1. An Introduction to Chemistry Chapter 1
Hein and Arena
Eugene Passer Chemistry Department Bronx Community College
© John Wiley and Sons, Inc
Version 2.0
12th Edition

2. Chapter Outline 1.1 Why Study Chemistry? 1.2 The Nature of Chemistry
1.3 Thinking Like A Chemist
1.4 A Scientific Approach to Problem Solving
1.5 The Scientific Method
1.6 The Particulate Nature of Matter
1.7 The Physical States of Matter
1.8 Classifying Matter

3. What is a Science?
The observation, identification, description, experimental investigation, and theoretical explanation of natural phenomena.

4. Natural Phenomena observe describe identify experimentally investigate
theoretically explain

5. Chemistry
The science of the composition, structure, properties and reactions of matter, especially of atomic and molecular systems.

6. composition
structure
Matter
properties
reactions

7. 1.4 A Scientific Approach to Problem Solving

8. Problem Solving
Define the problem by recognizing it and stating it clearly.
In science this is called an observation.
2. Propose solutions to the problem.
In science this is called a hypothesis.
3. Decide the best way to solve the problem.
In science we perform an experiment.

9. 1.5 The Scientific Method

10. Definitions

11. Hypothesis: A tentative explanation of certain facts that provide a basis for further experimentation.
Theory: Well-established hypothesis. An explanation of the general principles of certain phenomena with considerable evidence or facts to support it.
Law: Statement of natural phenomena to which no exceptions are known under the given conditions. A law is not an explanation.

12. Outline

14. Steps

15. 1. Collect facts or data that are relevant to the problem or question at hand. This is usually done by experimentation.
2. Analyze the data to find trends (regularities).
3. Formulate a hypothesis that will account for the data and that can be tested by further experimentation.

16. 4. Plan and do additional experiments to test the hypothesis.
5. Modify the hypothesis to ensure compatibility with the experimental data.

17. EXPLANATIONS

19. 1.6 The Particulate Nature of Matter

20. Matter is anything that has mass and occupies space.
Matter can be invisible.
Air is matter, but it cannot be seen.
Matter appears to be continuous and unbroken.
Matter is actually discontinuous. It is made up of tiny particles call atoms.

21. An apparently empty test tube is submerged, mouth downward in water
An apparently empty test tube is submerged, mouth downward in water. Only a small volume of water rises into the tube, which is actually filled with invisible matter–air.
1.3

22. 1.7 Physical States of Matter

23. SOLIDS
Shape
Definite - does not change. It is independent of its container.
Volume
Definite
Particles
Particles are close together. They cling rigidly to each other.
Compressibility
Very slight–less than liquids and gases.

24. A solid can be either crystalline or amorphous
A solid can be either crystalline or amorphous. Which one it is depends on the internal arrangement of the particles that constitute the solid.

25. LIQUIDS Shape Not definite - assumes the shape of its container.
Volume
Definite
Particles
Particles are close together.
Particles are held together by strong attractive forces. They stick firmly but not rigidly to each other.
They can move freely throughout the volume of the liquid.
Compressibility
Very slight–greater than solids, less than gases.

26. GASES Shape No fixed shape. Volume Indefinite. Particles
Particles are far apart compared to liquids and solids.
Particles move independently of each other.

27. GASES Compressibility
The actual volume of the gas particles is small compared to the volume of space occupied by the gas.
Because of this a gas can be compressed into a very small volume or expanded almost indefinitely.

28. ATTRACTIVE FORCES Solid Attractive forces are strongest in a solid.
These give a solid rigidity.
Liquid
Attractive forces are weaker in liquids than in solids.
They are sufficiently strong so that a liquid has a definite volume.

29. ATTRACTIVE FORCES Gas Attractive forces in a gas are extremely weak.
Particles in the gaseous state have enough energy to overcome the weak attractive forces that hold them together in liquids or solids.
Because of this the gas particles move almost independently of each other.

32. 1.8 Classifying Matter

33. Matter refers to all of the materials that make up the universe.

34. Substance
A particular kind of matter that has a fixed composition and distinct properties.
Examples
ammonia, water, and oxygen.

35. Homogeneous Matter Examples
Matter that is uniform in appearance and with uniform properties throughout.
Examples
ice, soda, pure gold

36. Heterogeneous Matter Examples
Matter with two or more physically distinct phases present.
Examples
ice and water, wood, blood

37. Homogeneous
Heterogeneous

38. Phase
A homogenous part of a system separated from other parts by physical boundaries.
Examples
In an ice water mixture, ice is the solid phase and water is the liquid phase.

39. Mixture
Matter containing 2 or more substances that are present in variable amounts. Mixtures are variable in composition. They can be homogeneous or heterogeneous.

40. Homogeneous Mixture (Solution)
A homogeneous mixture of 2 or more substances. It has one phase.
Example
Sugar and water. Before the sugar and water are mixed, each is a separate phase. After mixing the sugar is evenly dispersed throughout the volume of the water.

41. Heterogeneous Mixture
A heterogeneous mixture consists of 2 or more phases.
Example
Sugar and fine white sand. The amount of sugar relative to sand can be varied. The sugar and sand each retain their own properties.

42. Heterogeneous Mixture
A heterogeneous mixture consists of 2 or more phases.
Example
Iron (II) sulfide (FeS) is 63.5% Fe and 36.5% S by mass.
Mixing iron and sulfur in these proportions does not form iron (II) sulfide. Two phases are present: a sulfur phase and an iron phase.
If the mixture is heated strongly a chemical reaction occurs and iron (II) sulfide is formed.
FeS is a compound of iron and sulfur and has none of the properties of iron or sulfur.

43. Heterogeneous Mixture
solid phase1
solid phase2
liquid phase

44. Mixture of iron and sulfur
Compound of iron and sulfur
Formula
Has no definite formula: consists of Fe and S.
FeS
Composition
Contains Fe and S in any proportion by mass.
63.5% Fe and 36.5% S by mass.
Separation
Fe and S can be separated by physical means.
Fe and S can be separated only by chemical change.

45. Heterogeneous Mixture of One Substance
A pure substance can exist as different phases in a heterogeneous system.
Example
Ice floating in water consists of two phases and one substance. Ice is one phase, and water is the other phase. The substance in both cases is the same.

46. System Examples The body of matter under consideration.
In an ice water mixture, ice is the solid phase and water is the liquid phase. The system is the ice and water together.

47. Classification of matter: A pure substance is always homogeneous in composition, whereas a mixture always contains two or more substances and may be either homogeneous or heterogeneous.
1.6

48. The End