CHM1045L Complete Lab Manual PDF

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lab manual for chem 1 lab, detailing all procedures....

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THE SCIENTIFIC METHOD

DATA AND OBSERVATIONS

INTRODUCTION:

1. What are you investigating today?

One of the things that makes a science a science is the use of a logical, orderly approach in dealing with information. This approach is called the SCIENTIFIC METHOD. The first step in the scientific method involves making OBSERVATIONS of a behavior in a particular system. Observations can be made with any of the senses (sight, hearing, smell, touch and taste). Mechanical devices such as thermometers or balances may also be used in making observations.

2. Describe the object.

The next step is to examine the observations, looking for any pattern in the behavior observed. If such a pattern is found, it is stated formally as a LAW. A law may be written in words, or it may be expressed as a mathematical equation.

3. Can you collect data using the object? If so, record some data below.

The third step goes beyond simply collecting and organizing information, and involves developing a model that will explain why the observed behavior occurs. This model is called a HYPOTHESIS. A good hypothesis should not only explain the behavior already observed but also predict some behavior that has not yet been seen. This provides a way to test the hypothesis, which is the fourth step in the scientific method. If the hypothesis fails a test it must be modified or discarded. However, passing a test (or even many tests) does not prove a hypothesis - the hypothesis may still not be the correct description of the system in which the behavior has been observed. A hypothesis that has passed many tests is called a THEORY. To review, the four steps in the scientific method are:

4. Analyze the data and determine the average.

5. Write a hypothesis about this object?

6. Test your hypothesis. Do your observations support your hypothesis?

1. Make observations of a behavior. 2. Look for a pattern or trend in these observations; if one exists; express it in the form of a LAW. 3. Try to develop a model that accounts for the behavior. This model is a HYPOTHESIS. 4. Test the hypothesis. In this lab exercise you are going to practice using the scientific method. You will make observations, identify any patterns, and write these as laws. You will then attempt to develop a hypothesis to explain the behavior. Finally, you will suggest ways to test your hypothesis.

7. Can you develop a theory/law about this object? State it below.

MATH AND CALCULATOR EXERCISE Students in chemistry courses must be able to perform mathematical calculations and use a scientific calculator correctly. In this exercise, you will be introduced to the rules for expressing answers using significant figures and will be shown the proper way to use a scientific calculator. Scientific notation (expressing numbers as powers of 10’s) will also be presented. These topics will be crucial to your success in lab as well in lecture. Scientific Notation: The use of scientific notation is a method of conveniently expressing very large or small numbers as powers of tens. A general example is N x 10n where N is a number between 1 and 10 and n is the exponent, either positive or negative. Here are some typical examples: The number: 11 = 1.1 x 101 12 = 1.2 x 103 150 = 1.5 x 102 16000 = 1.6 x 104 0.1 = 1 x 10-1 0.001 = 1 x 10-3 You notice in the answers that there is only one digit other than zero to the left of the decimal. A number less than 1 will have a negative exponent and a number greater than 1 will have a positive exponent. The procedure of entering these numbers into a calculator is dependent on the brand or model of the calculator you use. Later, you will be given examples using various calculators. SIGNIFICANT FIGURES: The concept of significant figures (SF, sig figs) is based upon the degree of uncertainty that exists in any measurement and the error introduced into any calculation using that measurement. As a chain is no stronger than its weakest link, a calculation can be no more accurate that the least accurate number used in the calculation. The application of this concept can produce results that can be very confusing to students familiar with the precise requirements of classical mathematics. For example, you learned in grade school that 100 plus 1 will equal 101. Assuming that these to numbers represent the results of

measurements and reflect their accuracy, with the application of significant figure rules, the answer will be 100, not 101. Your instructor and text will explain why this rather strange phenomenon occurs. The purpose of this handout is to summarize the rules for using significant figures that you will have to apply in all calculations in chemistry class and lab. Before you can apply significant figure rules, you must be able to determine how many sig figs there are in any given number. The rules are: 1. All non-zero digits are significant in any number. 2. Zeroes between non-zero digits ARE always significant. 3. Trailing zeroes in non-decimal numbers ARE NOT significant. 4. Trailing zeroes in decimal numbers ARE significant. 5. Leading zeroes in decimal numbers less than one ARE NOT significant. 6. In decimal numbers greater than one, ALL digits are significant. 7. Exact numbers* have an infinite number of sig figs. 8. All digits in a number written in scientific notation are significant. 9. A line over a zero** indicates that that zero and all digits to the left are significant. Examples: Number 36,892 2,004 340,000 3.45000 0.000456 80.41 7 days/week 3.4050 X 1010

Sig Figs 5 4 2 6 3 4 ∞ 5

Rule(s) 1 2 3 4, 6 5 6 7 8

*Exact numbers are used in the calculation but their effect on the number of sig figs in the answer is ignored. **If you need to write 3000 to 3 sig figs, put a line over the middle zero (3000). The rules for determining the number of sig figs in the answer to a calculation are different depending upon which mathematical operation you are performing.

For multiplication, division, and exponentiation the answer can have no more sig figs than the LEAST number of sig figs in any number used in the calculation. For example:

To enter a number like 2.45 * 107, you should enter the following: 2.45 ee (or exp) 7

459 * 32 * 0.003512 = 52

51.584256 to 2 sigs

Because 32 only has 2 SF the answer must be rounded off to 2 SF. 3000∗ 3.2222∗ 652 3.589002

3.45 ee (-) 6 or 3.45 exp (-) 6 or 3.45 ee (+/-) 6 = 2,000,000

1756093.532 to 1 sig

Because 3000 has only 1 SF, the answer must be rounded off to 2 million. A note on rounding, digits to the right of the decimal are dropped when rounding, digits to the left of the decimal are replaced with zeroes, NEVER DROPPED! 34.5677 equals 34.6 when rounded to 3 sig figs. 45089.000 equals 45,000 when rounded to 2 sig figs For addition and subtraction, the rule is a little confusing. The answer can have no more significant PLACES than the least significant place in any number used in the calculation. For Example (Least significant place underlined): 12,000 345 � + 0.0034 12,000

To enter a number like 3.45 * 10-6, you should enter the following:

Good to the 1,000’s place Good to the 1’s place Good to the 10,000th place Answer good to the 1,000’s place

Notice the 345 and 0.0034 were essentially discarded. This is because the first number was inaccurately measured and was only good to the thousand’s place. Calculator Use: There is a button on every scientific calculator that should be used for entering numbers in scientific notation. That button will be either “EXP” or “EE”. On a few models, the button is “X10”. Do not use the “X” button followed by the “10”. This is totally wrong. When you enter using the EXP or EE button, this means “x 10”. To enter the number 2.4 x 104 you should enter 2.4, then hit EXP or EE, then 4. To enter 4.78 x 10-5, you should enter 4.78, then EXP or EE, then (-) or (+/-), then 5.

PRACTICE PROBLEMS Significant Figures 0.00000342 0.000300023 123.34507 1200054.33 100000333

____ ____ ____ ____ ____

2.4304 x 1011 2100000 123,456,000.68

____ ____ ____

Calculations with Significant Figures 123.4567 + 21.34 + 0.1

__________

210.1234 – 24.4

__________

21.45 * 3.6

__________

113.45 * 21.56 * 3.4 (1.256 * 107) (3.54 * 104)

__________ __________

(12.4734) / (2.11)

__________

(12.56 + 2.1) / (3.345) (5.1 * 10-7) (3.55 * 10-8)

__________ __________

√7.43 (4 * 972) + (76.4 * 29.3) – (12 * 7)

__________ __________

4.1 ∗ 10−3 – 6.9 ∗ 10−2 7.2 ∗ 10 −6 + 8.943 ∗ 104 �0 400� 0 x 200

__________

6.3 x 107

__________

CHM1045L – Experimental and Laboratory Methods Equipment and Glassware Beaker

Crucible tongs

Erlenmeyer flask

Florence Flask

Graduated Cylinder

Test tube holder

Suction Flask (sidearm)

Glass Funnel

Büchner Funnel

Test tube rack

Watch Glass

Buret

Instruments Analytical Balances, accurate to 0.0001 g

Top Loader Balance, accurate to 0.01 g

CHM1045L – Data Sheet and Lab Reports 1. Data sheets must be completed and the measured values recorded accurately with their respective units. You must show all of your work in the appropriate locations in the data sheet. a. Remember, all measurements (except electronic readings) must be estimated to an additional decimal place, especially thermometers and burets 2. Laboratory reports must be written in the past tense and follow the template below. Numerical values must be written with the correct number of significant digits and the appropriate units. Title of Experiment Your Name, Lab Partners Date Introduction - A couple of sentences describing the reason the experiment was performed Observations, Results, and Conclusions - A couple of paragraphs reciting the data that you collected (not how you collected it) and the results/conclusions from your calculations. For Example: Density of Solids and Liquids Domenick Grasso 8/30/2013 Introduction In this experiment we determined the density of two unknown solids and two unknown liquids. Density is an intrinsic physical property of a substance and is the relationship between its mass divided by its volume. Observations, Results, and Conclusions Metal 1 had a silver color, felt dense, and had a mass and volume of 12.3456 g and 1.26 mL, respectively. We calculated its density to be 9.79 g/mL. Metal 2 had a copper color and was bendable. It has a mass and volume of 7.8012 g and a volume of 1.01 mL, respectively. Its density was calculated to be 7.72 g/mL. Liquid 1 was clear and odorless and had a mass and volume of 9.9852 g and 10.00 mL, respectively. Its density was 0.9985 g/mL. Liquid 2 was green and smelled of alcohol. It had a mass and volume of 10.0021 g and 10.00 mL. Its density was calculated to be 1.000 g/mL.

DENSITIES OF SOLIDS AND LIQUIDS INTRODUCTION Density is defined as mass per unit volume. In this experiment, the densities of various unknown solid and liquid materials will be determined. For a solid having a smooth surface and a regular shape, the volume of the solid can be measured directly. However, for a solid that is irregular in shape, the best method for determining its volume is by the displacement of some liquid. If the volume of a liquid in a burette is measured and then the solid is immersed in it, the difference in the two volume readings is the volume of the solid. Mass is found beforehand by weighing the dry solid on the analytical balance. For a liquid, both the mass and volume are readily determined by weighing a known volume of the liquid in a volumetric container. Knowing the mass and volume of a material allows one to calculate the density of that material by the equation: Density =

Buret reading: 26.77 mL *Burets are read to 0.01 mL.

Mass Volume

Densities for liquids and solids will vary depending upon the temperature at which the measurements are taken. For example, the density of water at 0 °C is 0.99987 g/cm3 while at 25 °C it is 0.99704 g/cm3.

Filled volumetric flask *Volumetric flasks are accurate to 0.00 mL

MATERIAL AND EQUIPMENT Read all glassware to the meniscus and record data accurately with units. Measure all weights using an analytical balance.

Thermometer reading: 25.1 °C *Thermometers are read to 0.1 °C.

DENSITY PROCEDURES, DATA, AND OBSERVATIONS

Name:

Procedure For the density determination of solids

Data and Observations Metal 1

Metal 2

1. Describe each unknown metal you received.

____________ ____________

2. Make sure the metal is dry then weigh it on an analytical balance.

____________ ____________

3. Place the buret in the buret clamp at your workstation. Fill the buret about half full with tap water and record the buret reading.

____________ ____________

4. Tilt the buret and gently slip the metal unknown into the buret, taking care not to let it drop with force. Hold the buret in one hand and lightly tap the bottom of the buret with the other hand to dislodge any air bubbles adhering to the surface of the metal. Record the level of the buret.

____________ ____________

5. Calculate the volume of the metal.

____________ ____________

6. Calculate the density of each solid unknown using the volume and mass data that you collected.

____________ ____________

For the density determination of liquids

Liquid 1

Liquid 2

1. Describe each unknown liquid (e.g., color, smell, et cetera) you received.

____________ ____________

2. Use clean, dry 10.00 mL volumetric flasks. Record the weight of the volumetric flasks on an analytical balance.

____________ ____________

3. Fill the volumetric flask to the mark with the liquid unknown using an eyedropper (the bottom of the meniscus touching the mark), record the combined weight of flask plus its contents.

____________ ____________

4. Calculate the weight of the liquid by difference.

____________ ____________

5. Pour the unknown liquid from the volumetric flank back into the test tube from which it came. Place the thermometer in the liquids and record the temperature when stable.

____________ ____________

6. Calculate the density of each liquid unknown.

____________ ____________

Calculations: Using the space below show all calculations used to determine the density of unknown metal 1.

NOMENCLATURE NOMENCLATURE RULES INTRODUCTION: Binary Ionic Compounds Currently there are over 10 million known chemical substances. This list of known chemical substances continues to grow. Over the past years several different approaches have been taken to naming chemical substances. In ancient times pure elemental substances were named using Greek, Latin, Arabic, or Persian word roots that described certain characteristics of the element. Elements are also assigned chemical symbols. The chemical symbols use one or two letter notations to represent an atom of a specific element. Some atomic symbols have been assigned to certain elements based on their ancient Latin or Greek names. For example, the atomic symbol for copper is Cu, from the Latin word cuprum, which describes the reddish color of this metal. When a single letter is used for the chemical symbol, like "S" for sulfur, we write the symbol in the upper case (capital). However, when the chemical symbol has two letters, as in "Si" for silicon, we capitalize the first letter and use lower case for the second letter. The atomic symbol for mercury is Hg, from the Latin word hydrargyrum, for quicksilver. Naming compounds which are composed of more than one type of atom can be somewhat more complicated. Over the past 150 years chemists have refined their methods of naming chemical compounds. The guidelines for chemical nomenclature are based on the separation of substances into different types. These rules of nomenclature are determined by an organization of chemists called "The International Union of Pure and Applied Chemistry" or IUPAC for short. The primary division is between organic compounds and inorganic compounds. Organic compounds contain carbon, normally in combination with other elements, such as, hydrogen, oxygen, nitrogen, sulfur, and phosphorus. All other compounds are called inorganic compounds. Early chemists associated organic compounds with living organisms, while inorganic compounds were associated with the non–living matter. The next major division for simple inorganic substances is between ionic compounds and molecular compounds. In this laboratory exercise we will study the most current practices in chemical nomenclature for various types of simple chemical substances.

Binary ionic compounds are compounds formed by combining a Group 1A and 2A metals with a non– metal element. Usually these are the simplest compounds because the elements of Groups 1A and 2A all have valence electrons that are “s” electrons which allow them to have only one oxidation state. In addition to these two groups there are four other elements that tend to exhibit only one oxidation state each. These are aluminum (Al3+), cadmium (Cd2+), zinc (Zn2+) and silver (Ag+). The naming of compounds containing these elements with a monatomic anion or a polyatomic anion is straight forward. The name is composed of the cation name and the anion name. Nomenclature of Cations All metal cations are named by using the elemental name. Some of the metals exhibit a single oxidation state. They are named using the English name for the element with no additional information pertaining to the ionic charge because there is only one choice. Note that the correct naming of an ion’s name will include the word “ion” as in H+ is the hydrogen ion. In naming a compound or formula the word ion is not use as in NaCl is sodium chloride. Element Lithium Potassium Barium Calcium Zinc Aluminum Silver

Ion Li+ K+ Ba+2 Ca+2 Zn+2 Al3+ Ag+

Name lithium ion potassium ion barium ion calcium ion zinc ion aluminum ion silver ion

There are three transition elements that tend to have only one oxidation state; they are Cd2+ (cadmium), Zn2+ (zinc) and Ag+ (silver). The rest of the transition metals have variable oxidation states and require naming the oxidation state in the compounds name. In addition to the transition metals many of the semimetals and a few non–metals exhibit one or more oxidation states. Aluminum in the combined state is always 3+ so no indication is given in the name of aluminum compounds.

Naming of Anions Names for the monoatomic anions (negative) ions are composed of the elemental stem name followed by the –ide suffix. Ion I– Cl– O2– As3–

Element Iodine Chlorine Oxygen Arsenic

Name Stem Iod– Chlor– Ox– Arsen–

Ion Name iodide ion chloride ion oxide ion arsenide

For monoatomic variable charge cations the nomenclature rule above is insufficient. Naming of variable ions use either of two systems, the more common IUPAC approved Stock system or the older Latin system which is retained because of historical value and to fact that many compounds bear an old Latin derived name. You will need to learn both systems of nomenclature. Stock System The Stock System addresses all cations and anions. We will deal with the cations first and the anions later. Again the metals that have multiple oxidation states are best treated first. The Stock system uses the English ...