1. Chapter 1 Matter, Measurement, and Problem Solving
Honors Chem 1 A

2. 1.1 Atoms and Molecules
Atoms: submicroscopic building blocks of matter
Molecules: 2 or more atoms joined in a specific geometric arrangement
So….if atoms are submicroscopic, how do we know that they exist at all? How can we tell one atom from another?

3. The properties of matter depend on the atoms and molecules that compose them. Think of water. What is the formula for water? What would happen if the composition of water as changed slightly?
Molecular geometry is very important as well!
Chemistry: the science that seeks to understand the behavior of matter by studying the behavior of atoms and molecules!!!!

4. 1.2 The Scientific Approach to Knowledge
First, we need to think about what science is.
Right now as you sit in my class, your tummy is making a weird sound.
What is that sound?
Why is it happening?
What can you do to stop the sound?

5. Scientific Method!!!!
So, your tummy noise is a problem. It is a problem that needs to be solved! But you are a scientist, so lets solve it with science!
Problem: Tummy Noise
Hypothesis: If I eat a cookie, then my tummy will stop making noise
Experiment:
1. Record the number of tummy noises over 1 minute
2. Eat cookie
3. Wait 15 minutes
4. Record the number of tummy noises over 1 minute
5. Compare data
Data: Tummy made noise 4 times in one minute before cookie and 0 times in one mintute after cookie
Conclusion: Eating the cookie reduced/eliminated the number of tummy noises. Therefore, tummy noise was due to empty stomach/hunger. Eating a cookie filled stomach/eliminated hunger. Elimination of hunger reduced/eliminated tummy noise. Hypothesis was correct.

6. Law vs. Theory
This is a MAJOR pet peeve of mine….so learn it now. Or ELSE (evil laugh…..)
Scientific Law: simply states that something IS (sometimes called scientific principles) (summarizes accepted observations)
-Ex. Law of gravity: gravity exists
-Law of conservation of mass: mass is conserved.
NO WHY!!!!!

7. Well, what about why???
Scientific Theory: gives the reasons for the observed phenomenon
-Atomic theory: explains why atoms behave the way that they do and why matter exhibits certain properties
Theories can change and can be proven wrong. New information is discovered every day. However, that does not mean that theories should be taken lightly. A theory is ALMOST CERTAINTLY TRUE based on current available information/data. This is DIFFERENT from a hypothesis!!!

8. How can we use the scientific method?
Clearly identify your problem
Research the possible causes of this problem
Identify a possible solution to this problem (hypothesis)
-if…then statement
Test your solution and record your results (experiment)
Test it many, many times! The more times you test, the better your data. The better your data, the more supportable your conclusion.
If your data does not support your hypothesis, REPEAT!!!!
Very few hypotheses go on to become accepted theories.

9. So… lets summarize
Look at page 7 in your book. At the top of the page is a question. Think about this question. What is the best choice? Why? Do you agree with the book? Why/why not?

10. Homework
Pg 37, 1-7
Pg 38, (these are HARDER)

11. 1.3 Classification of Matter
Matter can exist in 3 different STATES
There are actually more states of matter, of 3 occur most commonly on earth.
SOLID: definite shape and volume. Atoms/molecules are densely packed. Vibrate! Atoms/molecules do not have enough energy to slide past one another.
Crystalline(regular, repeating pattern) or amorphous
Charcoal vs. Diamond

12. Liquid: definite volume, but no definite shape
Liquid: definite volume, but no definite shape. Will assume the shape of its container. Particles have enough energy to slide past one another, but not enough to move far apart from one another Gas: no definite shape, no definite volume. Particles have enough energy to bounce off of one another. Think about this: If you take a sample of a liquid and mass it. You then seal it in a closed container and vaporize it. What do you think will happen to the mass of the sample? Why?

13. Phase changes! Melting/freezing Vaporizing/condensing (and boiling)
Sublimation/deposition
What do you think has to happen for one state of matter to move to another state of matter?

14. Classifying Matter

15. Mixtures A physical blend of 2 or more substances is a mixture.
Can be Heterogeneous or homogeneous
Can be separated by physical change (THINK PHASE CHANGE)
Decanting, distillation, filtration
Volatile: easily vaporizable

16. 1.4 Physical and chemical changes/properties
Physical Property: can be observed without changing the chemical composition of the substance (color, odor, melting point)
Chemical Property: you must change the chemical composition of the substance to observe a chemical property
(flammability: can only test if you burn it!)
Physical Change: generally reversible
Chemical Change: Generally not reversible

17. Energy!
Physical and chemical changes are generally accompanied by energy changes.
Energy: the capacity to do work
Work: action of a force through a distance
Kinetic energy: energy of motion
Potential energy: energy of position/composition
Kinetic energy + potential energy = total energy
-Thermal Energy: energy associated with temperature of an object

18. Law of conservation of ENERGY
Energy is conserved… ENERGY CANNOT BE CREATED NOR DESTROYED!
Systems with high potential energy tend to change in a direction of lower potential energy, releasing energy into the surrondings.
Homework:
Pg 37, 8-19
Pg 38, 37-50

19. The importance of Measurement
Not all measurements give the same amount of information!
Qualitative measurements: give descriptive, non-numerical results (NO NUMBERS)
Think ‘quality’
Quantitative: give results with numbers and units (NUMBERS!
Think ‘quantity’. Quantity means numbers

20. Scientific notation
Scientific Notation: number is written as the product of two numbers: a number greater than or equal one and less than 10 and a power of 10. ex: 1,200 = x 103

21. How to write in scientific notation
1. Write the number you are starting with
2. Move the decimal behind the first digit in the number
Write ‘x 10’
Count the ‘jumps’ this is the ‘power of 10’
3. Decide if the old number is bigger or smaller than the new number
Bigger = positive jumps
Smaller = negative jumps

22. How to write in standard form
Standard form= this is just the regular way to write numbers
Look at the exponent (number on the x 10). If it is positive, make the number bigger. If it is negative, make the number smaller.
Count the jumps
Place new decimal

23. Write the following numbers in scientific notation:
1. 3,567,987
5. 6
6. 12,000

24. Write the following numbers in standard form
x 106
x 103
x 10-2
x 10-8
x 10 10 x

25. Accuracy in Measurements
Accuracy: a measure of how close a measurement comes to the actual value
Precision: a measure of how close a series of measurements are to one another
Error = experimental value – accepted value
Percent error = |error|/accepted value x 100

26. Significant Figures in measurements
Significant figures: all of the digits that are known, plus the last digit which is estimated.
Think of 3 different rulers:
One broken into whole meters
One broken into decimeters
One broken into centimeters
Which is more precise? WHY?

27. RULES FOR SIGNIFICANT FIGURES!!!!
1. Zeros between nonzero digits are significant
2. Zeros appearing in front of all nonzero digits are not significant
3. Zeros at the end of a number and to the right of a decimal point are significant
4. Zeros at the end of a number but to the left of a decimal point are significant

28. Exact Numbers!!!
Exact numbers have no uncertainty so there is no limit on significant figures.
Ex. 1 bird

29. You try it: Count the number of sig figs in each measurement
4) 23,013.00
5) 8,000,000
6) 303
1) 1, 7)
2) 19 birds in a tree 8) 3) 9)

30. Reading Instruments
In a lab, you will have to read instruments to the appropriate number of sig figs.
1. Read the instrument all the way
2. Estimate last digit (this is uncertain)
Ex. If your ruler goes to centimeters, your measurement will go to tenths of a centimeter
2.4 centimeters (the .4 centimeters is estimated)

31. Significant figures in calculations
Addition and Subtraction: round to the same number of DECIMAL PLACES as the number with the LEAST number of decimal places
Multiplication and Division: round to the same number of SIGNIFICANT FIGURES and the number with the LEAST number of significant figures

32. Let’s try it:
61.2 meters meters meters =
LOWEST NUMBER OF DECIMAL PLACES? OR 8.6 HAS ONE DECIMAL PLACE. ANSWER WILL BE ROUNDED TO ONE DECIMAL PLACE!!!
7.55 meters x 0.34 meter =
LOWEST NUMBER OF SIG FIGS? HAS 2 SIG FIGS. ANSWER WILL HAVE 2 SIG FIGS!!

33. International System of Units
International system of units (SI)= revision of the metric system. Every unit is a multiple of 10.
This makes converting from one unit to another EASY!

34. Units of Measurement
Length meter Volume cubic meter or liter Mass kilogram Density grams/cm3 Temperature kelvin or degrees Celsius Time second Pressure pascal or atmosphere Energy joule

35. Prefixes
Prefixes are used to tell you how many multiples of 10 something is. Sometimes the SI unit isn’t a practical measurement. For example: meters to measure the size of an atom. Atoms are way too small to be measured in meters, it’s not practical! We use a smaller multiple of meters and denote it with a prefix (nanometers)

36. Volume: the space occupied by a sample of matter
SI unit is milliliter or cubic centimeter
We often use liter to measure volume
WHY???

37. Mass Mass: the amount of matter in an object
Weight: measure of the pull of gravity on an object
Weight can change with location, mass can’t
SI unit = Kilogram
Kilo=1000
We often use grams in lab. WHY????

38. Density
Which is heavier, a pound of lead or a pound of feathers? What’s the difference between the two quantities? DENSITY!

39. Density = mass/volume
Matter with a lower density will float on matter with a higher density (liquids and gases)
Generally, density decreases as temperature increases.
water is most dense at 4 degrees Celsius
What happens with ice on water? What can you assume about the density of ice?
What happens when you fill a balloon with helium? What can you assume about the density of helium?

40. Intensive property vs. Extensive property Intensive does not depend on the amount of substance present -Color, density, flammability Extensive depends on the amount of substance present -Mass, energy content

41. 1. start with what you know 2. write conversion factor as a fraction
Dimensional Analysis
You must be able to convert from one unit to another. This skill will be used for the REST OF YOUR LIFE!!!!
1. start with what you know
2. write conversion factor as a fraction
3. what you want goes on top!
4. multiply or divide as necessary

42. Common Conversion factors
1 kilometer = 1,000 meters 1 meter = 100 cm 1 meter = 1,000 mm 1 hour = 60 minutes 1 yard = 3 feet 1 mole = 6.02 x 1023 particles

43. 273 meters to kilometers
start with what you know
273 meters
Write conversion factor as a fraction
1 km = 1000 meters 1km/ 1000meters
What you want goes on top!
I want to end with kilometers, so it goes on top!
Multiply or divide as necessary
273 meters x 1 km/ 1000 m = km

44. Convert to meters:
267 cm
273 mm
273 km
Convert to hours
273 minute
Convert to yards
273 feet

45. Temperature conversion
Degrees C = (degrees F -32)/ 1.8 K = degrees C What is Kelvin? (K)

46. Homework
Pg 37, 20-32
Pg 40, 51-92