1. Unit 7: matter & ENERGY

2. I. Important Terms Chemistry: Matter:
The study of matter and its changes
Any object that has a mass and volume (pretty much anything)

3. I. Important Terms Atom: -
Particle Diagram:
Smallest particle of an element that retains the properties of that element
Can not be broken down by a chemical reaction
2 different types of elements

4. I. Important Terms Compound: -
Particle Diagram:
Two ore more DIFFERENT ELEMENTS BONDED together
Can be broken down by a chemical reaction
1 type of compound

5. II. Phases of Matter The phases that matter is in depends on:
1. The space between atoms or molecules
2. The strength of the intermolecular force (IMF) between atoms

6. Phase Shape Volume Particle Diagram IMF Movement
Solid (s)
Liquid (l)
Gas (g)
strong
definite
vibrate
definite
Takes shape of con-tainer
Vibrate and rotate
Med-ium
definite
Vibrate and rotate and translate
Takes shape of con- tainer
Takes shape of con- tainer
weak

7. III. Classification of Matter
Element (Column A):
Compound (Column B):
Mixture (Column C):
Substance that cannot be changed into a simpler substance under normal conditions
Substance consisting of 2 or more different chemical elements that can be separated by chemical reactions (heat, electricity)
A physical blend of 2 or more substances that are NOT chemically combined (air, soup, tap water)

8. Aqueous Solution (aq):
a mixture of water and some other substance that can dissolve in water (salt water)
Element
Compound
Mixture

9. How Elements, Compounds, and Mixtures can be illustrated
N2 (g)
CO2 (g)
Air (N2, O2, H20, CO2(g)

10. IV. Physical vs. Chemical Properties
Properties of an element or compound that can be observed or measured WITHOUT a chemical reaction
Ability of an element or substance to undergo a chemical reaction (bond breaking) and form a NEW substance

11. IV. Physical vs. Chemical Properties
Color, texture, odor, density, melting/freezing temps, solubility, volume, mass
Reactivity, pH (acidity), ability to rust, decompose, ferment, combust
VIDEO EXAMPLE

12. IV. Physical vs. Chemical Properties

13. V. Physical vs. Chemical Changes
A change that does NOT produce a new substance, it just changes the position of the particles
Changing a substance into a NEW substance (bonds are broken and then new ones formed)
A color change may occur and a NEW s, l, or g is formed
A change that does NOT affect a substance’s chemical position

14. V. Physical vs. Chemical Changes
Any phase change (freezing, melting…), dissolving, mixing, cutting
Burning, rusting, fermentation, cooking/baking

15. V. Physical vs. Chemical Changes
liquid water freezes to ice
salt dissolves in water
liquid nitrogen in plastic bottle bursts open
rusting on a pan
potassium reacting with water to form potassium hydroxide and hydrogen gas
2K (s) + 2H2O (l) -> 2KOH (aq) + H2 (g)

16. V. Physical vs. Chemical Properties

17. A chemical reaction ALWAYS results in a new substance
Chemical Reaction Equation:
4 H
+2 O
6 Total
4 H
+2 O
6 Total =
Reactants
Products
2 H2 (g) + O2 (g)  2 H2O (l)
Coefficient
# of atoms
A chemical reaction ALWAYS results in a new substance

18. VI. Conservation of Mass:
Mass cannot be created or destroyed in a chemical reaction
The total mass of reactants equals the total mass of the products
Example (video clip)
Silver nitrate (AgNO3) and sodium chloride (NaCl) solutions before and after mixing

19. Examples: 50 g + Sodium+ 76 g 126 g Sodium Chloride Chlorine 
1) If 50.0 grams of sodium reacts with chlorine to form 126 grams of sodium chloride. How many grams of chlorine reacted?
2) If g of water is separated into hydrogen and oxygen gas, and the hydrogen gas has a mass of 20.0 g. What is the mass of the oxygen gas produced?
50 g +
Sodium+
76 g
126 g
Sodium Chloride
Chlorine 
178.8 g 
H2O 
20 g +
158.8 g
H2 + O2

20. VII. Classification of Matter
Substance
- Definite Composition (Homogenous)
Mixture of Substances
Physically
Separable

21. VII. 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

22. 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)

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

24. 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

25. VIII. 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

26. VIII. 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

27. 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.

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

29. IX. 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

30. 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

31. X. 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

32. 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

33. X. 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

34. X. 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

35. 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

36. XI. 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

37. XI. 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

38. XII. 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

39. 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).

40. XII. 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

41. XII. 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

42. XIII. 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

43. XIII. 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

44. XIV. 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

45. XIV. 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

46. XV. 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

47. Enthalpy: the amount of ________________________ used or released in a system.
+ ∆H = endothermic
- ∆H = exothermic
Systems in nature tend to be ____________________ (release energy).

48. Entropy: a measure of the ____________________ or randomness of a system.
+∆S = increasing entropy (more disorder)
-∆S = decreasing entropy (less disorder)
Systems in nature tend to ____________________ entropy.

49. Gibbs Free Energy
Gibbs Free Energy equation predicts reaction spontaneity: ∆G = ∆H - T∆S, it is ____________________ dependent.
Summary of Gibbs free energy
Enthalpy change (∆H)
Entropy change (∆S)
Gibbs free energy (∆G)
Spontaneity
Negative (exothermic)
Positive (increasing)
always negative
always spontaneous
Positive (endothermic)
depends on T, may be + or -
yes, if the temperature is high enough (i.e. melt ice)
Negative (decreasing)
yes, if the temperature is low enough (i.e. freeze ice)
always positive
never spontaneous (the surroundings must give energy to the system)
When ∆G = 0, the system is in equilibrium.