1. Unit 6: Physical Behavior of matter

2. Definite Composition (Homogeneous)
I. Classification of Matter Chart
Matter
Substance
Definite Composition (Homogeneous)
Physically
Separable
Mixture of Substances

3. I. Classification of Matter Chart
Substance
Element
Contains just a single element
Ex: Iron (Fe), Oxygen (O2)
Compound
Two or more different elements
Ex: Water (H2O), Carbon Dioxide (CO2)
Chemically
Separable

4. Checks for Understanding
A compound differs from an element in that a compound
Is homogeneous
Has a definite composition
Has a definite melting point
Can be decomposed by a chemical reaction
Which of the following substances cannot be separated by chemical change?
Nitrogen (g)
Sodium chloride (s)
Carbon dioxide (g)
Magnesium Sulfate (aq)

5. Nonuniform; distinct phases
I. Classification of Matter Chart
Mixture of Substances
Homogeneous
Uniform throughout (air, salt water, aqueous solutions (aq))
Heterogeneous
Nonuniform; distinct phases
(soup, concrete)

6. Check for Understanding
A pure substance that is composed only of identical atoms is classified as…
A compound
An element
A heterogeneous mixture
A homogeneous mixture
A heterogeneous material may be…
A mixture
Pure substance

7. II. Particle Diagrams

8. III. Separating Matter
Certain types of matter can be separated using various methods.
Monatomic Elements - _____________ be decomposed (broken apart) using _____________ or ______________ means.
Diatomic Elements and Compounds (ie – O2 and H2O) – can be decomposed using __________________ only
CANNOT
CHEMICAL
PHYSICAL
CHEMICAL MEANS

9. III. Separating Matter
Mixtures – can be separated using ___________________
Filtration –
Evaporation –
Chromatography –
Distillation –
PHYSICAL MEANS
Separation by particle size
Separation by boiling point
Separation by polarity
Separation by boiling point

10. Check for Understanding
Which of the substances could be decomposed by a chemical change?
A) sodium B) aluminum C) magnesium D) ammonia
A sample of a material is passed through a filter paper. A white deposit remains on the paper, and a clear liquid passes through. The clear liquid is then evaporated, leaving a white residue. What can you determine about the nature of the sample?
What are some of the differences between a mixture of iron and oxygen and compound composed of iron and oxygen?
It is a heterogeneous mixture
In a mixture the elements are not bonded with each other and can be physically separated. In a compound the elements are bonded and can only be separated through a chemical reaction.

11. Think about this
What would a particle diagram look like for each of these phases?
What happens to the spacing and speed of particles at each of the phases?
SOLID LIQUID GAS

12. IV. Forms of Mechanical Energy
Kinetic Energy
Energy of movement
(similar to temperature)
(how fast atoms are moving)
Potential Energy
Stored energy
(energy of position)
More spread out (gas) = High PE
Closer together (solid) = Low PE

13. V. Heating and Cooling Curves (animation)
ENDOTHERMIC
ABSORBED
Heating Curve: ___________ - Energy is being ________
 gas
l  g
 liquid
s  l
s  g
 solid
Sublimation (video)-
Solid changes directly to a gas
Heating Curve Animation

14. V. Heating and Cooling Curves
AB
BC
CD
DE
EF
Kinetic Energy
Potential Energy
Phase
Con-stant
Con-stant ↑ ↑ ↑
Con-stant
Con-stant
Con-stant ↑ ↑
gas
solid
l  g
boiling
s  l
melting
liquid

15. Check for Understanding
A substance begins to a melt. What happens to the potential and kinetic energy?
PE increase, KE stays the same
2. The temperature of a substance refers to what type of energy?
Kinetic energy
3. How does the speed and space of water molecules compare when in a liquid phase to a gas phase
Molecules move faster and more spread out in gas phase

16. V. Heating and Cooling Curves
EXOTHERMIC
RELEASED
Cooling Curve: ___________ - Energy is being ________
 gas
g  s
g  l
 liquid
l  s
 solid
Deposition -
Gas changes directly to a solid

17. V. Heating and Cooling Curves
AB
BC
CD
DE
EF
Kinetic Energy
Potential Energy
Phase
Con-stant
Con-stant ↓ ↓ ↓
Con-stant
Con-stant
Con-stant ↓ ↓
g  l
condensing
solid
gas
liquid
l  s
Freezing

18. Check for Understanding
As a substance condenses, what happens to its potential and kinetic energy?
PE decreases, KE stays the same
2. What phase is a substance in when it has its highest kinetic energy?
gas
3. How does the speed and space of water molecules compare when in a liquid phase to a solid phase
Molecules move slower and are closer together in solid phase

19. Vi. Temperature vs. Heat
Amount of energy transferred from one substance to another
Average kinetic energy of its particles (how fast they’re moving)
Joules (J) or Calories (cal)
1 cal = 4.18 J
Celsius (oC) or Kelvin (K)
(K = oC + 273) T q

20. Vi. Temperature vs. Heat K = oC + 273 K = Kelvin oC = degrees Celsius
Temperature Scales (See Ref. Tabs.):
K = oC K = Kelvin oC = degrees Celsius
Convert: 200 degrees Celsius to Kelvin
Law of Conservation of Energy:
Heat Transfer:
K = oC + 273
K = 200oC = 473 K
Energy (heat) cannot be created or destroyed. Energy (heat) can be TRANSFERRED.
HEAT ALWAYS MOVES FROM WARMER OBJECTS TO COLDER OBJECTS

21. Check for Understanding
You wake up in the morning and your barefoot touches the ceramic floor and it feels cold. Explain which way heat is being transferred.
Heat moves from your body (warm) to the floor (cold)
You are cooking pasta in a boiling metal pot of water. You grab the metal handles with your bare hands (ouch!). Explain which way heat is being transferred.
Why do you feel cold after you get out of a hot shower. (link)
Heat moves from metal handles (warm) to your hands (cold).

22. vii. Endothermic and exothermic (revisited)
Energy is either absorbed or released in chemical reactions
Remember:
- Breaking bonds is ____________________
- Heat is ___________ the reaction from the surroundings
- Ex) heat + Br2  Br + Br
- Creating bonds is _________________
- Heat is ___________ the reaction into the surroundings
- Ex) N+ N N2+ energy
ENDOTHERMIC
ENTERING
EXOTHERMIC
EXITING

23. vii. Endothermic and exothermic (revisited)
Where does the heat come from (or go to)? ______________________
For exothermic reactions, heat (energy) leaves the reaction and moves __________________________. Therefore, making the surrounding temperature _________________
Exothermic Chemical Equation:
The surroundings
into the surroundings
warmer
Reactant(s)  Product(s) + HEAT

24. vii. Endothermic and exothermic (revisited)
For endothermic reactions, heat (energy) leaves the surroundings and moves __________________________. Therefore, making the surrounding temperature _________________
Endothermic Chemical Equation:
Into the reaction
colder
Reactant(s) + HEAT  Product(s)
ENDOTHERMIC REACTION VIDEO

25. vii. Endothermic and exothermic (revisited)
Enthalpy: It is a measure of the heat released or absorbed in a reaction. If heat is released then the reaction is an exothermic reaction, heat will be on the products side and the ΔH will be negative. If heat is absorbed then the reaction is an endothermic reaction, heat will appear on the reactants side and the ΔH will be positive. (see Table ____ on Reference Tables).
Examples: Tell whether each of the following reactions are endothermic or exothermic (you may have to use table I). Then tell whether the temperature of the surroundings increases or decreases as a result.
Endo/Exothermic _______Surroundings
__________ C6H12O O2  6CO H2O + heat _______________ I
Exo (-∆H)
warmer

26. vii. Endothermic and exothermic (revisited)
Endo/Exothermic _______Surroundings
____________C6H12O6 (s) + 6O2(g)  6CO2 (g) + 6H2O + heat ____________
2. __________ 2H2O kJ  2H2 + O2 _______________
6. __________ 2KClO3(s)  2KCl(s) + O2(g) + heat _______________
7. __________ H+(aq) + C2H3O2(aq) + heat  HC2H3O2(l) _______________
exo(-∆H)
warmer
endo(+∆H)
colder
exo(-∆H)
warmer
endo(+∆H)
colder

27. VIiI. Vapor Pressure ______ VP = _________________________ Example:
______ VP = _________________________ Example:
Factors for Vapor Pressure 1. 2.
Pressure a liquid “feels” pushing it to evaporate (turn to gas)
HIGH
Evaporates easily
Ethanol has higher VP than H2O, so it evaporates more easily
Strength of intermolecular force: molecules held together by dipole-dipole (polar) have lower VP than Van der Waals (nonpolar)
(H2O does not evaporate as fast as methane (CH4))
Temperature/Pressure: increase temp., increase VP

28. VIiI. Vapor Pressure Vapor Pressure of Four Liquids (See Table _____)
Questions:
1. What is the vapor pressure of propanone at 35 oC? _______
2. What temperature does water boil at if the pressure is 70 kPa? ______
3. What is the normal boiling point of ethanoic acid? __________
Sublimation:
Example: “Dry Ice” CO2 (s)  CO2 (g) H
Measured in: kPa (101.3 kPa = 1 atm = 760 mmHg = 760 torr.)
48 kPa
90 degrees C
117 degrees C
Occurs because solids have very weak IMF (usually Van der Waals).
They go directly from s  g and have HIGH VP

29. Check for Understanding
How could you change the boiling point temperature of a substance without adding anything to the substance?
You spill a glass of water on the floor. How could you get the water to evaporate faster, without using a mop or something to soak it up?
Change your altitude (higher altitude (lower pressure), lower boiling point)
Increase the temperature of the room. Spread out the puddle (increase surface area).