Chapter 1. Matter, Measurement and Problem Solving. Chapter 1. Helleur. Principles of Chemistry: A Molecular Approach 1 st Ed.

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1 Why Clickers? to refresh the students during the lecture to assess student understanding of a topic to increase student participation in class to encourage student interaction in large lectures peer teaching to increase student attendance Principles of Chemistry: A Molecular Approach 1 st Ed. Nivaldo Tro Matter, Measurement and Problem Solving 1

2 What Is Chemistry? Chemistry is the study of the composition, structure, and properties of matter and energy, and of the changes that matter undergoes. (show demonstration with NaCl) Matter is anything that occupies space and has mass. Energy is the ability to do work. Experiment and Explanation Experiment and explanation are the heart of chemical research. An experiment is an observation of natural phenomena carried out in a controlled manner so that the results can be duplicated and rational conclusions can be obtained. After a series of experiments, a researcher may see some relationship or regularity in the results. Experiment and Explanation If the regularity or relationship is fundamental and we can state it simply, we call it a law. An example is the law of conservation of mass, which says that, for a chemical reaction, the sum of the masses of the reactants = sum of the masses of the products. example : Mercury + oxygen mercury (II) oxide (HgO) We know that 2.35 g Hg produces 2.73 HgO What is the mass of oxygen that reacted? The law states: Mass of substances before reaction = mass of substances after 2.35 g Hg + x g O g HgO Solve for x x= 0.38 g Problem: Law of Conservation of Mass #1. Zinc metal reacts with yellow crystals of sulfur in a exothermic reaction to produce a white powder of zinc sulfide. If a chemist determines that 65.4 g of Zn reacts with 32.1 g of S, how many grams of ZnS could be produced from 20.0 g of Zn metal? Answer is 29.8 g 2

3 Elements, Compounds, and Mixtures To understand how matter is classified by its chemical constitution we must first look at physical and chemical changes. A physical change is a change in the form of matter but not in its chemical identity. Physical changes are usually reversible. No new compounds are formed during a physical change. Melting ice and dissolving sugar in water are examples of a physical changes. Elements, Compounds, and Mixtures A chemical change, or chemical reaction, is a change in which one or more kinds of matter are transformed into a new kind of matter or several new kinds of matter. Chemical changes are usually irreversible. New compounds are formed during a chemical change. Vinegar, HC 2 H 3 O 3, reacting with baking soda NaHCO 3 is an example of a chemical change how do you know?? Elements, Compounds, and Mixtures A physical property is a characteristic that can be observed for material without changing its chemical identity. Examples are physical states (solid, liquid, or gas), melting point, and color. A chemical property is a characteristic of a material involving its chemical change. A chemical property of iron is its ability to react with oxygen to produce rust. Fe can be oxidized Which of the following represents a chemical change? a. freezing water to make ice cubes b. dry ice evaporating at room temperature c. toasting a piece of bread d. dissolving sugar in hot coffee e. crushing an aluminum can Clicker Question 3

4 Classifying Matter Elements, Compounds and Mixtures Which of the following represents a chemical change? a. freezing water to make ice cubes b. dry ice evaporating at room temperature c. toasting a piece of bread d. dissolving sugar in hot coffee e. crushing an aluminum can Elements, Compounds, and Mixtures A modern form of the periodic table. The elements An element is a substance that cannot be decomposed by any chemical reaction into simpler substances. i.e.,periodic table of Elements; zinc (Zn) The smallest unit of an element is the atom. Most substances are compounds. A compound is a substance composed of two or more elements chemically combined. Zinc sulfide (ZnS) The smallest unit of a compound is the molecule. Presentation of Lecture Outlines,

5 Elements, Compounds, and Mixtures Most of the materials we see around us are mixtures. A mixture is a material that may or may not be be separated by physical means into two or more substances. (sand and water) Mixtures are classified as heterogeneous if they consist of physically distinct parts (sand and water) or homogeneous when the properties are uniform throughout (seawater looks like one substance however is it composed on salt and water) Measurement and Significant Figures Measurement is the comparison of a physical quantity to be measured with a unit of measurement that is, with a fixed standard of measurement. The term precision refers to the closeness of the set of values obtained from identical measurements of a quantity. i.e., measurements such as 23.3, 23.4, 23.2 cm of the same object is precise Accuracy is a related term; it refers to the closeness of a single measurement to its true value. i.e, if the true value is 18.3 cm and the measurement was 23.4 cm, then the measurement is not accurate A student measures the mass of a penny 4 times and records the following data. What can be said about the data if the actual mass of the penny is g? a. The data is both accurate and precise. b. The data is neither accurate nor precise. c. The data is accurate but not precise. d. The data is not accurate but it is precise. Trial Number Mass, g A student measures the mass of a penny 4 times and records the following data. What can be said about the data if the actual mass of the penny is g? a. The data is both accurate and precise. b. The data is neither accurate nor precise. c. The data is accurate but not precise. d. The data is not accurate but it is precise. Trial Number Mass, g

6 Measurement and Significant Figures To indicate the precision of a measured number (or result of calculations on measured numbers), we often use the concept of significant figures. Significant figures (or sig. fig.) are those digits in a measured number (or result of the calculation with a measured number) that include all certain digits plus a final one having some uncertainty. Measurement and Significant Figures To count the number of significant figures in a measurement, observe the following rules: Count the first non-zero digit as the first significant figure. Zeros at the end of the number are only significant if there is a decimal point shown (four sig. figs) (only three sig. figs) (six sig. figs.) Measurement and Significant Figures In math calculations When multiplying and dividing measured quantities, give as many significant figures as the least found in the measurements used. When adding or subtracting measured quantities, give the same number of decimal places as the least found in the measurements used g /102.4 ml = g/ml only three significant figures (3 sig. fig.) 1.03g g g = 11.2g only one decimal place 6

7 Calculate the following with the correct number of significant figures. Calculate the following with the correct number of significant figures. a. 2 b. 1.4 c. 2.2 d e ( ) + (2.83 x 0.360) = a. 2 b. 1.4 c. 2.2 d e ( ) + (2.83 x 0.360) = Measurement and Significant Figures An exact number is a number that has an infinite number of significant digits. When you say there are twelve inches in a foot, you mean exactly twelve. or km = 1 mile (exactly) Note that exact numbers have no effect on significant figures in a calculation. Q# 1.2 (a) How many significant digits in each number? Where appropriate, use scientific notation to better express the number (b) Convert 295 miles to km given km = 1 mile NOTE: Marks will be taken off each time you have the wrong significant Figures 7

8 SI Units and SI Prefixes Table 1.1 SI Base Units In 1960, the General Conference of Weights and Measures adopted the International System of Units (or SI), which is a particular choice of metric units. This system has seven SI base units, the SI units from which all others can be derived. Quantity Unit Symbol Length meter m Mass kilogram kg Time second s Temperature kelvin K Amount of substance mole mol The last 2 are not used in chem 1010 Electric current ampere A Luminous intensity candela cd SI Units and SI Prefixes The advantage of the metric system is that it is a decimal system. A larger or smaller unit is indicated by a SI prefix that is, a prefix used in the International System to indicate a power of 10. The next slide lists the SI prefixes most commonly used. SI Prefixes Multiple Prefix Symbol 10 6 mega M 10 3 kilo k 10-1 deci d 10-2 centi c 10-3 milli m 10-6 micro μ 10-9 nano n pico p Show most common measurements in units of mass and length 8

9 The distance between atoms is sometimes given in picometers, where 1 pm is equivalent to 1 x m. If the distance between the carbon atoms in diamond i cm, what is this distance in picometers? 1) ) ) 154 4) 1,540 5) 154,000 3) 154 ( cm)(10 2 m/1 cm)(1 pm/ m) = 154 pm Temperature The Celsius scale (formerly the Centigrade scale) is the temperature scale in general scientific use. However, the SI base unit of temperature is the Kelvin (K), a unit based on the absolute temperature scale. The conversion from Celsius to Kelvin is simple since the two scales are simply offset by º K = o C K Derived Units The SI unit for speed is meters per second, or m/s. This is an example of an SI derived unit, created by combining SI base units. Volume is defined as length cubed and has an SI unit of cubic meters (m 3 ). Traditionally, chemists have used the liter (L), which is a unit of volume equal to one cubic decimeter. 3 1 L = 1 dm and 1 ml = 1 cm 3 9

10 Derived Units The density of an object is its mass per unit volume: d = m / V where d is the density, m is the mass, and V is the volume. Generally the unit of mass is the gram. The unit of volume is the ml for liquids; cm 3 for solids; and L for gases. A Density Example A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm 3. What is the density of galena? Density = mass volume = 12.4 g = = 7.56 g/cm cm3 3 sig. fig. A piece of germanium (mass = g) is placed in ml of chloroform in a 25 ml graduated cylinder. The chloroform level increases to ml. The best value for the density of germanium from these data is 1) 3.6 g/ml. 2) g/ml. 3) 5.35 g/ml. 4) 5.4 g/ml. 5) 8.0 g/ml. 3) 5.35 g/ml. Using the principle of volume displacement, the volume of the germanium is ml ml = 3.41 ml.. Density = g/ 3.41 ml = 5.35 g/ml 10

11 Units: Dimensional Analysis Q# 1. A substance has a density of g/ml. what volume will 9.37 g of this substance occupy? Ans= 6.32 ml Q#2. vinegar question will be given out Do third question (# 8) if time permits In performing numerical calculations, it is very good practice to associate units with each quantity. The advantage of this approach is that the units for the answer will come out of the calculation. If you make an error in arranging factors in the calculation, it will be apparent because the final units will be nonsense. Units: Dimensional Analysis Dimensional analysis (or the factor-label method) is the method of calculation in which one carries along the units for quantities. Suppose you simply wish to convert 20 yards to feet (1 yard = 3 feet) 3 feet 20 yards = 60 feet 1 yard Note that the units have cancelled properly to give the final unit of feet. Units: Dimensional Analysis The ratio (3 feet / 1 yard) is called a conversion factor. The conversion-factor method may be used to convert any unit to another, provided a conversion equation exists. In Chapter 3 we will be extensively using the concept of dimensional analysis and the Conversation factor NOTE: marks will be taken off if units are absent or wrong 11

12 Unit Conversion example Sodium hydrogen carbonate (baking soda) reacts with acidic materials such as vinegar to release carbon dioxide gas. Given an experiment calling for kg of sodium hydrogen carbonate, express this mass in milligrams kg x 103 g 1 kg x 10 3 mg 1 g = 3.48 x 10 5 mg Unit Conversion Suppose you wish to convert lb to grams. note that 1 lb = g so the conversion factor from pounds to grams is g / 1 lb. Therefore, lb g 1 lb = 248 g Copyright Houghton Mifflin Company. All rights reserved. Copyright Houghton Mifflin Company. All rights reserved. Operational Skills Q #1. Convert 95 km/hr to m/s Q#2. Convert 2.19 m to nm Identifying conversion factors as you go Using the law of conservation of mass. Using significant figures in calculations. Converting from one temperature scale to another. Calculating the density of a substance. Using the density to relate mass and volume. Converting units. Copyright Houghton Mifflin Company. All rights reserved. 12

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