Chapter 1 Introduction: Matter & Measurement

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

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 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. 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. 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. 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. Properties of Matter Physical and Chemical Changes

20. 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. 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. 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. 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 2.2046 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 = -273.15 oC.

29. Units of Measurement • Temperature • Celsius Scale • Also used in science. • Water freezes at 0 oC and boils at 100 oC. • To convert: K = oC + 273.15. • 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) -36.7 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: 0.0003 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. 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 67.50 in. What is this height in centimeters? • Perform the following conversions: (a) 2 days to s; (b) 20 Kg to g.

44. 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?