01 INTRODUCTION

Slide 1 :Chapter 01 Introduction


Slide 2 :Chemistry Chemistry - is the science that deals with the composition, structure, and properties of substances and with the transformations that they undergo Matter – anything that has mass and occupies space Energy – ability to do work to accomplish change Chemistry: Methods and Measurement 1.1 The Discovery Process


Slide 3 :Major Areas of Chemistry Biochemistry Organic Chemistry Inorganic Chemistry Analytical Chemistry Physical Chemistry 1.1 The Discovery Process


Slide 4 :Scientific Method Scientific Method - is a systematic approach to research Hypothesis – is a tentative explanation of a set of observations Theory – is a well substantiated explanation of some aspect of the natural world. Law - is a generalization that describes recurring facts or events in nature 1.1 The Discovery Process


Slide 5 :1.2 Matter and Properties Matter and Physical Properties Physical Change – produces a recognizable difference in the appearance of a substance without causing any change in its composition and identity. (ice to liquid water to water vapor). - Usually change is reversible. Physical Property – can be observed without changing the composition or identity of a substance (Boiling Point, Melting Point)


Slide 6 :1.2 Matter and Properties Matter and Chemical Properties Chemical Properties – properties of matter that need to change the composition of the substance to be observed. (flammability) Chemical Change – the identities of substances change and new substances form. - Change is irreversible. Example 1.1 Can the process that takes place when an egg is fried be described as a physical or chemical change? Chemical Change


Slide 7 :1.2 Matter and Properties Intensive and Extensive Properties Intensive Property – does not depend on the amount of the sample (density, boiling and melting points, and specific gravity) Extensive Property – depend on the amount of the sample (mass and volume) Example 1.3 Is temperature an extensive or intensive property? Intensive property


Slide 8 :1.2 Matter and Properties Classification of Matter water, ammonia, table sugar, Au, O2 air, soft drinks, milk, and cement salt dissolved in water sand + Fe filings Example 1.3 Is sugar in water a pure substance, a homogenous mixture, or a heterogeneous mixture? Homogenous Mixture


Slide 9 :1.3 Measurement in Chemistry Data, Results, and Units Data – individual result of a single experiment or observation. (mass, length, volume, time, temperature) Result – outcome of an experiment. Example 1.4 A drug is less stable if moisture is present, and excess moisture can hasten the breakdown of the active ingredient , leading to loss of potency. Therefore we may wish to know how much water a certain quantity of a drug gains when exposed to air. To do this experiment, we first weigh the drug, then expose it to the air for a period reweigh it. Which are the data and which are the results? Data: initial and final weights Results: difference in weight and the conclusions based on the observed change in weight.


Slide 10 :1.3 Measurement in Chemistry English and Metric Units Convert 1 foot to miles 1 foot = 12 inches = 0.33 yard = 1/5280 mile Convert 1 meter to milimeters 1 meter = 10 decimeter = 100 centimeters = 1000 millimeters Unit English Metric Mass pound (lb) gram (g) Length yard (yd) meter (m) Volume gallon (gal) liter (l)


Slide 11 :1.3 Measurement in Chemistry Table 1.1 Some Common Prefixes Used in the Metric System


Slide 12 :1.3 Measurement in Chemistry


Slide 13 :1.3 Measurement in Chemistry


Slide 14 :1.3 Measurement in Chemistry


Slide 15 :1.3 Measurement in Chemistry Unit Conversion: English and Metric Systems Conversion factor Factor label method or dimensional analysis Conversion of Units Within the Same System Table 1.2 Some Common Relationships Used in the English System A. Mass 1 pound = 16 ounces 1 ton = 2000 pounds B. Length 1 foot = 12 inches 1 yard = 3 feet 1 mile = 5280 feet C. Volume 1 gallon = 4 quarts 1 quart = 2 pints 1 quart = 2 pints Example 1.5 Convert 12 gallons to units of quarts 48 qt Example 1.6 Convert 10.0 centimeters to meters 0.100 m


Slide 16 :1.3 Measurement in Chemistry Conversion of Units from One System to Another Table 1.3 Commonly Used “Bridging” Units for Intersystem Conversions A. Mass 1 pound = 454 grams 2.2 pounds = 1 kilogram B. Length 1 inch = 2.54 centimeters 1 yard = 0.91 meter C. Volume 1 quart = 0.946 meter 1 gallon = 3.78 liters Example 1.7 Convert 4.00 ounces to kilograms. 0.114 kg Example 1.8 Convert 1.5 meters2 to centimeters2 1.5 x 104 cm2


Slide 17 :1.4 Significant Figures and Scientific Notation Significant Figures Ruler Object A Object B Ruler II: 4.4 cm 6.7 cm Ruler I: 4.45 cm 6.79 cm


Slide 18 :1-4 Significant Figures and Scientific Notation The significant figures (also called significant digits) of a number - meaningful digits in a measured quantity or calculated quantity. - indicates the accuracy of a measurement. - more significant numbers higher accuracy of measurement. Rules in Determining the Number of Significant Figures Zeros appearing between nonzero digits are significant, for example: 60.8 has three significant figures 39008 has five significant figures Zeros appearing in front of nonzero digits are not significant, for example: 0.093827 has five significant figures 0.0008 has one significant figure 0.012 has two significant figures Zeros at the end of a number and to the right of a decimal are significant, for example: 35.00 has four significant figures 8,000.000000 has ten significant figures Zeros at the end of a number without a decimal point may or may not be significant, and are therefore ambiguous, for example: 1,000 could have between one and four significant figures.


Slide 19 :1-4 Significant Figures and Scientific Notation Number 6.29 g 0.00348 g 9.0 1.0  10-8 100 eggs 100 g  = 3.14159 Significant Figures 3 3 2 2 infinite bad notation various Rounding Off ≥5, round up <5, round down


Slide 20 :1-4 Significant Figures and Scientific Notation Adding and Subtracting with Significant Figures Use the number of decimal places in the number with the fewest decimal places. 1.14 0.6 11.676 13.416  13.4 Multiplying and Dividing with Significant Figures Use the fewest significant figures 0.01208  0.236 = 0.05118644 = 0.0512


Slide 21 :1-4 Significant Figures and Scientific Notation Exercise 1-5 Find the number of significant figures in the following measurements 52.0 ml 52 ml 5200 ml 520 ml 0.000520 ml Exercise 1-6 Round off the following numbers to three significant figures. 10.071 0.008695 51,428


Slide 22 :1-4 Significant Figures and Scientific Notation Exercise 1-7 Perform the following mathematical operation. Give the answer with the correct number of significant figures. 11.73g + 6.8g +120g = _____ 150 ml – 6.8 ml = _____ 2.6 cm x 5.2 cm x 11.1 cm = _____ 8.238 g ÷ 0.92 mL = _____


Slide 23 :1-4 Significant Figures and Scientific Notation Scientific Notation 0.0024 0.0180 224 Error, Accuracy, Precision, and Uncertainty Error – difference between the true value and our estimation, or measurement of the value. Types of Errors Random Errors – causes data from multiple measurements of the same quantity to be scattered in a more or less uniform way around some average value. - Inherent in the experimental approach to the study matter and its behavior Systematic Errors – causes data to be either smaller or larger than the accepted value. Can be found and, in many cases, removed or corrected Uncertainty – degree of doubt in a single measurement


Slide 24 :1-4 Significant Figures and Scientific Notation Accuracy tells us how close a measurement is to the true value of the quantity that was measured. Precision refers to how closely two or measurements of the same quantity agree with one another.


Slide 25 :1.5 Experimental Quantities Mass Mass – describes the quantity of matter in an object Weight – the force of gravity on an object Balances – are instruments used to measure the mass of materials Atomic mass unit (amu) – convenient way of representing the mass of particles on the atomic level 1 amu = 1.661 x 10-24 g


Slide 26 :Volume Volume – space occupied by an object Standard metric unit of volume is the liter 1.5 Experimental Quantities


Slide 27 :1.5 Experimental Quantities Time the standard metric unit of time is the second.


Slide 28 :1.5 Experimental Quantities Example 1.16 Normal Body Temperature is 98.6 °F. Calculate the corresponding temperature in degrees Celsius and Kelvin. 37.0 °C 310.2 K Temperature – how hot or how cold something is. - measure of the kinetic energy of a sample of matter. Units of Temperature Celsius, °C Fahrenheit, °F Kelvin, K


Slide 29 :1.5 Experimental Quantities Energy – ability to do work or cause change - measure of the kinetic energy of a sample of matter. Types of Energy Kinetic Energy – energy due to motion Potential Energy – energy due to position Forms of Energy Light Heat Electrical Mechanical Chemical Characteristics of Energy In chemical reactions, energy cannot be created or destroyed Energy may be converted from one form to another Conversion of energy from one from to another always occurs with less than 100% efficiency.


Slide 30 :1.5 Experimental Quantities Units of Heat Energy 1 calorie (cal) = 4.18 joules (J) Calorie – is the amount of heat energy required to increase the temperature of 1 gram of water to 1°C. Concentration Concentration – is a measure of the number of particles of a substance, or the mass of those particles, that are contained in a specified volume.


Slide 31 :Which is more dense? a b Density is the ratio of mass to volume


Slide 32 :1.5 Experimental Quantities Common Units of Density g/mL = g/cm3 = g/cc Example 1.17 2.00 cm3 of aluminum are found to weigh 5.40 g. Calculate the density of aluminum in units of g/cm3. 2.70 g/cm3 Example 1.18 Air has a density of 0.0013 g/mL. What is the mass of a 6.0 L sample of air? 7.8 g air Example 1.19 Calculate the mass, in grams, of 10.0 mL of mercury (Hg) if the density of mercury is 13.6 g/mL. 136 g Hg Example 1.20 Calculate the volume, in milliliters, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 grams. 4.17 mL liquid


Slide 33 :1.5 Experimental Quantities Specific Gravity (SG)- is the ratio of the density of a given substance to the density of water. Dimensionless quantity if SG value is greater than 1 (Aluminum = 2.7) are denser than water, thus will sink in water. if SG value is lesser than 1 (Wood = 0.2) are less dense than water, thus will float on it.