1. Basic Concepts of Matter
Chapter 1
Basic Concepts of Matter

2. Classifying Matter Matter Mass vs. Weight Kinetic-Molecular Theory
Anything that has mass and occupies space
Mass vs. Weight
Kinetic-Molecular Theory
All matter consists of extremely tiny particles in constant motion

3. States of Matter Solid Liquid Gas
-Closely packed together with a definite ridged shape
-Vibrate back and forth in a confined space
-the particles are not able to move past one another
Liquid
-arranged randomly with a definite volume
-“fluid”
-the particles are not confined in space and can move past one another
Gas
-no definite shape or volume
-the particles are far apart and move very rapidly colliding with other particles and the container walls

4. Categorizing Matter Elements Pure Substance
-cannot be decomposed into simpler form via chemical reactions
-found on periodic chart
-atoms are the smallest particle that retains the characteristic properties of the elements
Pure Substance
-consists of all the same substance (pure gold, distilled water, etc)
-have a set of unique properties that identifies it

5. Categorizing Matter Chemical Compounds
-two or more elements in a definite ratio by mass with unique properties that separate them from the individual elements
-can be decomposed into the constituent elements by chemical reactions
-chemical compounds are held together by a chemical bound
Water– hydrogen and oxygen
Carbon dioxide – carbon and oxygen

6. Categorizing Matter Mixtures homogeneous mixtures (solution)
two or more pure substances in the same container
homogeneous mixtures (solution)
-uniform composition throughout
-single phase
-cannot be separated easily
heterogeneous mixtures
-nonuniform composition thoughout
-easily separated

7. Physical and Chemical Changes
Physical changes
changes in physical properties
-melting, boiling, and cutting
Chemical changes
changing one or more substances into one or more different substances (chemical reaction)
2H2 + O2 -> 2H2O

8. Chemical and Physical Properties
Chemical Properties
observed during a chemical reaction (change in chemical composition)
-rusting, oxidation, burning…
-chemical reactions
Physical Properties
observed without changing the substance’s composition
-allow for identification and classification
-density, color, solubility, melting point…

9. Classification of Physical Properties
Extensive Properties
depend on the amount of substance present
-mass or volume
Intensive Properties
do not depend on the amount of substance present
-melting point, boiling point, density…

10. Density Describes how compact a substance is Who “discovered” density?
Density = mass/volume or D = m/V

11. Density
Example: Calculate the density of a substance if 742 grams of it occupies 97.3 cm3.
1cm3 = 1mL => 97.3cm3 = 97.3mL

12. Density
Example: You need 125 g of a corrosive liquid for a reaction. What volume do you need?
liquid’s density = 1.32 g/mL

13. Units of Measure Qualitative measures Quantitative measures
Nonnumerical experimental observations describing the identity of a substance in a sample
Quantitative measures
Numerical experimental observations describing how much of a particular substance is in a sample
System International d’Unites (SI)
measurement system used in the sciences based on the metric system

14. Math Review and Measurements
We make measurements to understand our environment:
Human senses: sight, taste, smell, hearing…
Our senses have limits and are biased
Instruments: an extension of our senses meter sticks, thermometers, balances
These are more accurate and precise
All measurements have units
METRIC SYSTEM vs. British System

15. SI units Quantity Unit Symbol length meter m mass kilogram kg
time second s
current ampere A
temperature Kelvin K
amt. substance mole mol

16. Measurements in Chemistry
Name Symbol Multiplier
mega M (1,000,000)
kilo k (1,000)
deka da 10
deci d (0.1)
centi c (0.01)

17. Measurements in Chemistry
Name Symbol Multiplier
Milli m (0.001)
Micro  ( )
Nano n
Pico p
Femto- f

18. Units of Measurement Length Measure of space in any direction
-derived unit cm
-standard length is a meter (m)

19. Units of Measurement Volume Amount of space occupied by matter
-derived unit: mL or cm3 (cc)
-liter (L) is the standard unit

20. Units of Measurement Time (t) Mass (m) Weight (W)
Interval or duration of forward events
-standard unit is the second (s)
Mass (m)
measure of the quantity of matter in a body
Weight (W)
measure of the gravitational
attraction (g) for a body (w=m x g)
1 kg = 1000g kg = 2.2 lbs
1 g = 1000mg

21. Heat and Temperature Heat (q) vs. Temperature (T)
3 common temperature scales:
all use water as a reference
-Fahrenheit (F)
-Celsius (C)
-Kelvin (K)

22. Temperature Reference Points
Melting Point Boiling Point
of water of water
32 oF oF
0.0 oC cC
273 K K
Body temperature 37.0 oC or 98.6 oF
37.2 oC and greater—sick
41 oC and greater, convulsions
<28.5 oC hypothermia
Fahrenheit Celsius Kelvin

23. Temperature Scales

24. Fahrenheit to Centigrade Relationships
Temperature Scales
Fahrenheit to Centigrade Relationships
Example: Convert 211 oF to degrees Celsius.
Example: Express 548 K in Celsius degrees.

25. Precision and Accuracy
how closely measured values agree with the correct value
Precision
how closely individual measurements agree with each other
Precise
Accurate
Both
Neither

26. Mathematics in Chemistry
Exact numbers (counted numbers)
1 dozen = 12 things
Measured Numbers
Use rules for significant figures
Use scientific notation when possible
Significant figures
digits in a measured quantity that reflect the accuracy of the measurement
-in other words, digits believed to be correct by the person making the measurement
Exact numbers have an infinite number of significant figures
= 1 dozen

27. Significant figures (numbers/digits)
Why use significant numbers?
-Calculators give 8+ numbers
-People estimate numbers differently
-Dictated by the precision (graduation) on your measuring device
-In the lab, the last significant digit is the digit you (the scientist) estimate
Scientists have develop rules to help determine which digits are “significant”

28. Rules for Significant Figures
1. All Nonzero numbers are significant!!!
2. Leading zeroes are never significant
3. Imbedded zeroes are always significant
3.0604
4. Trailing zeroes may be significant
- You must specify significance by how the number is determined or even written
1300 nails - counted or weighed?
–How many significant figures?

29. Significant Figures Multiplication & Division rule:
The product retains the number of significant figures that corresponds to the multiplier with the smallest number of significant figure (sig. fig.)

30. Significant Figures Addition & Subtraction rule:
Answer retains the smallest decimal place value of the addends.

31. Locating the decimal and deciding when to count the zeros!!!
Scientific Notation
Express answers as powers of 10 by moving the decimal place right (-) or left (+)
Use of scientific notation is to remove doubt in the Significant Figures:
2000  2 x 103
15000  1.5 x 10?
0.004  4 x 10-3
 __.__ x 10?
In scientific notation, zeros are given if they are significant!!!
1.000 x 103 has 4 significant figures
2.40 x 103 has ? significant figures
Key to Sig. Figs…
Locating the decimal and deciding when to
count the zeros!!!

32. Review #2 Units of Measure Heat vs. Temperature Precision vs. Accuracy
-length
-volume
-time
-mass
-weight
Heat vs. Temperature
-three temperature scales
-temperature conversions
Precision vs. Accuracy
Significant Figures
Scientific Notation

33. Conversion Factors Length Volume See Text for more conversion factors
1 m = inches
2.54 cm = 1 inch
Volume
1 liter = 1.06 qt
1 qt = liter
See Text for more conversion factors

34. Conversion Factors Why do conversions?
-Scientists often must convert between units
Conversion factors can be made for any relationship of units
-Use known equivalence to make a fraction that can be used to “convert” from one unit to the other

35. Dimensional Analysis 1 inch = 2.54 cm
Use the ratio to perform a calculation so the units will “divide out”
Example: Convert 60 inches to centimeters

36. Dimensional Analysis Example: Express 9.32 yards in millimeters.
3 ft = 1 yard
1 ft = 12 in or
1 in = 2.54 cm
100 cm = 1 m
1000 mm= 1 m

37. Dimensional Analysis Example: Express 627 milliliters in gallons.
1 liter = 1.06 qt
1 qt = liter

38. Practice on your Own 1kg = 2.20 lbs Convert 25 g to lbs
Convert 1 mL to Liters
Convert 20 meters to cm

39. Dimensional Analysis Area = length x width
Area is two dimensional thus units must be in squared terms:
Express: 2.61 x 104 cm2 in ft2

40. Volume =length x width x height
Dimensional Analysis
Volume =length x width x height
Volume is three dimensional thus units must be in cubic terms
Express: 2.61 ft3 in cm3
this volume is used in medical measurements--cc

41. Percentage Percentage is parts per hundred of a sample % = x100
Example: A 335 g sample of ore yields 29.5 g of iron. What is the percent of iron in the ore?
g of substance
total g of sample