1. Matter and Energy

2. Chapter Outline Classification of Matters Physical/Chemical Properties
Physical/Chemical Changes
Energy
Temperature, Heat, and Specific Heat

3. In Your Room
Everything you can see, touch, smell or taste in your room is made of matter.
Chemists study the differences in matter and how that relates to the structure of matter.

4. What is Matter? Matter: anything that occupies space and has mass
Matter is actually composed of a lot of tiny little pieces: Atoms and Molecules

5. Atoms & Molecules Atom : Smallest piece of an element
Atom consists of subatomic particles
Movie of atoms:
Molecule : Smallest piece of a compound
made of atoms (electron density map)
Same molecules makes up a compound

6. Physical States of Matters
Matter can be classified as solid, liquid or gas based on what properties it exhibits

7. Why different States of a Matter? Structure Determines Properties
the atoms or molecules have different structures in solids, liquid and gases

8. Solids
Particles in a solid: packed close together and are fixed in position
though they may vibrate
___Compressibility
retaining their shape and volume
_______ to flow

9. Liquids
Particles are closely packed, but they have some ability to move around
 _____ compressibility
______ to flow, yet not to escape and expand to fill the container (not “antigravity”)

10. Gases
The particles have complete freedom from each other (not sticky to each other)
The particles are constantly flying around, bumping into each other and the container
There is a lot of empty space between the particles (low density)
 Compressible
 Able to flow and Fill space (“antigravity”)

11. Classifying Matter:
Copper, Sugar, Coke, Gasoline/Water

12. Pure substance vs. Mixture
Constant Composition
Homogeneous
Mixture
Variable Composition
Matter
Pure substance (A, B, C) is composed of only one kind of piece, like Diamond, 24K Gold, pure Sugar…
Mixture (D) is composed of different kinds of pieces, like Brass, Flour, Soda drink

13. Pure substance: Element vs. Compound
Elements: Substances which can not be broken down into simpler substances by chemical reactions. (A,B)
Compounds: Most substances are chemical combinations of elements. (C)
Examples: Pure sugar, pure water
can be broken down into elements
Properties of the compound not related to the properties of the elements that compose it

14. Elements 116 known, 91 are found in nature
others are man-made
Abundance = percentage found in nature
Hydrogen: most abundant in the universe (fuel for starlight, such as in the Sun)
Oxygen: most abundant element (by mass) on earth and in the human body
Silicon: abundant on earth surface
every sample of an element is made up of lots of identical atoms

16. Compounds Composed of elements in fixed percentages
water is 89% O & 11% H
billions of known compounds
Organic (sugar, glycerol) or inorganic (table salt)
same elements can form more than one different compound
water and hydrogen peroxide contain just hydrogen and oxygen
carbohydrates all contain just C, H & O (sugar, starch, glucose)

17. Mixture
Matter that is composed of different kinds of pieces. Different samples may have the same pieces in different percentages. (D)
Examples:
Solid: Flour, Brass (Copper and Zinc), Rock
Liquid: Salt water, soda, Gasoline
Gas: air

18. Classification of Mixtures
Homogeneous mixture is uniform throughout
appears to be one thing
every part has identical properties, though another sample with the same components may have different properties
solutions (homogeneous mixtures): Air; Tap water
Heterogeneous mixture is non-uniform throughout
contains regions with different properties than other regions: gasoline mixed with water; Italian salad dressing

19. Pure Substances vs. Mixtures
All samples have the same physical and chemical properties
Constant composition = all samples have the same pieces in the same percentages
Homogeneous
separation into components based on chemical properties
Temperature usually stays constant while melting or boiling
Mixtures
different samples may show different properties
variable composition = samples made with the same pure substances may have different percentages
homogeneous or heterogeneous
separate into components based on physical properties
temperature changes while melting or boiling because composition changes

20. Classifying Matter in Daily Life
Pure element or compound?  Mixture: Homo-/Heterogeneous?
Pb, pure element
Mixture, homogenous
H2O, pure compound
Water + vinegar + oil, heterogeneous
Ag, pure
Pure substance or Mixture:
Lead fishing weight
Tap water
Distilled water
Italian salad dressing
Silver dollar

21. How Matters Differ from each other?
Physical Properties are the characteristics of matter that can be changed without changing its composition
characteristics that are directly observable
Chemical Properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy
characteristics that describe the behavior of matter

22. Some Physical Properties

23. Some Chemical Properties

24. Iron: Physical Properties vs. Chemical Properties
easily corroded in moist air to form brownish rust
when added to hydrochloric acid, it disappears and generates hydrogen gas
more reactive than silver, but less reactive than magnesium
silvery solid at room temperature with a metallic taste and smooth texture
melts at 1538°C and boils at 4428°C
Density g/cm3
Can be magnetized
conducts electricity
thermal conductivity

25. Physical Changes in Matter
Changes in the properties of matter that do NOT effect its composition
Heating water
raises its temperature, but it is still water
Evaporating butane from a lighter
Dissolving sugar in water
even though the sugar seems to disappear, it can easily be separated back into sugar and water by evaporation

26. Chemical Changes in Matter
Definition: A change in the properties of matter that change its composition, involving a Chemical Reaction
Rusting of Iron: iron combining with oxygen to make iron(III) oxide
Burning Butane from a lighter: butane gas  carbon dioxide and water
Silverware tarnishing: Silver combines with Sulfur in the air to make tarnish

27. Is it a Physical or Chemical Change?
a physical change results in a different form of the same substance
the kinds of molecules don’t change
a chemical change results in one or more completely new substances
the new substances have different molecules than the original substances
you will observe different physical properties because the new substances have their own physical properties

28. Phase Changes: Physical Changes
Boiling = liquid to gas
Melting = solid to liquid
Subliming = solid to gas
Condensing = gas to liquid
Freezing = liquid to solid
Deposition = gas to solid
state changes require heating or cooling the substance
evaporation is not a simple phase change, it is a solution process

29. Separation of Mixtures
Separate mixtures based on different physical properties of the components
Physical change
Centrifugation &
Decanting
Density
Evaporation
Volatility
Chromatography
Adherence to a Surface
Filtration
State of Matter (solid/liquid/gas)
Distillation
Boiling Point
Technique
Different Physical Property

30. Filtration of Heterogeneous Mixture: filtering spaghetti from boiling water
Solid is “captured” by filter paper
Liquid flows through the filter paper

31. Distillation of Homogeneous Mixture
Separation of liquids through “boil/condense”
More volatile component
less volatile component
Heat makes liquid boil

32. Law of Conservation of Mass
Antoine Lavoisier: “Father of Modern Chemistry”: “Matter is neither created nor destroyed in a chemical reaction”
Massbefore = Massafter
Massreactants = Massproducts

33. Conservation of Mass 266 grams = 266 grams
Total amount of matter remains constant in a chemical reaction
58 grams of butane burns in 208 grams of oxygen to form 176 grams of carbon dioxide and 90 grams of water.
butane oxygen  carbon dioxide + water
58 grams grams  grams grams
266 grams = grams

34. Energy Energy do not have mass and volume (<-> Matter)
Energy is anything that has the capacity to do work
Chemistry is the study of matter, Chemical process involve energy change
it can cause physical and/or chemical changes in matter

35. Matter Possesses Energy
when a piece of matter possesses energy, it can give some or all of it to another object
all chemical and physical changes result in the matter changing energy

36. Kinds of Energy Kinetic and Potential
Kinetic Energy is energy of motion, or energy that is being transferred from one object to another (running water)
Examples: Heat
Potential Energy is energy that is stored (water above the dam)
Examples: Chemical Energy, Nuclear energy, Attraction between opposite charges.

38. Law of Conservation of Energy
“Energy can neither be created nor destroyed”
the total amount of energy in the universe is constant – there is no process that can increase or decrease that amount
We can transfer energy from one place to another: Solar energy from the Sun to the Earth
We can transform the energy:
Solar panel: Solar energy  ______
Gas engine: Chem energy  _______
Applying the brake: Kinetic energy  _______

39. Units of Energy
calorie (cal) is the amount of energy needed to raise one gram of water by 1°C
kcal = energy needed to raise 1000 g of water 1°C
food Calories = kcals
Energy Conversion Factors
1 calorie (cal) =
4.184 joules (J)
1 Calorie (Cal)
1000 calories (cal)
1 kilowatt-hour (kWh)
3.60 x 106 joules (J)

40. Energy Use Unit joule (J) 4.18 3.6 x 105 9.0 x 108 calorie (cal) 1
Energy Required to Raise Temperature of 1 g of Water by 1°C °C
Energy Required to Light 100-W Bulb for 1 hr
Energy Used by Average U.S. Citizen in 1 day
joule (J)
4.18
3.6 x 105
9.0 x 108
calorie (cal) 1
8.6 x 104
2.15 x 108
Calorie (Cal)
0.001 86
215,000
kWh
1.1 x 10-6
0.10
250

41. Conversion Problem: 955 Cal  Joule?
Information
Given: 955 Cal
Find: ? J
Conv. Fact.
1 cal = Joules
3.44 x 103 J

42. What is Heat
Heat is the exchange of thermal energy between samples of matter
How Heat flows?
Heat transfers from matter that has _______ temperature to matter that has ______ temperature
until they reach the same temperature
heat is exchanged through molecular collisions between two samples

43. Exothermic/Endothermic Processes
Physical/Chemical changes are often accompanied by energy change
Exothermic: A change _____________ heat.
Example: Combustion (burning); Condensation of water
Endothermic: A change ________ heat.
Example: Ice melting; Evaporation of alcohol

44. The Meaning of Temperature
Temperature is a measure of the average kinetic energy of the molecules in a sample
Not all molecules have in a sample the same amount of kinetic energy
a higher temperature means a ______ average kinetic energy

45. Fahrenheit Temperature Scale
Two reference points:
Freezing point of concentrated saltwater (0°F)
Average body temperature (100°F)
more accurate measure now set average body temperature at 98.6°F
Room temperature is about 75°F

46. Celsius Temperature Scale
Two reference points:
Freezing point of distilled water (0°C)
Boiling point of distilled water (100°C)
more reproducible standards
most commonly used in science
Room temperature is about 25°C

47. Fahrenheit vs. Celsius
a Celsius degree is 1.8 times larger than a Fahrenheit degree
Conversion between Fahrenheit and Celsius

48. The Kelvin Temperature Scale
Kelvin scale is an absolute scale, meaning it measures the actual temperature of an object
0 K = Absolute Zero: all molecular motion would stop
0 K is theoretically the lowest temperature in the universe
0 K = -273°C = -459°F
Absolute Zero is a theoretical value

49. Kelvin vs. Celsius
the size of a “degree” on the Kelvin scale is the same as on the Celsius scale
though technically, call the divisions on the Kelvin as kelvins, not degrees
that makes 1 K 1.8 times larger than 1°F
the 0 standard on the Kelvin scale is a much lower temperature than on the Celsius scale

50. Extremes of Temperature
On the Earth,
Lowest temperature recorded: -89.2°C ( °F, 184 K)
Highest air temperature recorded: ~60°C
In science lab,
the highest temperature: 4 x 1012 K (?)
the lowest temperature: ~10-10 K (?)

51. Temperature Scales Celsius Kelvin Fahrenheit Rankine 100°C 373 K 212°F
BP Water
298 K
75°F
534 R
Room Temp
25°C
0°C
273 K
32°F
459 R
MP Ice
-38.9°C
234.1 K
-38°F
421 R
MP Mercury
-183°C
90 K
-297°F
162 R
BP Oxygen
BP Helium
-269°C
4 K
-452°F
7 R
-273°C
0 K
-459 °F
0 R
Absolute
Zero
Celsius
Kelvin
Fahrenheit
Rankine

52. Example: Convert -80 °F into Celsius and Kelvin
°C = -62 °C (round to 2 significant figures)
K = 211 K

53. Example: Convert 80 K into Celsius and Fahrenheit
°C = -193°C , °F = -315 °F

54. Heat and Temperature
The temperature increase of an object (DT) depends on the amount of heat energy added (q).
When heating a pot of water, the more heat is given, the __________ (higher / lower) temperature increase (aka larger DT)
For the same amount of heat q, the temperature increase of an object DT depends on its mass (m).
Using the same heat, ______ (small / large) pot of water will have higher temperature change (aka, larger DT).

55. Heat and Temperature
Given the same heat, the temperature increase of an object (DT) depends on what the material it is made of.
Same amount of heat upon 1 kg of gold vs. 1 kg of aluminum: gold will be heated to higher temperature.
What makes different materials differ from each other is the Specific Heat.

56. Specific Heat: A Physical Property
Specific heat c = heat absorbed by 1 gram of the substance to raise the temperature by 1 °C
c = q / (m DT)
cal/g°C or J/g°C
waters specific heat = J/g°C for liquid
or cal/g°C
less for ice and steam

57. About Specific Heat Specific heat is a property of the type of matter
it can be used to identify the type of matter
Higher specific heat, more heat needed to increase temperature
water’s high specific heat  a good cooling agent
it absorbs a lot of heat for a relatively small mass
Heavier atoms tend to have lower specific heat.

58. What difference does Specific Heat make?
Diamond and graphite are both made from pure carbon. From 0C to 100C, which object absorbs more heat, 1.0 kg diamond or 1.0 kg graphite?
Given 100. J heat, which object will have higher temperature change, 1.0 kg silver or 1.0 kg aluminum?

59. Heat Gain or Loss by an Object
Heat energy gained or lost by an object depends on
how much material there is
what the material is, and
how much the temperature changed
The sign of temperature change and calculated heat
If T > 0 (temperature increase), q > 0 is for heat absorbed (gained)
If T < 0, q < 0 is for heat released (lost)

60. A. Calculating Heat: During a shower, 5
A. Calculating Heat: During a shower, 5.0 gallons of water is heated from 11.0°C to 49.0°C. The specific heat of water is J/g°C. How much heat is needed? 1 gal = 3.78 L.
Given: V = 5.02 gal;
Ti = 11.0°C;
Tf = 49.0°C;
C = J/g°C
Q = ?
Mass = 1.90 x 104 g
q = 3.02 x 106 J (keep 3 SF)

61. A2. Calculating Heat: Refrigerator absorbs heat from food
A2. Calculating Heat: Refrigerator absorbs heat from food. How much heat will a cup of water (234 g) lose so that its temperature falls from 25.0°C to 4.0°C. The specific heat of water is J/g°C.
Given: mass = 234 g;
Ti = 25.0°C;
Tf = 4.0°C;
C = J/g°C
q = ?
Heat q = x 104 J (keep 3 SF)

62. B. Finding Specific Heat: A sample of 11 g liquid absorbs 56 J heat
B. Finding Specific Heat: A sample of 11 g liquid absorbs 56 J heat. The temperature of liquid rose from 10.4°C to 12.7°C. what is the specific heat of this liquid?
Information
Given: m = 11 g; Ti = 10.4°C;
Tf = 12.7°C; q = 56 J
Find: c
Eq’n: q = m ∙ C ∙ DT
Solution Map: q, m, DT  c
2.2 J/g °C (keep 2 SF)

63. C. Find the temperature change: 57 g Lead absorbs 611 J heat
C. Find the temperature change: 57 g Lead absorbs 611 J heat. Initial Temperature 47°C. The heat capacity of Lead C = J/g°C. What is the final temperature?
Given: m = 57 g; Ti = 47°C;
q = 611 J;
C = J/g°C
Find: Tf = ?°C
Eq’n: q = m ∙ C ∙ DT
Solution map: q, C, m DT Tf
DT = Tf - Ti
Tf = 131 °C