Matter: Properties

1
Matter Properties Change

• Chapter 3

2
A. Matter

• Matter anything that has mass and takes up
  space
• Everything around us
• Mass measurement that reflects the amount of
  matter (usually in grams)
• Volume the amount of space something takes up
• Chemistry the study of matter and the changes
  it undergoes

3
B. Four States of Matter

• Solids
• particles vibrate but cant move around
• fixed shape
• fixed volume
• incompressible

4
B. Four States of Matter

• Liquids
• particles can move around but are still close
  together
• variable shape
• fixed volume
• Virtually incompressible

5
B. Four States of Matter

• Gases
• particles can separate and move throughout
  container
• variable shape
• variable volume
• Easily compressed
• Vapor gaseous state of a substance that is a
  liquid or solid at room temperature

6
B. Four States of Matter

• Plasma
• atoms collide with enough energy to break into
  charged particles (/-)
• gas-like, variableshape volume
• stars, fluorescentlight bulbs, TV tubes

7

• II. Properties Changes in Matter (p.73-79)
• Extensive vs. Intensive
• Physical vs. Chemical

8
A. Physical Properties

• Physical Property
• can be observed measured without changing the
  identity of the substance

9
A. Physical Properties

• Physical properties can be described as one of 2
  types
• Extensive Property
• depends on the amount of matter present (example
  length, mass, volume)
• Intensive Property
• depends on the identity of substance, not the
  amount (example scent, density, melting point)

10
D. Chemical Properties

• Chemical Property
• describes the ability of a substance to be
  observed reacting with or changing into another
  substance

11
E. Physical vs. Chemical Properties

• Examples
• melting point
• flammable
• density
• magnetic
• tarnishes in air

• physical
• chemical
• physical
• physical
• chemical

12
F. Physical Changes

• Physical Change
• changes the form of a substance without changing
  its identity
• properties remain the same
• Examples cutting a sheet of paper, breaking a
  crystal, all phase changes

13
F. Phase Changes Physical

• Evaporation
• Condensation
• Melting
• Freezing
• Sublimation
• Deposition

• Liquid -gt Gas
• Gas -gt Liquid
• Solid -gt Liquid
• Liquid -gt Solid
• Solid -gt Gas
• Gas -gt Solid

14
G. Chemical Changes

• Process that involves one or more substances
  changing into a new substance
• Commonly referred to as a chemical reaction
• New substances have different compositions and
  properties from original substances
• Reaction involves reactants reacting to
  produce products

15
G. Chemical Changes

• Signs of a Chemical Change
• change in color or odor
• formation of a gas (bubbles)
• formation of a precipitate (solid)
• change in light or heat

16
H. Physical vs. Chemical Changes

• Examples
• rusting iron
• dissolving in water
• burning a log
• melting ice
• grinding spices

• chemical
• physical
• chemical
• physical
• physical

17
What Type of Change?
18

• Exothermic- heat energy EXITS the system
• surroundings usually feel warmer
• 1 g H2O (g) ? 1 g H2O (l) 2260 J
• ex. Combustion, evaporation of water

19

• Endothermic- heat energy ENTERS the system
• - heat absorbed from surroundings
• - surroundings usually feel cooler
• - 1 g H2O (s) 333 J ? 1 g H2O (l)
• - 1 g H2O (l) 2260 J ? 1 g H2O (g)
• - ex. Cold packs, melting ice

20
What Type of Change?
?

• ?

?
?
21
I. Law of Conservation of Mass

• Although chemical changes occur, mass is neither
  created nor destroyed in a chemical reaction
• Mass of reactants equals mass of products

massreactants massproducts
A B ? C
22

• III. Classification of Matter (pp. 80-87)
• Matter Flowchart
• Pure Substances
• Mixtures

23
A. Matter Flowchart
MATTER
yes
no
Can it be physically separated?
Homogeneous Mixture (solution)
Heterogeneous Mixture
Compound
Element
24
A. Matter Flowchart

• Examples
• graphite
• pepper
• sugar (sucrose)
• paint
• soda

• element
• hetero. mixture
• compound
• hetero. mixture
• solution

25
B. Pure Substances

• Element
• composed of one type of identical atoms
• Atom Composed of protons, electrons, and
  neutrons. Smallest particle of matter that can
  be identified as one element
• EX copper wire, aluminum foil

26
B. Pure Substances

• Compound
• composed of 2 or more elements in a fixed ratio
  (bonded together)
• properties differ from those of individual
  elements
• EX table salt (NaCl)

27
C. Mixtures

• Variable combination of 2 or more pure
  substances, each retains its chemical identity
  properties.

Heterogeneous
Homogeneous
28
C. Mixtures

• Homogeneous are uniform throughout
• Solutions
• very small particles
• particles dont settle
• EX rubbing alcohol, gasoline, soda

29
C. Mixtures

• Heterogeneous
• medium-sized to large-sized particles
• particles may or may not settle
• EX milk, fresh-squeezed lemonade

30
C. Mixtures

• Examples
• tea
• muddy water
• fog
• saltwater
• Italian salad dressing

• Answers
• Solution
• Heterogeneous
• Heterogeneous
• Solution
• Heterogeneous

31
Separation Methods

• Ways to separate mixtures Chapter 3 Matter
  Its Properties

32
Separating Mixtures

• Substances in a mixture are physically combined,
  so processes bases on differences in physical
  properties are used to separate component
• Numerous techniques have been developed to
  separate mixtures to study components
• Visually
• Magnetism
• Filtration
• Distillation
• Crystallization
• Chromatography

33
Filtration

• Used to separate heterogeneous mixtures composed
  of solids and liquids
• Uses a porous barrier to separate the solid from
  the liquid
• Liquid passes through leaving the solid in the
  filter paper

34
Distillation

• Used to separate homogeneous mixtures
• Based on differences in boiling points of
  substances involved

35
Crystallization

• Separation technique that results in the
  formation of pure solid particles from a solution
  containing the dissolved substance
• As one substance evaporates, the dissolved
  substance comes out of solution and collects as
  crystals
• Produces highly pure solids
• Rocky candy is an example of this

36
Chromatography

• Separates components of a mixture based on
  ability of each component to be drawn across the
  surface of another material
• Mixture is usually liquid and is usually drawn
  across chromatography paper
• Separation occurs because various components
  travel at different rates
• Components with strongest attraction for paper
  travel the slowest

37
Thermochemistry

• Chapter 171
• Pages 505 510

38
A. Vocabulary
ThermochemistryThe study of energy changes
that occur during chemical reactions and changes
in state
EnergyThe capacity to do work or produce heat
39
A. Vocabulary

• TEMPERATURE is a measure of the amount of kinetic
  energy an object/substance has
• HEAT is energy that transfers from one
  object/substance to another (thermal energy)
• Represented by the symbol q
• Transfers because of difference in temperature
• Always flows from a warmer to a cooler object

40
A. Vocabulary

• Which has more thermal energy?

41
B. Heat Transfer

• Why does A feel hot and B feel cold?

• Heat flows from A to your hand hot

• Heat flows from your hand to B cold

42
C. Types of Energy

• Potential due to position or composition can
  be converted to work
• Kinetic due to motion of the object
• KE ½ mv2
• (m mass, v velocity)
• Law of Conservation of Energy energy can be
  neither created nor destroyed

43
C. Types of Energy

• Energy that is stored in the chemical bonds of a
  substance is called CHEMICAL POTENTIAL ENERGY
• Types of atoms and their arrangement determine
  amount of energy stored in substance

44
D. Exothermic and Endothermic

• System the reaction (our focus)
• Surroundings everything around the reaction
  (rxn container, room, etc)
• Surroundings
• Universe System Surroundings

System
45
D. Exothermic and Endothermic

• Exothermic process heat is released into the
  surroundings
• Exo Exit
• Exothermic processes are represented by a
  negative q

46
B. Exothermic and Endothermic

• Combustion of Methane

47
D. Exothermic and Endothermic

• Endothermic Process heat is absorbed from the
  surroundings
• Endo Into
• Endothermic processes are represented by a
  positive q

48
D. Exothermic and Endothermic

• Formation of Nitric Oxide

49
D. Exothermic and Endothermic

• Sign of (q) is a signal to indicate the
  direction of the heat transfer

• Exothermic
• - q heat transferred from a substance

• Endothermic
• q heat transferred into a substance

50
E. Measuring Heat Flow

• Two Common Units
• Joule
• calorie
• 4.184 J 1 cal
• 1J 0.2390 cal
• 1Calorie 1 kilocal 1000 cal

51
F. Heat Capacity

• the amount of heat needed to increase the
  temperature of an object by 1oC
• Heat Capacity depends on
• The mass of the object
• The chemical composition of the object

52
G. Specific Heat Capacity

• Specific heat capacity amount of heat needed to
  raise the temperature of 1g of a substance by 1oC
• Specific Heat of H2O (l) 4.184 J/goC
• Specific Heat of H2O (s) 2.02 J/goC
• Molar heat capacity amount of heat needed to
  raise the temperature of 1 mole of a
  substance by 1oC

53
G. Specific Heat Capacity
q m ? C ? ?T
q heat (J) m mass (g) C specific heat (J/gK)
or (J/goC) ?T change in temperature (K or C)
54
G. Specific Heat Capacity

• q m x C X ?T
• q heat (joules or calories)
• m mass (grams)
• C Specific Heat
• ?T change in temperature
• The change in temperature can be measure in
  Kelvin or degrees Celsius

55
G. Heat Transfer Problem

• A 32.0-g silver spoon cools from 60.0C to
  20.0C. How much heat is lost by the spoon?

GIVEN m 32.0 g Ti 60.0C Tf 20.0C q ? C
.235 J/gC
WORK q mC?T ?T 20C - 60C 40C q
(32.0g)(-40C)(.235J/gC) q 301 J
Exo
56
G. Heat Transfer Problem

• The temperature of a 95.4g piece of copper
  increases from 25.0 oC to 48.0 oC when it absorbs
  849 J of heat. What is the specific heat of
  copper?

WORK q mC?T C q/m ?T ?T 48C 25C
23C C 849J/(95.4g)(23C) C 0.387 J/gC
GIVEN m 95.4 g Ti 25.0C Tf 48.0C q 849
J C ?
Endo
57
G. Heat Transfer Problem

• What is the molar heat capacity of copper?

WORK C 0.387 J
GIVEN C 0.387 J/gC MM 63.55 g/mol
63.55 g Cu 1 mol Cu

g oC
24.6 J/mol oC
58
Thermochemistry

• Chapter 172
• Pages 511 517

59
A. Calorimetry

• Measures the heat flow into or out of a system
• Heat released by the system is equal to heat
  absorbed by the surroundings

• Calorimeter Insulated device used to measure
  absorption or release of heat

60
A. Calorimetry

• Enthalpy (H) the heat absorbed or released of a
  system at constant pressure
• ? H is the heat of a reaction
• Heat or enthalpy change are used interchangeably
  here
• q ? H

61
B. Thermochemical Equations

• In a thermochemical equation, the enthalpy of
  change for the reaction can be written as either
  a reactant or a product
• Endothermic (positive ?H)
• 2NaHCO3 129kJ Na2CO3 H2O CO2
• ?H 129 kJ
• Exothermic (negative ?H)
• CaO H2O Ca(OH)2 65.2kJ
• ?H - 65.2 kJ

62
A. Calorimetry

• ?H Hfinal - Hinitial OR
  
• ?Hreaction (rxn) Hproducts Hreactants
•  
• For endothermic reactions HfinalgtHinitial ?H
  is positive (?H)
• For exothermic reaction HfinalltHinitial and ?H is
  negative (-?H)

63
A. Calorimetry

• When solving for the heat transfer of a system
  (between 2 objects), assume that
• q initial soln - qfinal rxn
• heat goes in heat goes out
• ?H -?H or
• CH20 x m H20 x ?T H20 -1(Cmetal x mmetal x
  ?Tmetal)

64
A. Calorimetry

• Example
• 2 H2 O2 ? 2 H2O 483.6 kJ
• ?H - 483.6 kJ
• Exothermic

65
A. Calorimetry

• How to measure ? H for a reaction in aqueous
  solution
• Dissolve chemicals in water
• Place in calorimeter
• Measure temperature change

66
A. Calorimetry

• qsurr m x C x ?T
• qsys ?H - qsurr -m x C x ?T
• Negative enthalpy exothermic
• Positive enthalpy endothermic

67

• When 25.0mL of water containing HCl at 25.0 oC is
  added to 25.0mL of water containing NaOH at 25.0
  oC in a calorimeter a rxn occurs. Calculate the
  enthalpy change (in kJ) during the rxn if the
  highest temperature observed was 32.0 oC. Assume
  all densities 1.00g/mL
• KNOWN
• Cwater 4.184 J/g oC
• V 25.0mL 25.0mL
• ?T 32.0 25.0 7.0 oC
• Density 1.00g/mL
• M (50mL) x (1.00g/mL) 50g
• ?H ?

• ?H -mC?T
• -(50.0g)(4.184J/goC) (7.0oC)
• ?H -1463 J -1460J

Exo
68
C. Heat of Combustion

• The heat of reaction for the complete burning of
  one mole of a substance
• Written the same way as change in enthalpy

69

• Write the thermochemical equation for the
  oxidation of Iron (III) if its ?H -1652 kJ
• Fe(s) O2(g)? Fe2O3(s) 1652 kJ
• How much heat is evolved when 10.00g of Iron is
  reacted with excess oxygen?

Exo
4
3
2
1 mol
10.00g Fe
1652 kJ
73.97 kJ of heat
4 mol Fe
55.85g Fe
70
Thermochemistry

• Chapter 173
• Pages 520 526

71
A. Heat of Fusion

• Heat of Fusion (?Hfus)
• Heat absorbed by one mole of a solid substance
  when it melts to a liquid at a constant
  temperature
• ?Hfus of ice 6.009 kJ/mol

• Heat of Solidification (?Hsolid)
• Heat lost by one mole of a liquid substance when
  it solidifies at a constant temperature
• ?Hfus - ?Hsolid
• ?Hsolid of water -6.009 kJ/mol

72
B. Heating Curves
73
B. Heating Curves

• Temperature Change
• change in KE (molecular motion)
• depends on heat capacity

• Phase Change
• change in PE (molecular arrangement)
• temp remains constant

74
C. Heat of Vaporization

• Heat of Vaporization (?Hvap)
• energy required to boil 1 gram of a substance at
  its b.p.
• ?Hvap for water 40.79 kJ/mol
• usually larger than ?Hfuswhy?

• EX sweating, steam burns

75
D. Practice Problems

• How much heat energy is required to melt 25 grams
  of ice at 0oC to liquid water at a temperature of
  0oC?

ice
6.009 kJ 1 mol H2O
25 g H2O
1 mol H2O 18.02 g H2O
8.3 kJ
76
D. Practice Problems

• How much heat energy is required to change 500.0
  grams of liquid water at 100oC to steam at 100oC?

steam
40.79 kJ 1 mol H2O
500.0 g H2O
1 mol H2O 18.02 g H2O
1132 kJ
77
D. Practice Problems

• How many kJ are absorbed when 0.46g of C2H5Cl
  vaporizes at its normal boiling point? The molar
  ?Hvap is 26.4 kJ/mol.

26.4 kJ 1 mol C2H5Cl
0.46 g C2H5Cl
1 mol C2H5Cl 64.52 g C2H5Cl
0.19 kJ
78
E. Heat of Solution

• During the formation of a solution, heat is
  either released or absorbed
• Enthalpy change caused by dissolution of 1 mol of
  a substance is the molar heat of solution ?Hsoln
• Examples hot packs, cold packs

79
E. Heat of Solution

• NaOH(s) ? Na(aq) OH-(aq)
• ?Hsoln -445.1 kJ/mol
• NH4NO3 ? NH4(aq) NO3-(aq)
• ?Hsoln 25.7 kJ/mol