Honors Chemistry Matter and Energy ppt

1. The Physical
Behavior of
Matter

2. What is Matter?
Anything that takes up space & has mass
 Substances – variety
of matter that has
the same
composition and
properties
throughout.
 Two types are
elements (Na) and
compounds (NaCl).
 Mixtures – two or
more substances
mixed together (not
united).
 Two types are
homogeneous
(salt water) and
heterogeneous
(sand and sugar).

3. Substances
Elements
 Cannot be
decomposed
(broken down)
 Ex: hydrogen (H2),
Oxygen (O2) and
nitrogen (N2)
Compounds
 Can be decomposed
by a chemical change
 Chemically united
 Definite proportions
 Different properties
 Ex: H2O, NaCl, CO2
 A Binary Compound is a
compound that only has
two elements (NaCl)

4. Differences between Mixtures
Homogeneous
 Solutions that are
considered one thing.
 Example: White vinegar
Heterogeneous
 Uneven mixture of two
different things.
 Example: milk

5. Ways To Separate Mixtures (1)
 Distillation- a mixture of liquids can be
separated by their boiling points.
 Examples: distillation of salt water.
distillation of petroleum liquids.

6. Ways to Separate Mixtures (2)
• Filtration- separates the solid and liquid
parts of a mixture.
• Example: coffee filter which separates the
coffee grounds from the brewed coffee.

7. Ways of Separating Mixtures (3)
 Chromatography- way of separating different
molecules in a mixture.
 Example: separating components of chlorophyll.

8. Ways of Separating Mixtures (4)
 Centrifuge- a spinning machine that
pushes the most dense particles to the
bottom of the tube.
 Example: separate isotopes such as separating
uranium hexafluoride, and uranium-235.

9. Ways to Separate Mixtures (5)

10. Properties
Physical Properties
 Can be found without
changing the substance to
something else.
 Ex. Color, hardness, phase,
solubility, odder, density,
mass, volume
Chemical Properties
 Are found by making a
substance react, and form a
new substance.
 Ex. Burning, reaction w/ water
or acid, changing to a new
substance.

11. Energy
 The ability to do work.
 Exothermic- energy given off in a
chemical reaction.
 Endothermic- energy absorbed in a
chemical reaction.
 You measure energy in joules (J).

12. Table T
Heat
q= mCΔT q=heat
m= mass
q= mHf C= specific heat capacity(table B)
ΔT= change in temperature
q= mHv Hf= heat of fusion
Hv= heat of vaporization

13. Sample Problem:
How much heat energy in joules if
absorbed by 100g of water when it is
heated from 20ºC to 30ºC?
q= m· C · ΔT
q= 100g x 4.18 Joules/gºC x (30 – 20) ºC
q= 100g x 4.18 J/gºC x 10ºC
q= 4,180 joules

14. Heating Curves
Heat of Fusion- is the amount of heat needed to
change a solid into a liquid at a constant
temperature.
Heat of Vaporization- is the amount of heat
needed to change a liquid into a gas at a constant
temperature.
Heat of Fusion 334 J/g
Heat of Vaporization 2260 J/g
Specific heat capacity of H2O(l) 4.18 J/gºC
Table B
Physical Constants for Water

15. Temperature
 The measure of the average kinetic energy
of the molecules.
 The higher the temperature the more kinetic
energy it has.
 Heat flows from a higher temperature to a
lower temperature until they are the same
temperature.
 Measured with a thermometer.

16. Temperature continued
 Boiling Point- when the vapor pressure equals
the atmospheric pressure.
 Freezing Point- the temperature at which a
liquid solidifies under a specified pressure.
 Absolute Zero= -273oC or 0K
 Kelvin-
K= Kelvin ºC = degrees Celsius
K = oC + 273

17. Changes
Physical Change
-change in
appearance, but no
new substance is
produced
Ex- tearing a piece of
paper, heating ice
Chemical Change
- Produces a new
substance with
different properties
Ex- burning
magnesium

18. Solids
 Definite shape, definite volume, and crystalline
structure, geometric pattern.
 Closely packed particles that vibrate but don’t
change position.
 Melting point- temperature when a solid changes
into a liquid.
 Sublimation- the change from a solid directly to a
gas. Ex. Dry ice (CO2) & Iodine

19. Liquids
 Definite volume, takes shape of its container.
Particles are close together and move :water.
 Evaporation- when liquid changes into a gas.
Ex: water vapor.
 Vapor Pressure- the pressure that the vapor
exerts on the sides of the container.

21. Gases
 No definite volume or shape.
 Particles are far apart, and can expand
anywhere.
* When there
is a phase
change from
solid to liquid
to gas entropy
increases

22. Gas Laws
 Boyle’s Law
As pressure increases, volume decreases at constant
temperature

23. Gas Laws
 Charles’ Law
As volume increases, temperature increases at
constant pressure

24. What is STP?
 Standard Temperature and Pressure
 When you have a combined gas law at STP, use 273
K as your temperature and 101.3 kPa as the pressure.

25. Kinetic Molecular Theory (Ideal Gas Law)
A model that tells how gasses should behave
• Ideal gas- perfect gas that agrees with John Daltons 5
assumptions
• Tiny particles
• Elastic collisions
• Gases are in constant motion
• No force of attraction
• Temperature is related to speed

26. Ideal Gas
• Particles have no
volume.
• No attractive forces
• Examples: H2, He
Real Gas
• Particles have
volume
• Attractive
• Examples: Cl2,
H2O(g)

27. How do the Gas Laws relate to
the Kinetic Molecular Theory?
Boyle’s Law
• Boyle’s law states that
pressure and volume are
inversely proportional.
• If you have a million
molecules in a container and
you decrease that container
the molecules will hit twice as
often, therefore twice the
pressure.
Charles’ Law
• Charles’ Law states that as
temperature increases,
volume increases.
• If you heat the air in a
balloon, there will be more
pressure on the sides. This
makes the balloon bigger in
volume.

28. Combined Gas Law
P1
x V1 = P2
x V2
T1 T2

29. THE END