1. Introduction: Matter & Measurement
CHEMISTRY The Central Science
Introduction: Matter & Measurement
Cpt. 1 and 2

2. What is Chemistry? The central science
The study of the matter, its composition, properties, and the changes it undergoes.

3. The Molecular Perspective of Chemistry
Matter is the physical material of the universe. It has mass and occupies space.
On the microscopic level, matter consists of atoms and molecules.
Atoms combine to form molecules.
Molecules may consist of the same type of atoms or different types of atoms.

4. The Molecular Perspective of Chemistry

5. Why Study Chemistry?
Chemistry is central to our understanding of other sciences.
Chemistry is also encountered in everyday life.

6. States of Matter
Matter can be a gas, a liquid, or a solid.
Gases have no fixed shape or volume.
Gases can be compressed to form liquids.
Liquids have no shape, but they do have a defined volume. Liquids are fluid.
Solids are rigid and have a definite shape and volume.

7. The Three States of Matter

8. Elements and compounds are referred to as Pure Substances
Elements and compounds are referred to as Pure Substances. They contain one type of particle.
Mixtures are not considered pure substances. They contain more than one type of particle.

9. Pure Substances and Mixtures

10. Elements
If a pure substance cannot be decomposed into something else, then the substance is an element.
There are 118 elements known.
Each element is given a unique chemical symbol (one or two letters).
Elements are building blocks of matter.

11. Elements

12. Elements
Chemical symbols with one letter have that letter capitalized (e.g., H, B, C, N, etc.)
Chemical symbols with two letters have only the first letter capitalized (e.g., He, Be).
C Cu Na U

13. Compounds
Most elements interact to form compounds.
Law of Constant Composition (or Law of Definite Proportions):
The composition of a pure compound is always the same.

14. Substances
If water is decomposed, then there will always be twice as much hydrogen gas formed as oxygen gas.
Pure substances that can be decomposed are compounds.

15. Mixtures
Heterogeneous mixtures are not uniform throughout.
Homogeneous mixtures are uniform throughout.
Homogeneous mixtures are called solutions.

17. Physical and Chemical Properties
Intensive physical properties do not depend on how much of the substance is present.
Examples: density, temperature, color, melting point.
Extensive physical properties depend on the amount of substance present.
Examples: mass, volume, pressure.
Chemical Properties describe the chemical behavior of a substance
-Example: Flammable, corrodes metal, basic

18. Physical and Chemical Changes
When a substance undergoes a physical change, its physical appearance changes.
Ice melts: a solid is converted into a liquid.
Physical changes do not result in a change of composition.
When a substance undergoes a chemical change, it changes its composition:
When pure hydrogen and pure oxygen react completely, they form pure water.

19. Chemical Changes

20. Separation of Mixtures
Mixtures can be separated if their physical properties are different.
Solids can be separated from liquids by means of filtration.
Homogeneous liquid mixtures can be separated by distillation.

21. Distillation

22. Separation of Mixtures
Chromatography can be used to separate mixtures that have different polarity.
Chromatography can be used to separate the different colors of inks in a pen.

23. Gas Chromatography

24. The Periodic Table

25. Periodic Table Basics Most elements are metals – left of zig-zag
Non metals are on the right of zig-zag
Borderline elements called metalloids
Horizontal rows are called periods
Vertical columns are called groups or families. (similar properties – more later)

26. Metals and Nonmetals and Their Ions
Good conductors of heat and electricity.
Malleable and ductile.
Moderate to high melting points.
Nonmetals
Nonconductors of heat and electricity.
Brittle solids.
Some are gases at room temperature.

27. There are two types of units:
SI Units
There are two types of units:
fundamental (or base) units;
derived units.
There are 7 base units in the SI system.

28. Powers of ten are used for convenience with smaller or larger units in the SI system.

29. SI Units

30. There are three temperature scales: 1. Kelvin Scale
SI Units
Note the SI unit for length is the meter (m) whereas the SI unit for mass is the kilogram (kg).
1 kg weighs lb.
Temperature
There are three temperature scales:
1. Kelvin Scale
Used in science.
Same temperature increment as Celsius scale.
Lowest temperature possible (absolute zero) is zero Kelvin.
Absolute zero: 0 K = -273 oC.

31. Temperature 2. Celsius Scale 3. Fahrenheit Scale Also used in science.
Water freezes at 0 oC and boils at 100 oC.
To convert: K = oC
3. Fahrenheit Scale
Not generally used in science.
Water freezes at 32 oF and boils at 212 oF.
To convert:

32. Temperature

33. Scientific Notation
Numbers written in scientific notation include a numeral with one digit before the decimal point, multiplied by some power of ten (6.022 x 1023)
In scientific notation, all digits are significant.
You should be able to convert from non-scientific notation to scientific and vice-versa (on calc. as well).

34. Examples Convert to scientific notation: a. 450 000 000b
Convert to standard notation:
a x 10-6
b x 105

35. Derived Units
Derived units are obtained from the 7 base SI units.
Example:

36. The units for volume are given by (units of length)3.
SI unit for volume is 1 m3.
Common: 1 mL=1 cm3.
Other volume units:
1 L = 1 dm3 = 1000 cm3 = 1000mL.

37. Volume

38. Density
Used to characterize substances.
Defined as mass divided by volume:
Units: g/cm3 or g/mL
Originally based on mass (density was defined as the mass/vol. of 1.00 g of pure water).

39. Examples:
An object with a mass of g occupies a volume of 11.8 mL. What is its density?
A sample with a density of 3.75 g/cm3 occupies a volume of cm3. What is the mass of the sample?
A graduated cylinder is filled with 15.0 cm3 of water. An object with a mass of g causes the total volume to increase to 23.4 mL. What is the density of the sample?
1.52g/mL g g/ml

40. Dimensional Analysis
Method of calculation utilizing a knowledge of units. Conversion factors are simple ratios (fractions):

41. Using a Conversion Factor
Example: convert length in meters to length in centimeters:
3.25 meters x 100 cm = cm
1 m

42. Using Two or More Conversion Factors
Example to convert length in meters to length in inches:
3.00 meters x 100 cm x 1 inch = in
1 m cm

43. Problem solving
In dimensional analysis always ask three questions:
What data are we given?
What quantity do we need?
What conversion factors are available to take us from what we are given to what we need?

44. Uncertainty in Measurement
All scientific measures are subject to error.
These errors are reflected in the number of figures reported for the measurement.
Precision and Accuracy
Measurements that are close to the “correct” value are accurate.
Measurements that are close to each other are precise.

45. Precision and Accuracy

46. Significant Figures
The number of digits reported in a measurement reflect the accuracy of the measurement and the precision of the measuring device.
All the figures known with certainty plus one extra figure (estimated digit) are called significant figures.
In any calculation, the results are reported to the fewest significant figures (for multiplication and division) or fewest decimal places (addition and subtraction).

47. Significant Figures Rules:
Non-zero numbers are always significant.
Zeros between digits are always significant.
Initial zeros are not significant. (Example: has only 1 sf.)
Zeros at the end of the number are significant IF there is a decimal point.(24.0 cm has 3 sf)

48. Examples: How many significant figures are in each of the following?kg s
507 people
230,050 cm A

49. Reporting Uncertainties and Errors in Measurements
All instruments have limitations. When recording a measurement, you should report the measurement to the limits of the instrument.
The uncertainty (or tolerance) of an instrument may be found on the instrument ( given as ±) , or you may be told what it is.

50. Percent Error – used to compare data (observed values) to accepted values in a meaningful way
│ Obs – Acc│ x = % error
Acc

51. Direct Inverse Proportions Proportions