1. Chapter 3: Airbags

2. Airbags
This chapter will introduce the chemistry needed to understand how airbags work
Section 3.1: States of matter & Phase Diagrams
Section 3.2: Properties & Changes of matter
Section 3.3: Density
Section 3.4: Counting Molecules
Section 3.5: Gas Behavior & Gas Laws

3. Intro—Airbags

4. How do airbags work in your car?
Nylon bag inside your steering wheel
Solid sodium azide (NaN3) with is ignited with electricity when a crash sets off the trigger
2 NaN3 (s)  2 Na (s) + 3 N2 (g)
The nitrogen gas fills the airbag

5. Problems with this reaction?
NaN3 is very toxic
It produces sodium metal, which reacts with water to form hydrogen gas & enough heat to ignite that hydrogen gas
Reaction produces heat, so gas is very hot in airbag. An exothermic reaction.

6. The FIX! Addition of potassium nitrate!
It reacts with the sodium metal to form potassium oxide & sodium oxide and nitrogen gas
Following up with addition of silicon dioxide which reacts with the oxides to form silcates (glass).

7. Why do we use it?
It produces the gas very quickly,(at about 35 ms) – 4x faster than a blink of an eye
Reactants are small to store before needed
Amount of dangerous chemicals is minimal
Heat from reaction is absorbed, in part, by the physical components of the airbag system

8. Section 3.1

9. Solid Closely packed strong attractive forces
Vibrate in place; low kinetic energy of particles
Can’t switch places
Definite shape
Definite volume
Incompressible

10. Liquid Particles more spread out than solid Weaker attractions
Particles are free to move past each other: flow
Indefinite shape – takes shape of container
Definite volume
Slightly compressible

11. Gas Particles are very spread out; weakest attractive forces
Rapid, random motion; High kinetic energy of particles
Indefinite shape—take shape of container
Indefinite volume—they will fill container
Highly compressible

12. Changes in State: Endothermic
Kinetic energy must be put INTO the substance in order to increase the motions of the molecules thus breaking the intermolecular forces holding the particles together.
Melting: change of state from a solid to a liquid
Vaporization(Boiling or Evaporation): change of state from a liquid to a gas
Sublimation: change of state directly from a solid to a gas

13. Physical Changes: Names of the Phase Changes
Solid
Gas
Liquid

14. Changes in State: Exothermic
Kinetic energy must be taken OUT of the substance in order for the molecules to slow down so the intermolecular forces can begin draw & hold the particles closer together.
Freezing: change of state from a liquid to a solid
Condensation: change of state from a gas to a liquid
Deposition: change of state from a gas to a solid

15. Changes in State Endothermic Gas
Increasing molecular motion (temperature)
Changes in State
Gas
Sublimation
Boiling or
Evaporating
Endothermic
Liquid
Melting
Deposition
Condensing
Solid
Freezing
Exothermic

16. Temperature of state changes
Freezing point (fp) is the temperature at which a liquid turns into a solid
Melting Point (mp) is the temperature at which a solid turns into a liquid
freezing point is the same as melting point
Example: water has a melting point or freezing point of 0°C

17. Temperature of state changes
Boiling point (bp) is the temperature at which a liquid turns into a gas
Condensation Point (cp) is the temperature at which a gas turns into a liquid
boiling point is the same as condensation point
Example: water has a boiling point or dew point of 100°C

18. Physical Property
All substances have their own freezing and boiling points which make this property a great way to identify a unknown substance.

19. Atmospheric Pressure vs Vapor Pressure
Force per unit area created as gas molecules collide with objects
Force per unit area exerted against a surface by the weight of the air molecules above the surface
Force per unit area of the gas molecules above a liquid colliding
Usually measured in N/m2 but in chemistry we use atm or millimeters of mercury (mm Hg)
The more air molecules above a surface, the more molecules to exert a force and thus higher air pressure
The lower the attractive forces, the higher the vapor pressure
At sea level, atmospheric pressure equals 1 atm
Substance with high vapor pressure are called volatile

20. Temperature Controls Vapor Pressure
Only 2 factors control the Vapor Pressure of a Liquid: Nothing ELSE!!!!
Temperature
2. Attractive forces of the liquid

21. Vaporization: Differences between Evaporation and Boiling
Evaporation occurs
spontaneously at all temperatures
at the surface of the liquid
Boiling occurs
when extra kinetic energy(heat) is added.
at only 1 temperature dependent on pressure
It takes place within the body of the liquid

22. Real Definition of Boiling Point
external atmospheric pressure = vapor pressure of the liquid

23. Real Definition of Boiling Point
Since atmospheric pressure changes at various altitudes, “normal” boiling point is used to describe the temp at which a LG at “1 atm or 760 mmHg” of pressure
29, Mount Everest
69.0

24. Important Ideas The higher the altitude, the lower the atmospheric
pressure!
At higher altitudes, the boiling point is lower
It takes longer to cook foods at higher altitudes (lower atmospheric pressures)

25. Boiling Water By Changing Pressure

26. Heating & Cooling Curves
shows how solids, liquids & gases change state when temperature is changed
Plateaus : the changes of state (freezing, melting, boiling & condensation)
Freezing Point & Melting Point are at the same temperature or plateau
Boiling Point & Condensation Point are at the same temperature or plateau
Slopes= pure states (solid, liquid, gas)

27. At the plateaus, kinetic energy remains constant (temp
At the plateaus, kinetic energy remains constant (temp. remains constant) while potential energy changes
At the slopes, kinetic energy changes (temp. changes) while potential energy remains constant
*** DANGER!!***
Notice that a gas
can get higher
than boiling point!

28. Phase Changes BP 100 KE is changing/PE in constant  FP
 KE is constant/PE in changing

29. Heating Curve Examples: 1. What is the boiling point of the substance?
2. What letter represents the solid state only?
3. What letter represents the melting process?
100°C A B

30. Cooling Curve
While the substance is cooling during the liquid phase, the average kinetic energy of the molecules of the substance:
a) decreases
b) increases
c) remains the same 

31. Cooling Curve
Examples: 1. What is the freezing point of this substance?
70°C
2. How long does it take for the gas to completely liquefy?
9 -2= 7 min

32. Animation of Heating Curve

33. Phase Diagrams
shows how solids, liquids & gases change state as both temperature and pressure are changed
Crossing a line between states determines the change state (boiling, melting, etc)
A point directly on a line will identify the pressure and temperature (boiling point, melting point, etc.) of the change

34. Important Points on Phase Diagram
Triple Point is the temperature and pressure in which all 3 of the states coexist
Critical Point is the temperature & pressure at which a gas can no longer liquefy

36. Phase Diagrams of Water & Carbon Dioxide

37. Phase Diagram of Water 273 K
Temp. of line B at 1 atm (freezing point):
273 K
Temp. of line C at 1 atm (boiling point):
373K
D is the triple point

38. Phase Diagrams of Water
E is the critical point
What change of state happens when you cross line B at a constant pressure of 10 atm & increasing temp? melting
What change of state occurs when you cross line A at constant pressure of .001 atm? sublimation
What change of state happens when you cross line C at 400 K to 300K at approx. 5 atm?
condensation
Phase Diagrams of Water

39. Section 3.2
What are the properties & changes occurring within the airbag?

40. Physical versus Chemical Properties
Physical Property
Chemical Property
Characteristic that can be determined or measured without changing the substance’s identity
Characteristic that can only be determined or measured as the substance changes into different substances

41. Examples of Physical Properties
COLOR ODOR
TEXTURE BOILING POINT
DENSITY SOLUBILITY
VOLUME MASS
HARDNESS
MALLEABILITY /BRITTLENESS

42. Physical Properties

43. Examples of Chemical Properties
BURNING/COMBUSTING
RUSTING
ROTTING
FLAMMABILITY
REACTIVITY
NEUTRALIZATION
DECOMPOSING

44. Intensive and Extensive Properties
Intensive Property
Extensive Property
Size of the sample DOES NOT matter— a big piece & a small piece are the same with respect to the property
Size of the sample DOES matter
—a big piece & a small piece would be different with respect to the property

45. Intensive Extensive Melting point/boiling point Mass Volume Density
Color/Smell
Conductivity
Hardness
Extensive
Mass
Volume
Energy
Length
Shape

46. SELF CHECK Flammability Boiling point Solubility Malleability
Reactivity with oxygen

47. Flammability
Boiling point
Solubility
Malleability
Reactivity with oxygen
Chemical
Physical, Intensive

49. Physical VS Chemical Changes
Physical Change: the chemical structure of the substances is not changed.
H2O(l) H2O(g)

50. Physical VS Chemical Changes
Chemical Change: the chemical structures of the substances are changed.
H2O(l) H2(g) + O2(g)
RECALL: Another name for a chemical change is called a CHEMICAL REACTION.

51. Physical & Chemical Changes
Physical changes
do not produce new
substances
breaking, dissolving, distilling, cutting,
Changes in state (boiling, condensing, melting & freezing)
Chemical changes
do produce new
substances
rusting, burning, metabolizing food, oxidation or reduction, reacting with oxygen, etc.

53. Physical & Chemical Changes
Also…if a change can be un-done by a physical change, then the original change was physical as well.
If salt is dissolved in water, it seems to disappear…
many people think this is a chemical change.
But if the water is evaporated (a physical change), the salt is left in the container.
Since the original change was un-done with a physical change, then the original change (the dissolving) was a physical change as well.

54. Confusing changes People often use the following terms incorrectly.
Definition
Type of Change
Melting
Changes a solid into a liquid
Physical
Burning
Reacting with oxygen to produce CO2 and H2O
Chemical
Dissolving
Adding one substance to another to form a homogeneous mixture
Physical
Drying
Heating a sample to evaporate the water
Physical

55. RECALL the Possible Signs of Chemical Changes
Gas production (bubbling)
Energy change (getting hot or cold)
Color change
Light given off
Formation of a precipitate (an insoluble substance formed from two soluble substances)

56. They’re “Possible” signs
Sometimes these “signs” accompany physical changes as well!
Gas production (bubbling). Bubbles are formed during boiling (a physical change)
Energy change (getting hot or cold). Energy changes accompany changes in state (physical changes)
Color change. Color change can occur due to dissolving a substance (a physical change)
However, some of these signs also accompany physical changes, so you must take into account many observations to determine if the change was in fact chemical.

57. Section 3.3
Do you want high or low density in your airbag?

58. Density the ratio of mass to volume of a sample
How heavy is it for its size?
Lead = high density…small size is very heavy
Air = low density…large sample has very little mass

59. Density Mass Density m D = V Volume
In grams (g)
Density
In g/L or g/mL
D = m V
Volume
In liters (L) or mL
Don’t try to cancel out the units…density has “2 units” – a mass unit over a volume unit!

60. Floating
Objects float when they are less dense than the substance they are in!
Is vegatable oil more or less dense than water?
Fewer particles in the same space = less dense
More particles in the same space = More dense

61. Density Values: the larger the value, the more dense

62. Density Varies with Temperature
WHY?
Most substances will expand when heated, increasing the volume & decreasing the density.
Water is an exception: As water is cooled, it expands, increasing the volume & decreasing the density.
Thus, ICE is less dense than WATER!

63. Calculating Volume using Water Displacement
The volume is the difference between the final volume and the initial volume of water.
What is the volume of the dinosaur? ______________

64. Example 1—Solving for Density
What is the density of a sample with a mass of 2.50 g and a volume of 1.7 mL?

65. Example 1—Solving for Density
m = 2.50 g
V = 1.7 mL
Example:
What is the density of a sample with a mass of 2.50 g and a volume of 1.7 mL?

66. Example 2—Solving for Mass
What is the mass of a 2.34 mL sample with a density of 2.78 g/mL?

67. Example 2—Solving for Mass
V = 2.34 mL
D = 2.78 g/mL
Example:
What is the mass of a 2.34 mL sample with a density of 2.78 g/mL?
2.34 mL ×
× 2.34 mL

68. Example 3—Solving for Volume
A sample is 45.4 g and has a density of 0.87 g/mL. What is the volume?

69. Example 3—Solving for Volume
m = 45.4 g
D = 0.87 g/mL
V = ?
Example:
A sample is 45.4 g and has a density of 0.87 g/mL. What is the volume?
V ×
× V
0.87 g/mL
0.87 g/mL

70. Is it aluminum? The metal has a mass of 612 g and a volume of 345 cm3.
SELF CHECK
Example:
Is it aluminum? The metal has a mass of 612 g and a volume of 345 cm3.
The accepted density of aluminum is 2.70 g/cm3

71. Graphing Density Density
Volume (mL)
Mass (g)
Density
If we make the y-axis mass and the x-axis volume then…
Then the slope equals Density!