Lecture Presentation. Chapter 1. Introduction: Matter and Measurement. James F. Kirby Quinnipiac University Hamden, CT

Transcription

1 Lecture Presentation Chapter 1 Introduction: Matter and Measurement James F. Kirby Quinnipiac University Hamden, CT

2 Why do we study chemistry? My parents want me to study chemistry. I need to graduate. For fun. I do not know. It has a considerable impact on society. It s part of your curriculum. Chemistry is the central science. 2

3 Chemistry Is the study of properties of materials and changes that they undergo. The Atomic and Molecular Perspective of Chemistry: we study matter, its properties, and its behavior. 3

4 Matter We define matter as anything that has mass and takes up space. Give me some examples of matter. Book, pen, mountain, chemistry building, iphone,.. 4

5 States of Matter Gas no fixed volume or shape conforms to shape of the container Liquid compressible volume independent of container no fixed shape incompressible Solid volume and shape independent of container rigid incompressible 5

6 Atom, Element and Compound Atoms are the building blocks of matter. Each element is made of the same kind of atom. A compound is made of two or more different kinds of elements. 6

7 Pure Substance and Mixtures Pure substance: matter that has distinct properties and composition that does not vary from sample to sample Water (H 2 O) and table salt (NaCl) Mixtures: two or more substances combined together in which each substance retains its own chemical identity Heterogenous mixtures have inconsistent, non-uniform composition Example: Salad, jelly beans Homogenous mixtures (solutions) have consistent, uniform composition; cannot distinguish individual substances Example: sweet tea, coffee 7

8 Classification of Matter 8

9 Test: Classification of Matter Classify each of following as a pure substance or a mixture: table salt, oxygen, distilled water, air, milk shake 9

10 Properties and Changes of Matter 10

11 Properties of Matter Physical properties: can be observed without changing the identity and composition of the substance Color, odor, density, melting and boiling points, etc. Chemical properties: Can only be observed when a substance is changed into another substance. Flammability, corrosiveness, reactivity with acid, etc. Intensive Properties: independent of the amount of the substance that is present (useful in identifying substances) Density, boiling point, color, etc. Extensive Properties: depend upon the amount of the substance present (relates to the amount of substance present) Mass, volume, energy, etc. 11

12 Types of Changes Substances can undergo two types of changes: physical and chemical Physical change: physical appearance of a substance changes with NO CHANGE in chemical composition (same substance before and after the change, chemically) Changes of states, temperature, volume, etc. Chemical change: a substance is transformed into a chemically different substance, which result in new substances Combustion, decomposition, oxidation, etc. Chemical change is also called a chemical reaction 12

13 Example of Physical Change Ice is solid H 2 O. As ice melts, it becomes liquid H 2 O. But H 2 O is still H 2 O! There is no change in chemical composition. Only the physical properties of the substance have changed. Similarly, when the liquid H 2 O evaporates, it becomes vapor or gaseous H 2 O, which is still H 2 O. 13

14 Example of Chemical Reactions In the course of a chemical reaction, the reacting substances are converted to new substances. 14

15 Sample Problem Physical change or chemical change? 15

16 Separating Mixtures Mixtures can be separated based on physical properties of the components of the mixture. Some methods used are filtration. distillation. chromatography.

17 Filtration Separates components of a mixture based upon differences in particle size. Most often used to separate a precipitate from a solution. 17

18 Distillation Distillation uses differences in the boiling points of substances to separate a homogeneous mixture into its components. 18

19 Chromatography This technique separates substances on the basis of differences in the ability of substances to adhere to the solid surface, in this case, dyes to paper.

20 Numbers and Chemistry Numbers play a major role in chemistry. Many topics are quantitative (have a numerical value). Concepts of numbers in science Units of measurement Quantities that are measured and calculated Uncertainty in measurement Significant figures Dimensional analysis

21 SI Units SI units: Système International d Unités (International System of Units) A different base unit is used for each quantity Units used are those of the metric system 21

22 Units of Measurement Metric System The base units used in the metric system Mass: gram (g) Length: meter (m) Time: second (s or sec) Temperature: degrees Celsius ( o C) or Kelvins (K) Amount of a substance: mole (mol) Volume: cubic centimeter (cc or cm 3 ) or liter (l)

23 Units of Measurement Metric System Prefixes Prefixes convert the base units into units that are appropriate for common usage or appropriate measure.

24 Measurements in Chemistry The metric system uses a series of prefixes to denote small and large numbers. Name Symbol Multiplier mega M 10 6 kilo k 10 3 hecta h 10 2 deka da 10 Base b 1 deci d 10-1 centi c 10-2 milli m 10-3 micro m 10-6 nano n 10-9 pico p Remember this table! femto f

25 Sample Problem Metric conversion Perform the following metric conversion kg to g? 4.7 L to ML? 38.9 km to pm? 25

26 Temperature Definition: temperature is a measure of the average kinetic energy of the particles in a sample. Temperature scales o Celsius scale o Kelvin scale o Fahrenheit scale Often used in scientific measurement 26

27 Temperature: Celsius In scientific measurements, the Celsius and Kelvin scales are most often used The Celsius scale is based on the properties of water: 0 C is the freezing point of water 100 C is the boiling point of water 27

28 Temperature: Kelvin (K) The Kelvin is the SI unit of temperature It is based on the properties of gases There are no negative Kelvin temperatures K = C Note: its not K, just K (NO degree symbol) 28

29 Temperature: Fahrenheit The Fahrenheit scale is not used in scientific measurements Conversions to KNOW: F = (9/5)( C) + 32 C = (5/9)( F 32) K = C

30 Derived Units - Volume Volume, V = length (L) x Height (H) x Width (W) Units for volume are cubed. Example: cm 3, m 3 Volume is also measured in ml (liquids) KNOW: 1 ml = 1 cm 3 1 L = 1 dm 3 Know how to calculate both volume and density 30

31 Derived Units - Density Density is a physical property of a substance. Density = mass Volume d = m V Units for density: g/cm 3 or g/ml 31

32 Sample Problem - Density What is the mass in grams of a cube of gold (density = g/cm 3 ) if the length of the cube is 2.00 cm?

33 Numbers Encountered in Science Exact numbers are counted or given by definition. For example, there are 12 eggs in 1 dozen. Inexact (or measured) numbers depend on how they were determined. Scientific instruments have limitations. Some balances measure to ±0.01 g; others measure to ±0.0001g.

34 Uncertainty in Measurements Different measuring devices have different uses and different degrees of accuracy. 34

35 Precision and Accuracy Precision (or reproducibility) - how close the individual measurements are to each other. Accuracy - Refers to how close a measurement (or the average of repeated measurements) is (are) to the real value What can you say about the accuracy and precision of the darts of each dartboard? 35

36 Significant Figures The term significant figures refers to the digits that were measured All digits of a measured quantity including the uncertain ones are called significant figures When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers The greater the number of sig. figs, the greater the certainty implied for the measurement 2.2 g and g 36

37 How to Determine Significant Figures 1. All non-zero digits ARE significant 2. Zeroes embedded between non-zero digits ARE significant. (eg: 1005 kg) 3. Leading zeros ARE NOT significant. (eg: 150 km) 4. Zeros at the beginning of a number are NEVER significant; they merely indicate the position of the decimal point. (eg: cm) 5. Zeros at the end of a number are significant if a decimal point is written in the number. (eg: g) 37

38 Rounding Off Numbers Less than 5, the preceding number is unchanged. (eg: to 2 significant figures) Greater or equal to 5, the preceding number is increased by 1. (eg: to 3 significant figures gives) 38

39 Significant Figures Scientific notation makes it easier to determine the correct number of significant digits. 2.4 x 10 2 g 2.40 x 10 2 g x 10 2 g x 10 2 g The number before the decimal point is always a single digit between 0 and 9, inclusive. 39

40 Sample Problem Significant figures How many significant figures are in each of the following numbers (assume that each number is a measured quantity): (a) 4.003, (b) , (c) 5000? 40

41 Sig. Figs in Calculations The correct number of significant figures must be maintained when carrying out mathematical calculations with measurements. The rules depend on what type of calculation is done! Multiplication and division The answer must contain the same number of significant figures as there are in the number being used that has the LEAST significant figures x round off to x round off to

42 Addition and subtraction Sig Figs in Calculations Now we focus on the number of decimal places, not the total number of significant figures! The answer must have the same number of decimal places as the number being used that has the LEAST number of decimal places 83.5 ml ml ml = ml ml ml ml = ml 42

43 Practice problem The measured quantity contains significant figures. a. 3 b. 4 c. 5 d. 6

44 Practice problem 6.03 g g = g. a. 13 b c d

45 Practice problem 7.1 m x 6.03 m = m 2. a. 43 b c d

46 Practice Problem Of the following, is the smallest mass. A. 2.5 x 10 9 fg B. 25 kg C pg D ng E mg Chemistry 1145

47 Dimensional Analysis We use dimensional analysis to convert one quantity to another Most commonly dimensional analysis utilizes conversion factors e.g.: 1 in = 2.54 cm or 1000 mm = 1 m 1 in 2.54 cm or 2.54 cm 1 in 1000 mm 1 m or 1 m 1000 mm 47

48 Dimension Analysis Use the form of the conversion factor that puts the desired unit in the numerator Given unit desired unit given unit desired unit Conversion factor These are called conversion factors or unit factors. 12 in 1 1ft 12 in 1ft Only the units cancel, not the numbers!!! 48

49 Sample Problem Dimensional analysis Express 8.00 meters in inches(1 in. = 2.54 cm) Given unit: m Desired unit: in Conversion factors (what we know): 1 in = 2.54 cm and 1 m = 100 cm How many sig. figs should there be in the answer? Note: Conversion factors do not play a role in determining sig. figs!! 49

50 Sample Problem Dimensional analysis Express 627 milliliters in gallons(1 gal = 3.8 L) Given unit: ml Desired unit: gal Conversion factors (what we know): 1 gal = 3.8 L 1 L = 1000 ml 50

51 Dimensional analysis is your friend! Learn it! Use it! It will take you a LONG way in this course! 51