1. Chemistry: The Science of Matter
A-D; A.5; D.2

2. Unit Objectives Classify matter according to its composition
Distinguish among elements, compounds, homogenous mixtures, and heterogeneous mixtures
Relate the properties of matter to structure

3. Proper Lab Safety Procedures

4. Appropriate Dress for Lab
Long hair must be tied back
Goggles must be worn at all times
Close-toed shoes must be worn

5. Safety Equipment Shower Fire Blanket Eye Wash Station
Fire Extinguisher

6. At your lab table… You may not have food or drink
You must notify a teacher of any chemical spill…don’t touch it!
You may not have bring any backpacks or other personal items other that your lab guide, a periodic table, a writing utensil, and a calculator
Keep the area free from clutter…when your done with it, return it or dispose it!
Spray and wipe down your lab tables.
Don’t run, horse play, and always be aware of your surroundings.

7. Lab Measurement

8. Using the Metric System - Length
Using Rulers – always use the “cm” or “mm” side.
You can estimate out one decimal place beyond the actual marking.
For each of the following, draw a downward arrow pointing to the correct value and label the arrow.
Where is 4.6 cm?
Where is 46 mm?
Where is 6.85 cm?
*Where is 35.5 mm?

9. Common Metric Units of Length
Millimeter (mm): 1000 mm = 1 m
Centimeter (cm): 100 cm = 1 m
Meter
Kilometer (km) : 1 km = 1000 m

10. Using the Metric System - Mass
Weighing a substance
All digital scales are in grams and read out to two decimal places.
ALWAYS put chemicals onto a weighing dish and NEVER directly on the scale.
Find the mass of the weighing dish first, then press the “zero” button

11. Common Metric Units of Mass
Milligram (mg): 1,000 mg = 1 g
Gram (g)
Kilogram (kg): 1 kg = 1000 g

12. Using the Metric System - Volume
Beaker
Erlenmeyer Flask
Measured using graduated cylinder, Erlenmeyer flasks, and beakers.
Just like a ruler, You can estimate out one decimal place beyond the actual marking.
Volume in our lab will be mostly measured in mL (milliliters)
Always read the bottom of the meniscus.
Graduated Cylinders

13. Reading Volume

14. Common Metric Units of Volume
Milliliter (mL) : 1,000 mL = 1 L
Centimeters cubed : 1 cm3 = 1 mL
Liter (L)

15. Using the Metric System - Temperature
We ALWAYS use the unit Celcius (never Fahrenheit!!)
Hint: Degrees Celcius will seem much smaller than Fahrenheit.
Always make sure your digital thermometer is on the “°C” setting

16. Assignment 2 Part B – Complete in Groups

17. Accuracy and Precision
How close an experimental measurement is to the actual/correct value
If the correct measurement is 5.65 g, which actual value is the most accurate: g, 5.63 g, 5.7
How close a series of measurements are to one another
Which series of measurements are the most precise?
5.15 g, 5.55 g., 5.59 g.
5.15 g., 5.99 g., 6.86 g.

18. % Error
Gives you a mathematic “gauge” as to how close (accurate) your results are to the accepted value

20. Observation Quantitative Qualitative
An observation made that involves a number.
Examples:
The mass of the rock is 5 g.
The length of the block is 4 cm.
The chemical changed color 5 times.
It took 2 minutes for the reaction to happen.
The temperature changed from 20 °C to 30°C
An observation made that does not involve a number.
Examples:
The chemical reaction produced a gas.
The color of the substance changed from clear to blue.
The reaction got warmer.
Block A feels heavier than Block B.

22. What is matter? Matter is anything that takes up space and has mass.
Where is matter? Everywhere! Can you name some things around you that have matter? That don’t have matter?
Mass is the measure of the amount of matter that an object contains
What has mass? Doesn’t have mass?
Not matter – heat, light, the internet

23. Properties of Matter
Properties of matter – describe the characteristics and behavior of matter, including the changes that matter undergoes.
Example (Figure 1.3, p. 6) – Iron is…
Able to rust
Melted at high temps
Malleable (bendable)
Ductile (stretchable)
Magnetic

24. Macroscopic vs. Microscopic
Matter that is large enough to be seen
Examples?
Matter that can not be seen without a microscope
Examples:
Bacteria
Cells
Atoms
Subatomic particles - Protons, Neutrons, Electrons

25. Pure substance vs. mixtures – The two categories of matter
Substance – matter with the same fixed composition and properties.
Any sample of pure matter is a substance.
A pure substance is either a:
Compound
Element

26. Pure Substances
Pure Substance that cannot be broken down into any other substances by chemical or physical means
Gold - element
Manganese Dioxide - compound

27. Pure Substance Element composed of identical atoms
Are found on the periodic table
EX: copper wire, aluminum foil
Courtesy Christy Johannesson

28. Pure Substances Compound
composed of 2 or more elements in a fixed ratio
properties differ from those of individual elements
Chemical bonds hold the elements together
EX: table salt (NaCl)
Courtesy Christy Johannesson

29. Pure Substances - FYI Law of Definite Composition
A given compound always contains the same, fixed ratio of elements.
Two different compounds, each has a definite composition
Courtesy Christy Johannesson

30. Mixtures
Mixture – a combination of two or more substances in which the basic identity of each substance is not changed
Most of the matter you encounter every day is a mixture.
Mixtures can either be:
Homogeneous
Heterogeneous

31. Mixtures
Variable combination of two or more pure substances. Each keep individual properties
Heterogeneous – Can see different parts (different)
Homogeneous- Evenly Mixed cannot see different parts. (Same)
Courtesy Christy Johannesson

32. Tyndall Effect The scattering of light by particles in a mixture

33. Homogeneous Mixtures
A mixture that is evenly mixed in such a way that you cannot see its individual parts.
A homogenous mixture is also known as a solution.
Example: Apple Juice

34. Mixtures Solution homogeneous very small particles no Tyndall effect
particles don’t settle
EX:
rubbing alcohol (ethyl alcohol and water)
Air (nitrogen and oxygen)
Courtesy Christy Johannesson

35. Heterogeous Mixture
A mixture in which you can distinguish the various components
Types of homogenous mixtures:
Colloids
Suspensions

36. Mixtures Colloid heterogeneous medium-sized particles Tyndall effect
particles don’t settle
Particles scatter light
EX:
Milk
Clouds
Smoke
mayo
Courtesy Christy Johannesson

37. Mixtures Suspension heterogeneous large particles Tyndall effect
particles settle
EX:
fresh-squeezed lemonade
Sand in water
Courtesy Christy Johannesson

38. Classification of Matter
(gas. Liquid,
solid, plasma)
Separated by
PURE
SUBSTANCES
MIXTURES
physical means into
Separated by
COMPOUNDS
ELEMENTS
HOMOGENEOUS
MIXTURES
HETEROGENEOUS
MIXTURE
chemical
means into
Kotz & Treichel, Chemistry & Chemical Reactivity, 3rd Edition , 1996, page 31

39. Classification of Matter
                                                                                                                                                      
Matter
Physically
separable
Substance
Definite composition
(homogeneous)
Mixture of
Substances
Variable composition
Basis for separation: different components, different properties.
Strategy: devise a process that discriminates between components with different properties.
high density / low density reactive / inert volatile / nonvolatile
soluble / insoluble polar / nonpolar magnetic . nonmagnetic
Chemically
separable
Element
(Examples: iron, sulfur,
carbon, hydrogen,
oxygen, silver)
Compound
(Examples: water.
iron (II) sulfide, methane,
Aluminum silicate)
Homogeneous mixture
Uniform throughout,
also called a solution
(Examples: air, tap water,
gold alloy)
Heterogeneous mixture
Nonuniform
distinct phases
(Examples: soup,
concrete, granite)

40. Elements Compounds Mixtures
Both elements and compounds have a definite makeup and definite properties.
Elements
only one kind
of atom; atoms
are bonded it
the element
is diatomic or
polyatomic
Compounds
two or
more kinds
of atoms
that are
bonded
Mixtures
two or more
substances
that are
physically
mixed
substance
with
definite
makeup
and
properties
two or
more
kinds of
and
Packard, Jacobs, Marshall, Chemistry Pearson AGS Globe, page (Figure 2.4.1)

41. Classification of Matter
hetero-
geneous
mixture no
uniform
properties? no
solution
fixed
composition? no
element
chemically
decomposable?
yes
compound

42. Matter Flowchart Examples: graphite pepper element sugar (sucrose)
paint
soda
element
hetero. mixture
compound
hetero. mixture
Graphite image: geology.about.com/.../bl/images/blgraphite.htm
solution
homo. mixture
Courtesy Christy Johannesson

43. The Composition of Air Air Nitrogen Helium Oxygen Neon Water vapor
Carbon
dioxide
Argon
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 34

44. Top Ten Elements in the Universe
Percent
Element (by atoms)
Hydrogen 73.9
Helium 24.0
Oxygen
Carbon
Neon
Iron
Nitrogen
Silicon
Magnesium
Sulfur
A typical spiral galaxy
(Milky Way is a spiral galaxy)
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 26

45. Physical Properties & Changes
A physical property is a characteristic of a sample of matter that can be observed or measured without any change in its identity.
Examples/key words – solubility, melting point, boiling point, color, density, electrical conductivity, and physical state (solid, liquid, gas) (they are all generally adjectives)
A physical change is a change in matter that does not involve a change in the identity of the substance
Examples/key words – boiling, freezing, melting, evaporating, dissolving, separating, and crystallizing (they are all generally verbs)
A.2;

46. Chemical Properties & Changes
A chemical property is the ability of a substance to undergo a change such as reacting with other substances or decomposing. (generally adjectives)
A chemical change is the change of one or more substances into other substances. This is also referred to as a chemical reaction.
Examples & key words: decompose, explode, rust, oxidize, corrode, tarnish, ferment, burn, react, changes color, bubbles/fizzes, or rot (generally verbs)

47. Physical or Chemical Property
Physical or Chemical Property? -Identify each property below as either a chemical or physical property of a substance:
Questions:
Answers:
Aluminum bends easily
Copper sulfate dissolves in water
Magnesium burns in air
Gold jewelry is unaffected by perspiration
Basking soda is a white powder
Fluorine is a highly reactive element 1. 2. 3. 4. 5. 6.

48. Launch Lab – Why is the mass different? – p. 3
Record your observations.
Answer the “Analysis” questions on p. 3
Class Discussion: How can three objects with the same volume have different masses?

49. Density formula: D = m v Density Pyramid:
Density is the amount of matter (mass) contained in a unit of volume.
Examples:
A foam cup has less density than a stone because it has more mass in a identical-sized sample.
Density formula: D = m v Density Pyramid:
Demo -
Lab Challenge -

50. Density Practice Set #1 – Solving for density
1. Calculate the density of an object that has a mass of 2.53 g and a volume of 4.54 mL
2. Calculate the density of an object that has a mass of 16.0 g and a volume of 25.3 mL.
*3. Calculate the density of an object that has a mass of 3.01 g and a volume of 5.08 cm3
Which substance above has the greatest density?
…Least density?
0.56 g/mL
O.63 g/mL
0.59 g/cm3

51. Density Practice Set #2 – solving for mass or volume
1. What is the mass of a sample that has a density of 2.0 g/mL and a volume of 4.6 mL?
2. What is the volume of a sample that has a density of 0.23 g/mL and a mass of 2.5 g?
Fill in the missing cells on the chart below:
9.2 g
10.9 mL
Chart: d = 0.18 g/mL; m = 10.7 g; v = 3.2 mL
Mass (m)
Volume (v)
Density (d)
1.4 g
7.6 mL
5.1 mL
2.1 g/mL
24.2 g
7.5 g/mL

52. States of Matter
The states of matter on Earth are solid, liquid gas – these are physical properties of a substance
A change of state is the temperature at which a substance changes from one state to another.
Water freezes/melts at 0°C
Table salt (sodium chloride) freezes/melts at 804°C
Oxygen freezes/melts at -218°C
A change in state is a physical change
C.2

53. Changing States of Matter

54. Law of Conservation of Mass
While atoms of substances do and can change, they never disappear or appear from no where.
The law of conservation of mass (or matter) states that in a chemical change, matter is neither created nor destroyed.

55. Before = After … ALWAYS!!

56. Endothermic vs. Exothermic Reactions
All chemical changes also involve some sort of energy change.
Energy is either absorbed (endothermic reaction) or released (exothermic reaction) as a chemical reaction is taking place. C
Endothermic Reaction
C.2
Boring:
Do a demo?Ammonium chloride and barium hydroxide – endothermic (p. 41)
Exothermic – sugar and sulfuric acid