1. Chapter 1 Introduction: Matter & Measurement
CHEMISTRY The Central Science 10th Edition
Chapter 1 Introduction: Matter & Measurement

2. Why Study Chemistry
Chemistry is the study of the properties of materials and the changes that materials undergo.
Chemistry is central to our understanding of other sciences.
Chemistry is also encountered in everyday life.

3. Chemistry: Catastrophe Prevention?
The space shuttle Columbia disintegrated in 2003 upon reentry into the Earth’s atmosphere due to a damaged thermal protection system.

4. The Molecular Perspective of Chemistry
The Study of Chemistry
The Molecular Perspective of Chemistry
Matter is the physical material of the universe.
Matter is made up of relatively few elements.
On the microscopic level, matter consists of atoms and molecules.
Atoms combine to form molecules.
As we see, molecules may consist of the same type of atoms or different types of atoms.

5. Molecular Perspective of Chemistry

6. Classification of Matter
States of Matter
Matter can be a gas, a liquid, or a solid.
These are the three states of matter.
Gases take the shape and volume of their container.
Gases can be compressed to form liquids.
Liquids take the shape of their container, but they do have their own volume.
Solids are rigid and have a definite shape and volume.

7. Pure Substances and Mixtures
Classification of Matter
Pure Substances and Mixtures
Elements consist of a unique type of atom.
Molecules can consist of more than one type of element.
Molecules that have only one type of atom (an element).
Molecules that have more than one type of atom (a compound).
If more than one atom, element, or compound are found together, then the substance is a mixture.

8. Pure Substances and Mixtures

9. Classification of Matter
Pure Substances and Mixtures
If matter is not uniform throughout, then it is a heterogeneous mixture.
If matter is uniform throughout, it is homogeneous.
If homogeneous matter can be separated by physical means, then the matter is a mixture.
If homogeneous matter cannot be separated by physical means, then the matter is a pure substance.
If a pure substance can be decomposed into something else, then the substance is a compound.

10. Classification of Matter
Elements
If a pure substance cannot be decomposed into something else, then the substance is an element.
There are 114 elements known.
Each element is given a unique chemical symbol (one or two letters).
Elements are building blocks of matter.
The earth’s crust consists of 5 main elements.
The human body consists mostly of 3 main elements.

11. Classification of Matter
Elements

12. Classification of Matter
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).

13. Classification of Matter
Compounds
Most elements interact to form compounds.
Example, H2O
The proportions of elements in compounds are the same irrespective of how the compound was formed.
Law of Constant Composition (or Law of Definite Proportions):
The composition of a pure compound is always the same.

14. Classification of Matter
Compounds
If water is decomposed, then there will always be twice as much hydrogen gas formed as oxygen gas.
Pure substances that cannot be decomposed are elements.

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

17. Physical vs. Chemical Properties
Properties of Matter
Physical vs. Chemical Properties
Physical properties can be measure without changing the basic identity of the substance (e.g., color, density, odor, melting point)
Chemical properties describe how substances react or change to form different substances (e.g., hydrogen burns in oxygen)
Intensive physical properties do not depend on how much of the substance is present.
Examples: density, temperature, and melting point.
Extensive physical properties depend on the amount of substance present.
Examples: mass, volume, pressure.

18. Physical and Chemical Changes
Properties of Matter
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 changes its composition, it undergoes a chemical change:
When pure hydrogen and pure oxygen react completely, they form pure water. In the flask containing water, there is no oxygen or hydrogen left over.

19. Physical and Chemical Changes
Properties of Matter
Physical and Chemical Changes

20. Separation of Mixtures
Properties of Matter
Separation of Mixtures
Mixtures can be separated if their physical properties are different.
Solids can be separated from liquids by means of filtration.
The solid is collected in filter paper, and the solution, called the filtrate, passes through the filter paper and is collected in a flask.

21. Separation of Mixtures
Properties of Matter
Separation of Mixtures
Homogeneous liquid mixtures can be separated by distillation.
Distillation requires the different liquids to have different boiling points.
In essence, each component of the mixture is boiled and collected.
The lowest boiling fraction is collected first.

22. Separation of Mixtures

23. Separation of Mixtures
Units of Measurement
Separation of Mixtures
Chromatography can be used to separate mixtures that have different abilities to adhere to solid surfaces.
The greater the affinity the component has for the surface (paper) the slower it moves.
The greater affinity the component has for the liquid, the faster it moves.
Chromatography can be used to separate the different colors of inks in a pen.

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

25. Units of Measurement
Base SI Units

26. Selected Prefixes used in SI System
Units of Measurement
SI Units
Selected Prefixes used in SI System

27. Class Practice Examples
What is the name given to the unit that equals (a) 10-9 grams; (b) 10-6 second; (c) 10-3 meter
What fraction of a meter is a nanometer?

28. Units of Measurement 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:
Kelvin Scale
Used in science.
Same temperature increment as Celsius scale.
Lowest temperature possible (absolute zero) is zero Kelvin.
Absolute zero: 0 K = oC.

29. Units of Measurement Temperature Celsius Scale Fahrenheit Scale
Also used in science.
Water freezes at 0 oC and boils at 100 oC.
To convert: K = oC
Fahrenheit Scale
Not generally used in science.
Water freezes at 32 oF and boils at 212 oF.
To convert:

30. Class Practice Example
Make the following temperature conversions: (a) 68 oF to oC; (b) oC to oF

31. Units of Measurement
Temperature

32. Units of Measurement Derived Units
Derived units are obtained from the 7 base SI units.
Example:

33. Units of Measurement Volume
The units for volume are given by (units of length)3.
SI unit for volume is 1 m3.
We usually use 1 mL = 1 cm3.
Other volume units:
1 L = 1 dm3 = 1000 cm3 = 1000 mL.

34. Units of Measurement
Volume

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

36. Class Practice Examples
Answer the following problems:
(a) Calculate the density of mercury if 1.0 x 102 g occupies a volume of 7.36 cm3.
(b) Using the density for mercury, calculate the mass of 65.0 cm3 of mercury.

37. Uncertainty in Measurement
All scientific measures are subject to error.
These errors are reflected in the number of figures reported for the measurement.
These errors are also reflected in the observation that two successive measures of the same quantity are different.
Precision and Accuracy
Measurements that are close to the “correct” value are accurate.
Measurements that are close to each other are precise.

38. Precision and Accuracy

39. Uncertainty in Measurement
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 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).

40. Uncertainty in Measurement
Significant Figures
Non-zero numbers are always significant.
Zeros between non-zero numbers are always significant.
Zeros before the first non-zero digit are not significant. (Example: has one significant figure.)
Zeros at the end of the number after a decimal place are significant.
Zeros at the end of a number before a decimal place are ambiguous (e.g. 10,300 g).

41. Dimensional Analysis
Method of calculation utilizing a knowledge of units.
Given units can be multiplied or divided to give the desired units.
Conversion factors are used to manipulate units:
Desired unit = given unit  (conversion factor)
The conversion factors are simple ratios:

42. Using Two or More Conversion Factors
Dimensional Analysis
Using Two or More Conversion Factors
Example to convert length in meters to length in inches:

43. Class Practice Problem
A person’s height is measured to be in. What is this height in centimeters?
Perform the following conversions: (a) 2 days to s; (b) 20 Kg to g.

44. Using Two or More Conversion Factors
Dimensional Analysis
Using Two or More Conversion Factors
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?