1. Unit 6: Physical Behavior of matter

2. I. Classification of Matter
Substance
- Definite Composition (Homogenous)
Mixture of Substances
Physically
Separable

3. I. Classification of Matter
Element
(Fe, K, Ca, Ne)
Compound
Two or more different elements bonded
Chemically
Separable
Ionic
(Metal and Nonmetal)
Molecular (covalent bond)
(Nonmetals)
Individual atoms

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. I. Classification of Matter
Heterogeneous
Nonuniform; distinct phases
Homogenous
Uniform throughout (air, tap water, solutions)

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

7. II. 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
PHYSICAL
CHEMICAL
CHEMICAL MEANS

8. II. 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

9. 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.

10. Think about this
What happens to the spacing and speed of particles at each of the phases?
SOLID LIQUID GAS

11. III. 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

12. IV. 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

13. IV. 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

14. 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

15. IV. 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

16. IV. 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

17. 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

18. V. 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

19. V. 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

20. VI. Measurement of Heat Energy
The amount of heat given off or absorbed in a reaction can be calculated using the following equation: (See Ref. Tabs. on Table ______ )
Specific Heat:
Specific Heat for water: __________ (Found on Table _______ in Ref. Tabs) T
q = heat (J)
m = mass (g)
c = specific heat (J/g*oC)
∆T = change in temperature (oC)
q = m c ∆T
The amount of heat it takes to raise the temperature of 1g of a substance 1oC
4.18 J/g*K B
Specific heat for concrete is 0.84 J/(gK) – why concrete is much hotter than water on a sunny summer day

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. VI. Measurement of Heat Energy
Example: How many joules are absorbed when 50.0 g of water are hater from 30.2 oC to 58.6 oC?
m = 50.0 g
Ti = 30.2 oC
Tf = 58.6 oC
q = ?
q = m c ∆T
q = (50.0 g) (4.18 J/goC ) (58.6 oC – 30.2 oC)
q = J  5940 J

23. VI. Measurement of Heat Energy
Example: How many joules of heat energy are released when 50.0 g of water are cooled from 70.0 oC to 60.0 oC?
Example: 50.0 g of water goes from K to K.
A) Is heat energy released or absorbed? B) Calculate the amount energy.
q = m c ∆T
m = 50.0 g
Ti = 30.2 oC
Tf = 58.6 oC
q = ?
q = (50.0 g) (4.18 J/goC ) (60.0 oC – 70.0 oC)
q = J ( - means heat is released)
m = 50.0 g
Ti = K  16.6 oC
Tf = K  36.6 oC
q = ?
q = m c ∆T
q = (50.0 g) (4.18 J/goC ) (36.6oC – 16.6oC )
q = J

24. VII. Heat Of FUSION q = mHf
Heat of Fusion for water: ____________ (Found on Table ________ in Ref. Tabs)
Equation: (Found on Table _______ in Ref. Tabs)
Amount of heat absorbed (endothermic) to change a substance from s to l at its melting point
334 J/g B T
q = mHf

25. VII. Heat Of FUSION q = mHf m = 255 g Hf = 334 J/g q = ?
Example: How many joules are required to melt 255 g of ice at 0.00oC?
Example: What is the total number of joules of heat needed to change 150 g of ice to water at 0.00oC?
q = mHf
m = 255 g
Hf = 334 J/g
q = ?
q = (255 g) (334 J/g)
q = 85,170 J  85,200 J
q = 50,100 J  5.0 x 10 4 J or 50. kJ

26. VIII. Heat Of Vaporization
Heat of Vaporization for water: ____________ (Found on Table ________ in Ref. Tabs)
Equation: (Found on Table _______ in Ref. Tabs)
Amount of heat absorbed (endothermic) to change a substance from l to g at its boiling point
2260 J/g B T
q = mHv

27. VIII. Heat Of Vaporization
Example: How many joules of energy are required to vaporize 423 g water at 100 oC and 1 atm?
Example: What is the total number of joules required to completely boil 125 g of water at 100 oC at 1 atmosphere?
q = mHv
m = 423 g
Hv = 2260 J/g
q = ?
q = (423g) (2260 J/g)
q = 955,980 J  956,000 J
q = 282,500 J  283,000 J

28. IX. Calorimetry - Measure the amount of heat given off in a reaction.
Used to:
- Measure the amount of heat given off in a reaction.
- Use q = m c ∆T to find the amount of heat lost or gained in a sample