1. INTRODUCTION

2. CHEMISTRY
Is the study of the composition of materials and the changes that these materials undergo
Wood burns
Plants grow
Iron rusts
Bread bakes
Human digestion

3. Branches of Chemistry Analytical chemistry Physical chemistry
Answers questions about composition (what is the level of lead in that drinking water?)
Physical chemistry
Deals with energy transfers in chemical processes (how much energy is stored in a battery?)
Biochemistry
Deals with the analysis of chemical processes in living things (cellular process such as photosynthesis and respiration)
Organic chemistry
Study of carbon containing molecules (overlap with biochem)
Inorganic chemistry
Study of the chemistry of nonliving things (geological chemistry, the analysis of the rate of corrosion of metals)

4. The Scientific Method State the problem Make observations
Form a hypothesis
Perform an experiment
Variable you change – called the “independent” or “manipulated” variable
Variable you observe – called the “dependent” or “responding” variable
Develop a theory (a well-tested explanation for observations)
Continue to experiment
Modify theory as needed
Develop a scientific law (a summary of many observations and experiments)

5. Chapter 2: Matter & Change
Matter – anything that has mass and volume
_____________________ – require measurements
Mass – amount of matter, measured in grams (g)
Volume – amount of space an object occupies, measured in:
liters (L)
milliliters (ml)
cubic centimeters (cm3)
_________________ – descriptions, not measured
Color, texture, odor, opaque, clear, dense, sound, taste (don’t taste anything in the lab!)

6. Classification of matter

7. Matter: Substances and Mixtures
Substances – pure composition
Two types:
__________ – only one kind of atom, such as iron (the symbol is on the periodic chart)
__________ – a combination of atoms, such as iron oxide (which can be separated)
Mixtures – contains 2 or more substances
______________ – the same throughout; dissolved (saltwater)
______________ – the sample may different depending where you take it in the container (orange juice with pulp)

9. States of Matter
Solid- rigid shape, incompressible, fixed volume, usually particles tightly packed and organized)
Liquid- indefinite shape but fixed volume, incompressible, flows, usually less tightly packed than solid (by a few %), less organized
Gas- indefinite shape and volume, flows, particles far away from each other, disorganized, compressible

10. CHEMICAL REACTIONS Chemical Reaction Reactant Product
A change in which ________________________________________________________________________
Reactant
A ___________________________________
Product

11. Physical Properties and Changes
Physical Property – characteristics that describe a substance
Can be observed without changing the substances composition.
Physical Change – a change that does not affect the composition of the substance
The change may be reversible (melting) or irreversible (tearing)

12. Recognizing Chemical Changes
You know a chemical change has occurred if:
you see a change in color
there is a change in temperature (hot or cold)
a gas is produced
a solid precipitate is formed
an odor is produced

13. Separation of a Mixture Difference in physical properties can be used to separate mixtures.
Filtration- separates solid from a liquid (passes through a filter (pores))
Distillation-liquid is boiled to produce a vapor that is then condensed into a liquid

14. Law of Conservation of Mass
During any physical or chemical change, the mass of the products is always equal to the mass of the reactants.
MASS is NEITHER CREATED or DESTROYED
Example: making pizza
Physical Change: the mass of the dough + sauce + cheese separately = the mass of all three together
Chemical Change: the mass of the ingredients before it’s cooked = the mass after it’s cooked (as long as it didn’t burn!)

15. DENSITY Mass Volume Density
Amount of matter in an object
Volume
Amount of space an object occupies
Density
Mass per unit volume
Percent Error = /theoretical-experimental/ x 100%
theoretical

16. Density Practice
A copper penny has mass of 3.1g and volume of 0.35cm3. What is the density of copper?

17. Measurements Three types of measurement Length: meter (m)
Mass: gram (g)
Volume: Liter (L)

18. Measurements expressed in SCIENTIFIC NOTATION
The expression of numbers in terms of M x 10n
M ≥ 1.00 and M < 10
n is an integer
Convert to whole numbers
1 x 103 = _______
1 x 102 = _______
1 x 101 = _______
1 x 100 = _______
1 x 10-1 = _______
1 x 10-2 = _______
1 x 10-3 = _______

19. TRY THESE
25,000
______________
0.0100

20. Metric Conversions
Kilo . Hecto . Deka . (unit) . Deci . Centi . Milli … Micro … Nano m g L
Write the following in regular notation and scientific notation
How many mm in a m? _____________
How many cm in a m? _____________
How many m in a km? _____________

21. Measurements

22. ACCURACY/PRECISION Accuracy Precision
How close a measurement is to the actual value
Precision
How close a set of measurements are to each other

23. Temperature Official” unit = Kelvin (K)
Most widely used for measurement = degrees Celsius (ºC)
Temp in Kelvin = ºC (º C = K -273)
Examples:
ice water = 0 ºC = 273 K
Boiling water = 100 ºC = 373 K
Body temp = 37 ºC = 310 K (about 99 ºF)

24. SIGNIFICANT DIGITS or Figures
This is a process used to determine the number of digits to round to when measuring an object.
Use this process when
Measuring mass (on the scale) – g, kg, etc.
Measuring volume (in a graduated cylinder) – ml L, etc.
Measuring length (with a ruler) – cm, m
Is used to communicate to other scientists how accurate your measurement is:
Does your scale measure to the hundredths place, tenths place or whole number?
Referred to as “Sig Figs”

25. How to determine the number of Sig Figs in a measured value
Atlantic-Pacific Method
A = decimal Absent, begin counting from right
P = decimal Present, begin counting from left
Try these:
1,000
1 sig fig
0.001
0.0010
2 sig fig
1000.0
5 sig fig

26. Rules for Using Sig Figs
Multiplication/Division
Do all calculations, then round to the same number of digits as the number with the smallest number of sig figs
4.56 x =
Round to 2 sig figs: 6.4
8.315/298 =
Round to 3 sig figs:
Addition/Subtraction
Do the calculations, then round to the place of the number with the smallest number of decimal places =
Round to 31.1
88.88 – 2.2 = 86.68
Round to (note: if the number after 6 is > 5, round up)

27. Rules for Using Sig Figs
Multiple step calculations
Use an overbar to keep track of the significant figures from step to step.
Round only when reporting the final answer
Example:
88.88 – (calculator
answer)
Based on 2.22, round to 3 sig figs
If the number after the place you want to round to is > 5, round up (in this case 7). Ignor the other 7.
Answer = .0250
The zero after the 5 is significant. You must show it! = =