperature Scales Specific Heat
Classification of Matter
Physical & Chemical Properties
Physical & Chemical Changes
Conservation of Mass

3. ENERGY & HEAT Potential (stored) Kinetic (moving)
Energy is defined as the capacity of matter to do work.
Work is defined as the result of a force acting on a distance.
There are two types of energy:
Potential (stored)
Kinetic (moving)

4. ENERGY & HEAT
Energy possesses many forms (chemical, electrical, thermal, etc.), and can be converted from one form into another.
In chemistry, energy is commonly expressed as heat.
PE is converted to KE

5. ENERGY & HEAT Energy is conserved.
The law of conservation of energy states that energy is neither created nor destroyed.
The total amount of energy is constant.
Energy can be changed from one form to another.
Energy can be transferred from one object to another.

6. UNITS OF ENERGY
The SI unit of energy is the joule (J), named after the English scientist James Joule (1818–1889).

7. Heat & temperature are NOT the same thing!
HEAT vs. TEMPERATURE
Heat is measured in SI units of joule or the common unit of calorie.
Heat & temperature are NOT the same thing!
1 cal = J
Although the same amount of heat is added to both containers, the temperature increases more in the container with the smaller amount of water.

8. HEAT vs. TEMPERATURE The difference between Heat and Temperature
Heat is the total energy of all particles of matter
A form of energy associated with particles of matter
Temperature is the average kinetic energy of particles of matter
A measure of the intensity of heat or how hot or cold a substance is

9. TEMPERATURE
Temperature is a measure of how hot or cold a substance is.
Thermometer is an instrument used for measuring temperature, and is based on thermometric properties of matter (i.e. expansion of solids or liquids).
Three scales are used for measuring temperature.

10. TEMPERATURE SCALES Fahrenheit 32 - 212 Celsius 0 - 100 Kelvin

11. TEMPERATURE SCALES
To convert from one scale to another the following relationships can be used:
K = C + 273
F = (1.8 x C) + 32
C = (F - 32) ÷ 1.8
Alternately,
F = [(C + 40) x 1.8]-40
C = [(F + 40) ÷ 1.8]-40

12. Example 1:
The melting point of silver is C. What is this temperature in Kelvin?
TK = TC + 273
TK =
= 1234 K
= K

13. Example 2:
Pure iron melts at 1800 K. What is this temperature in Celsius?
TC = TK - 273
TC =
= 1527 C

14. Example 3:
On a winter day, the temperature is 5 F. What is this temperature on the Celsius scale?
TC = [(5 +40) ÷ 1.8]- 40
= -15 C

15. Example 4:
To make ice cream, rock salt is added to crushed ice to reach temperature of -11 C. What is this temperature in Fahrenheit?
TF = [( ) x 1.8]- 40
= 12 F

16. SPECIFIC HEAT
The specific heat of a substance is the amount of heat required to change the temperature of 1 g of that substance by 1C.
Units of specific heat are:
Different materials have different capacities for storing heat.
s = J / g ºC
s = cal / g ºC

17. Specific Heat of Some Substances
(cal/gC)
(J/gC)
Aluminum
0.214
0.897
Copper
0.0920
0.385
Iron
0.0308
0.129
Ammonia
0.488
2.04
Ethanol
0.588
2.46
Water
1.00
4.184
Most substances have substantially lower specific heats compared to water

18. SPECIFIC HEAT
When heated, substances with low specific heat get hot faster, while substances with high specific heat get hot at a slower rate.
When cooled, substances with low specific heat cool faster, while substances with high specific heat cool at a slower rate.

19. Specific heat of substance Change in its temperature
CALCULATING HEAT
The amount of heat lost or gained by a substance is related to three quantities:
Mass of substance
Specific heat of substance
Change in its temperature
Heat = x x
Q = m x s x T

20. Example 1:
How much heat is needed to raise the temperature of 200. g of water by 10.0 C. (Specific heat of water is J/gC)
Q = m x s x T
m = 200. g
s = J/gC
T = 10.0 C
Q = ???
Q = (200. g)(4.184 J/gºC)(10.0 ºC)
Q = J or kJ

21. Example 2:
Ethanol has a specific heat of 2.46 J/gC. When 655 J are added to a sample of ethanol, its temperature rises from 18.2 C to 32.8 C. What is the mass in grams of the ethanol sample?
Q = 655 J
s = 2.46 J/gC
T = 14.6 C
m = ???
m = 18.2 g

22. ENERGY & NUTRITION
In the laboratory, foods are burned in a calorimeter to determine their energy. A sample of food is burned in the calorimeter, and the energy released is absorbed by water surrounding the calorimeter.
The energy of the food can be calculated from the mass of the food and the temperature increase of the water.

23. Example 3:
A 2.3-g sample of butter is placed in a calorimeter containing 1900 g of water at a temperature of 17 C. After the complete combustion of the butter, the water has a temperature of 28 C. What is the energy value of butter in Cal/g?
1. Calculate heat absorbed by water
Heat absorbed by water
Heat released by butter =
2. Calculate energy value of butter

24. Example 3: Q = 21000 cal = 21 Cal Q = m x s x T
1. Calculate heat absorbed by water
m = 1900 g
s = 1.00 cal/gC
T = 11 C
Q = ???
Q = m x s x T
Q = (1900 g)(1.00 cal/gºC)(11 ºC)
Q = cal = 21 Cal
2. Calculate energy value of butter
21 Cal
2.3 g
= 9.1 Cal/g

25. ENERGY IN CHEMICAL CHANGES
Higher energy systems are less stable than lower energy systems.
In all chemical changes, matter either absorbs or releases energy.

26. ENERGY IN CHEMICAL CHANGES
When energy is released during a chemical change, it is said to be exothermic.
When energy is gained during a chemical change, it is said to be endothermic.
Exothermic reactions heat up
Endothermic reactions cool down

27. Which is exothermic and which is endothermic?
EXOTHERMIC vs. ENDOTHERMIC
energy is absorbed
higher potential energy
energy is given off
lower potential energy
Which is exothermic and which is endothermic?
4.4

28. CLASSIFICATION OF MATTER
Matter is anything that has mass, and occupies space.
Matter can be classified by its physical state as solid, liquid or gas.

29. SOLIDS
Solids are densely packed particles with definite shape and volume.
Solid particles have strong forces of attraction towards each other.
Solids are not very compressible.
Ice, diamond, quartz, and iron are examples of solid matter.

30. LIQUIDS
Liquids are loosely packed particles with definite volume but indefinite shape.
Liquid particles have moderate forces of attraction towards each other and are mobile.
Liquids are slightly compressible.
Water, gasoline, alcohol, and mercury are all examples of liquid matter.

31. GASES
Gases are very loosely packed particles with indefinite shape or volume.
Gas particles have little or no forces of attraction towards each other.
Gases are very compressible.
Oxygen, helium, and carbon dioxide are examples of gases.

32. GASES ARE COMPRESSIBLE
  Since the atoms or molecules that compose gases are not in contact with one another, gases can be compressed.

33. SUMMARY OF PROPERTIES OF MATTER

34. CLASSIFICATION OF MATTER
Anything that has mass
MIXTURE
Variable composition & properties
PURE SUBSTANCE
Fixed composition & properties
Mixtures can be converted into pure substances by simple physical processes (e.g. filtration, evaporation)

35. MIXTURES Also called solutions Tea, Coke Ink Salad dressing Cement
Variable composition & properties
HOMOGENEOUS
Uniform composition & properties
HETEROGENEOUS
Non-uniform composition & properties
Also called solutions
Tea, Coke Ink
Salad dressing Cement

36. PURE SUBSTANCES hydrogen, copper, gold
Fixed composition & properties
COMPOUNDS
2 or more elements chemically combined
ELEMENTS
Composed of one type of atom
Properties are unique compared to their components
Smallest particle is a molecule
hydrogen, copper, gold
Compounds can be converted into elements by chemical processes or reactions (e.g. electrolysis)
water, salt aspirin

37. separation of compound through chemical methods (electrolysis)
PURE SUBSTANCES
separation of compound through chemical methods (electrolysis)

38. CONCEPT CHECK Mixture: made of two or more types of substances
Classify each substance below as element, compound or mixture.
Mixture: made of two or more types of substances
Element: only one type of atom
Compound: composition is fixed
Element: only one type of atom

39. MIXTURES
Mixtures are 2 or more substances physically combined together.
Mixtures possess properties similar to those of their components.
Mixtures can be separated easily by a physical process.
Two types of mixtures are possible:
heterogeneous
homogeneous

40. HETEROGENEOUS MIXTURES
Heterogeneous mixtures are non-uniform in their composition.
Heterogeneous
Examples include vegetable soup, cement and salad dressing.

41. HOMOGENEOUS MIXTURES
Homogeneous mixtures are uniform in their composition.
Homogeneous mixtures are called solutions.
Homogeneous
Examples include gasoline, soda pop and salt solution.

42. MIXTURES vs. COMPOUNDS
List 3 differences between compounds & mixtures.
Composition
Compounds have fixed composition while mixtures have varied composition
Properties
Compounds have unique properties while mixtures have blended properties

43. MIXTURES vs. COMPOUNDS
List 3 differences between compounds & mixtures.
Make-up
Compounds are chemically combined (cannot be easily separated) while mixtures are physically combined (easily separated)

44. PHYSICAL & CHEMICAL PROPERTIES
The characteristics of a substance are called its properties.
Physical properties are those that describe the matter without changing its composition.
Examples are density, color, melting and boiling points, and electrical conductivity.

45. PHYSICAL & CHEMICAL PROPERTIES
The characteristics of a substance are called its properties.
Chemical properties are those that describe how matter behave in combination with other matter, and involve change in its composition.
Examples are flammability, corrosion, and reactivity with acids.

46. Examples:
Identify each of the following properties as physical or chemical:
Oxygen is a gas
Helium is un-reactive
Water has high specific heat
Gasoline is flammable
Sodium is soft & shiny
Physical
Chemical
Physical
Chemical
Physical

47. PHYSICAL CHANGES
Changes in physical properties of matter that do not involve change in its composition are called physical changes.
Examples are melting, evaporation and other phase changes.
Physical changes are easily reversible.

48. CHEMICAL CHANGES
Chemical changes are not easily reversible, and are commonly called chemical reactions.
A change that alters the chemical composition of matter, and forms new substance is called a chemical change.
Examples are burning, rusting, and reaction with acids.

49. Examples:
Identify each of the following changes as physical or chemical:
Cooking food
Mixing sugar in tea
Carving wood
Burning gas
Food molding
Chemical
Physical
Physical
Chemical
Chemical

50. Mass of Reactants Mass of Products =
CONSERVATION OF MASS
Similar to the law of conservation of energy, the law of the conservation of mass states that matter is neither created nor destroyed.
The total mass of substances does not change during a chemical reaction.
Mass of Reactants
Mass of Products =

51. → CONSERVATION OF MASS 266 g reactant 266 g product mass reactants =
Suppose that we burn 58 g of butane in a lighter. It will react with 208 g of oxygen to form ??? g of carbon dioxide and 90 g of water.
266 g reactant →
266 g product
mass reactants =
mass products
The number of substances and their properties may change, but the total amount of matter remains constant.

52. Burning: a Chemical Change
The butane molecules react with oxygen molecules in air to form new molecules, carbon dioxide and water.
This is a chemical change.

53. THE END