Introduction to Chemistry

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Term List

Objectives

• Explain why a knowledge of chemistry is central to many human endeavors.
• List and describe the steps of the scientific method.
• Explain the basic safety rules that must be followed when working in the chemistry laboratory.
• Identify the metric units of measurement used in chemistry.
• List the SI base units.
• Explain what causes uncertainty in measurements.
• Compare accuracy and precision.
• Explain how to use significant digits and scientific notation.
• Calculate percent error.
• Define density and explain how it is calculated.
• Use chemistry equations to solve problems.
• Explain how dimensional analysis and conversion factors are used to solve problems in chemistry.
• Solve problems using dimensional analysis.

What is chemistry?

Chemistry is the study of all substances and the changes that they undergo.  Like all sciences, chemistry involves making observations about the world and then investigating these observations.  Chemistry is part of all that we do, from the foods that we eat, to the cars that we drive.  Because chemistry touches so many areas of our lives it is often called the Central Science.  During this year we will use physics, biology, chemistry, history, and english to learn about the processes that touch our lives everyday.

Lesson 1:  The Scientific Method

Objectives

• Explain why a knowledge of chemistry is central to many human endeavors.
• List and describe the steps of the scientific method.

The Scientific Method

The scientific method is a way of answering questions about the world we live in.  The scientific method is a process that all scientists use to investigate the observations that they make.  The steps of the scientific method are:

1. Observation
2. Question
3. Hypothesis
4. Experiment
5. Conclusion
6. Natural Law
7. Theory

A scientist makes an observation:  "Leaves on the trees are green."  The scientist takes that observation and forms a question:  "Why are the leaves on the tree green?"  A hypothesis is a possible explanation for the answer to the question.  The hypothesis must be able to be tested through experiments.  If you are unable to create an experiment then you need to develop a different hypothesis.  A possible hypothesis to our question is: "The green color of the leaves is caused by the brown soil surrounding the trees."  Now that we have an hypothesis, we must develop an experiment.  In the case of our trees, we could grow them in black soil, brown soil, red clay, white sand, and brown sand.  The independent variable, variable that you change, is the color or type of soil.  The dependent variable, variable that changes from the experiment, is the green color of the leaves.  You would have to grow at least three trees in each type of soil in order to verify our results using statistics.  Your experiment would either prove or disprove your hypothesis.  A conclusion is the result of our experiments.  We know that our experiment would disprove our hypothesis.  The color of the soil has nothing to do with the green color of leaves, rather, the color results from a chemical called chlorophyll.  Since our experiment disproved our hypothesis, we would have to form a different hypothesis and a new set of experiments.  Scientists often spend their entire life searching for the answer to their question. 

If a set of experiments prove your hypothesis correct, then you must develop a new set of experiments that could be used to prove your hypothesis.  If all of your experiments prove your hypothesis correct, then you may be able to develop a natural law, which describes how nature behaves but does not describe why nature behaves in that particular way.

After many experiments and many different approaches to the question, the scientist may be able to develop a theory.  The theory explains why nature behaves in the way described by the natural law.  It answers not only the original question, but also any other questions that were raised during the process.  The theory also predicts the results of further experiments, which is how it is checked.  Theories are not the end of the process.  Theories can continued to be proved or they can possibly be disproved.  Theories are not facts, but our best guess as to how the world works, based on the scientific evidence we have collected.  Many theories that we will discuss during the year could be disproved as our understanding of the universe continues to expand.

Lesson 2:  Safety in the Laboratory

Objectives

• Explain the basic safety rules that must be followed when working in the chemistry laboratory.

Laboratory Safety

Chemistry is a science of investigation.  We will be conducting investigations using chemicals and laboratory equipment.  Before you can do experiments in the laboratory you must understand the safety rules and procedures.

1. Follow all rules of the Teacher.
2. Notify your teacher of problems.
3. Know how to use the safety equipment in the laboratory.
4. Wear goggles at all times
5. If you have long hair, tie it back
6. Avoid awkward transfers.
7. If it's hot, let it cool.
8. Carry chemicals defensively.
9. Dispose of chemical wastes properly.
10. Clean up afterward.

You will be given one warning, after that you will be asked to leave the laboratory and will receive a zero on the lab.  Safety in the laboratory is very important, because we are not only concerned with your safety but the safety of those around you.

Lesson 3:  Units of Measurement

Objectives

• Identify the metric units of measurement used in chemistry.
• List the SI base units.

Measurements are an integral part of science.  In chemistry we use a system that is based on an international system of measurement called the metric system. 

International System of Units

In 1960, the units of the metric system were streamlined by an international conference.  The conference created a system of units called the International System of Units, abbreviated SI.  The SI is built upon a set of seven metric units, which are called the base units of the SI. 

SI Base Units
Physical Quantity	Unit Name and Symbol
mass	kilogram, kg
length	meter, m
time	second, s
count, quantity	mole, mol
temperature	kelvin, K
electric current	ampere, A
luminous intensity	candela, cd

 

Derived Units

Scientists are able to use the seven base units to make or derive all other units that are needed.

Derived Units Commonly Used in Chemistry
Physical Quantity	Unit Name and Symbol	Units Derived From
area	square meter, m2	meter
volume	cubic meter, m3	meter
force	newton, N	meter, kg, second
pressure	pascal, Pa	Newton, square meter
energy	joule, J	Newton, meter
power	watt, W	Joule, second
voltage	volt, V	Watt, ampere
frequency	hertz, Hz	seconds
electric charge	coulomb, C	Ampere, second

Volume is important in chemistry, but is not an SI unit.  It is derived or made from the SI unit of meter.  A cube measuring 1 decimeter is called a liter.  

Mass and weight are often used interchangeably, which is not correct.  Mass is a measure of the amount of force required to move an object. All objects have mass, even in space.  Weight is a measure of the force exerted on a mass by gravity. Weight changes depending on the amount of gravity. Our weight on earth is different than on the moon. In space we do not have any weight.

Metric Units

Lesson 4:  Uncertainty in Measurement

Objectives

• Explain what causes uncertainty in measurements.
• Compare accuracy and precision.

Making a measurement always involves estimating part of the number. As we will see, part of a measurement is certain and part of a measurement is uncertain.

Making Measurements

Instruments such as electronic balances have digital displays.  The final digit on a digital display is the estimated digit.  With a digital display, the estimation of the uncertain digit is done for you electronically.  When you record a measurement from a digital display, you will always write all of the digits that are shown.

Instruments such as a graduated cylinder or a thermometer have scales marked or etched on the glass.  These are called analog scales and require that you estimate the final digit in a measurement.  A graduated cylinder that has marks in 1-mL can be read to the nearest 0.1 mL.  The estimated digit is the tenth's place.  You will always estimate 1/10th of the smallest increment on the instrument.

Examples:

1. Graduated cylinder:  certain digit, 1-mL, estimated 0.1 m; 12.1 mL, 10.5 mL, 100.1 mL
2. Beaker: certain digit 10 mL, estimated 1 mL; readings: 21 mL, 52 mL, 105 mL
3. Flask: certain digit 100 mL, estimated 10 mL; reading: 110 mL, 250 mL, 310 mL

Reliability in Measurement

Precision:  One way to check your work is to repeat the procedure several times.  Precision refers to the degree of agreement among several measurements of the same quantity.  If you get the same answer each time you perform the procedure, then your work is precise.  You do not have to have the correct answer, just the same answer.

Accuracy:  The second way to check your work is to compare your answer to an accepted standard.  Accuracy refers to the agreement of a particular value with the accepted value.  The closer your answer is to the accepted value, the better your work.

You want to be both precise and accurate, but you could be precise and not accurate, or accurate and not precise.

Lesson 5:  Working With Numbers

Objectives

• Explain how to use significant digits and scientific notation.
• Calculate percent error.
• Define density and explain how it is calculated.

 A measurement is rarely meaningful by itself.  More commonly, measurements are combined by adding, subtracting, multiplying, or dividing them to produce the values of mass, volume, or temperature needed in a scientific investigation.  In this section, you are going to use several mathematical tools to help you with this process.

Significant Digits

What is a significant digit?  Significant digits tell us the precision of your measuring device.  If you use two different graduated cylinders, one marked in 1-mL and one marked in 10-mL.  The one marked in 1-mL increments is more precise because you will be able to make the same measurement over and over again with little difficulty.  The one marked in 10-mL increments is less precise because it will be more difficult to record the same measurement over and over again. The one marked in 1-mL will have more significant digits than the one marked in 10-mL increments.

When measurements are combined mathematically, the uncertainty of the separate measurements must be correctly reflected in the final result.  A set of rules exists for this task, which depends on keeping track of the significant digits, or significant figures, in each separate measurement.

What digits are significant in a measurement?  All of the digits in a measurement that you are certain of, plus the one that you estimate are significant.  If you are using an instrument with a digital display, all of the digits that are shown on the display are significant.  If you are using an instrument with an analog scale, then you will have to properly estimate the final digit for that measurement.  If you have written the answer properly, then all of the digits will be significant. 

Sometimes a digit may not be significant.  We have a set of rules for you to follow when determining the number of significant figures in a measurement.

1. All non-zero digits are significant.
2. Zeros that are between non-zero digits or are between significant digits are themselves significant.
3. Zeros at the beginning of a number are not significant.
4. Zeros at the end of a number are significant if there is a decimal in the number.
5. Zeros at the end of a number are not significant if there is not a decimal in the number.
6. If the number is written in proper scientific notation, then all digits listed in the mantissa are significant.  The exponent is not used when counting significant figures.
7. If a number is the result of counting then we do not use that number to determine significant figures.  The number has infinite significant figures.
8. Exact number, numbers that arise from a definition, have infinite significant figures and are not counted.

Examples:

1. 256 mL                   3 significant figures (See Rule 1)
2. 206 V                      3 significant figures (See Rule 2)
3. 0.01 m                    1 significant figure (See Rule 3)
4. 10.00 g                   4 significant figures (See Rule 4 and Rule 2)
5. 120 kg                    2 significant figures (See Rule 5)
6. 1.230 x 1023 ug     4 significant figures (See Rule 6)
7. 1230 basketballs    infinite significant figures (See Rule 7)
8. 1000 g = 1 kg         infinite significant figures (See Rule 8)

Significant Digits in Calculations

When measurements are used in a calculation, our answer cannot have more significant figures that the single measurement with the least significant figures.  In other words your answer cannot be more precise than the least precise measurement that was made.

Multiplication and division:  Find the number with the least number of significant figures and round your final answer to that number of significant figures.

 

Volume	=	length	x	width	x	height
	=	3.05 m	x	2.10 m	x	0.75 m
	=	4.80375 m3
	=	4.8 m3

In the above example, the final answer has more significant figures (six) than any number that was used to calculate the answer.  The height has two significant figures, therefore the answer is rounded to two significant figures.

Addition and Subtraction:  Find the number with the least number of decimal places and round your final answer to that number of decimal places.  Do not count significant figures.

 

Shoes	951.0     g
clothing	1407        g
ring	23.911 g
eyeglasses	158.18   g
total	2540.       g

In the above example, the mass of the items was added together.  The final answer was rounded to zero decimal places because the mass of the clothing did not have any decimal places listed.

If a process has both multiplication and addition, then use the rule for multiplication for rounding your final answer.

q_sur = m x C x (Tf-Ti)

q_sur = 75.0 g  x 4.184 J/g C x (31 C - 21 C)

q_sur = 3138 J

q_sur = 3100 J

In the above example the heat of the surroundings was calculated to be 3138 J.  The temperatures used in the calculation have 2 significant figures.  The final answer should then be rounded to two significant figures.

Scientific Notation

1,300,000 can be expressed 1.3 x 10⁶

The number means to multiply 1.3 x 10 x 10 x 10 x 10 x 10 x 10.

Moving the decimal to the left gives a positive exponent, to the right gives a negative exponent.

Numbers greater than 1000 and less than 0.001, should always be written in scientific notation.

Numbers in scientific notation should have 1 nonzero number before the decimal.

Percents and Percent Error

We will use several calculations to evaluate the accuracy of your measurement.

Percent True

Percent true is used to determine how close you are to the accepted value.  This calculation is the same as the one that you would use for determining your grades.  All values must be greater than zero and the closer to 100% the measurement the more accurate.

 

Percent True =	measured value	x 100
	accepted value

Percent error

Percent error is also used to determine the accuracy of your measurement, but the closer the value is to zero the more accurate the measurement.

 

Percent Error =	measured value - accepted value	x 100
	accepted value

Density

Density is one of the important properties of matter.  Density is a measurement of the amount of matter in a certain amount of space. 

The formula for density is given:

Density =	mass
	volume

The units for density are:  solids (g/cm³), liquids (g/cm³ or g/mL), gases (g/L)

The density of water is given as 1.00 g/mL.  This is the same as saying 1.00 g per 1.00 mL.

 

1 g	=	1.00 g
mL		1.00 mL

You need to keep this in mind when you are working with density in problems.  The units in chemistry are treated the same as if they were variables in algebra.

Lesson 6:  Solving Equations

Objectives

• Use chemistry equations to solve a problem.

Solving equations in chemistry is the same as solving equations in algebra.  It appears to be more difficult in chemistry because we do not use "x", "y" and "z" as part of every equation that we solve.  Sometimes we will solve equations with a symbol (g, c, t...), sometimes we will solve equations that contain words (helium, density, pressure) and sometimes we will have equations that consist of a combination of symbols, numbers and words.  Do not make this any harder than it is.  Follow your rules for solving equations that you learned in algebra.  

Example:  

The equation that relates the wavelength of electromagnetic radiation to its frequency is given below:

 

where

solving for frequency results in 

solving for c results in

Solving equations is a simple rearrangement of the variables.  It doesn't matter the form of the variables, just that you follow the rules the same everytime.

Equations used in Chemistry

The following are some of the equations that we will be using throughout the year.  It is not important at this time to fully understand each equation, but to be able to recognize which equation we are using based on the information that is given in the problem.

 

Molarity (M) =	moles of solute
	liters of solution

"moles of solute" is the number of moles of the compound that is dissolved in the solvent.
"liters of solution" is the total number of liters that the solution occupies.

Molality (m) =	moles of solute
	kilograms of solvent

"moles of solute" is the number of moles of the compound that is dissolved in the solvent.
"kilograms of solvent" is the total number of kilograms of solvent that is in the solution.

Mole Fraction =	moles of substance
	total number of moles in solution

 

Normality (N) =	equivalents
	liters of solution

 

Density	Mass
	Volume

Lesson 7:  Solving Problems

Objectives

• Explain how dimensional analysis and conversion factors are used to solve problems in chemistry.
• Solve problems using dimensional analysis.

The majority of the work in this course involves solving problems.  The problems will revolve around laboratory situations, and will require you to choose the correct formulas, pick the information that is important to the problem, set up the problem correctly and to report the final answer to the correct number of significant figures.

Number Line

The number lines below can be used to convert from one unit of measurement to another.  The first number line is used to convert "normal" units of measurement.  This is to say that the units are not squared or cubed.  If you move from the left to the right you will divide, and moving right to left you will multiply.

Normal Measurements are based on the number 10.  Each interval on a metric number line represents 10 units.

Example 1:  Convert 45 nanometers to meters.  

Step 1:  Counting the number of spaces from nanometers to meters, we find that there are 9 intervals.
Step 2:  Since we are moving from the left to the right, we will divide.  There were 9 intervals, so we will be dividing by 1 x 10⁹.
Step 3:  Dividing 45 by 1 x 10⁹ gives us:  4.5 x 10^-8 meters.

TI-83 Keys:

4

5

x

1

2nd

EE

9

 TI-83 Screen Results:  45*1E9

Example 2:  Convert 3.45 Megagrams to micrograms.

Step 1:  Counting the number of spaces from Megagrams to micrograms, we find that there are 12 intervals.
Step 2:  Since we are moving from the right to the left, we will multiply.  There are 12 intervals, so we will be multiplying by 1 x 10¹².
Step 3:  Multiplying 3.45 by 1 x 10¹² gives us 3.45 x 10¹² micrometers.

 

Example 3:  Convert 24.5 mm² to hm².

Square units are based on the number 100 (10²).  Each interval consists of 100 units.

Step 1:  Counting the number of spaces from mm² to hm², we find 5 intervals.
Step 2:  Since we are moving from the left to the right, we will divide.  There are 12 intervals, so we will be dividing by 1 x 100⁵.
Step 3:  Dividing 24.5 by 1 x 100⁵ gives us 2.45 x 10^-9 hm².  Note that the answer is in proper scientific notation. 

TI-83 Keys

2

4

.

5

÷

(

1

x

1

0

0

^

5

)

Note the use of parenthesis when using the "^" symbol.

TI-83 Screen Results:  24.5/(1*100^5).  

 

Example 4:  Convert 678 cm³ to m³.

Cube units are based on the number 1000 (10³).  Each interval consists of 1000 units.

Step 1:  Counting the number of spaces from cm³ to m³, we find 3 intervals.
Step 2:  Since we are moving from left to right, we will divide.  There are 3 intervals, so we will be dividing by 1 x 1000³.
Step 3:  Dividing 678 by (1 x 1000³) gives us 6.78 x 10^-7 m³. Note that the answer is in proper scientific notation.

Dimensional Analysis

Dimensional analysis is a process for solving problems and for converting from one unit to another.  Once you learn how to use dimensional analysis, you will be able to solve the majority of the problems you will encounter in this course.  In order to solve a problem using dimensional analysis you must have or must set up unit equalities.

Unit Equalities

A unit equality is an equation that shows how different units are related.

 

Unit Equalities
Metric to Metric
1000 Mm = 1 Gm	10 dm = 1 m	1 L = 1000 cm3
1000 km = 1 Mm	100 cm = 1 m	1 ml = 1 cm3
1000 m = 1 km	10 cm = 1 dm
1000 mm = 1 m	10 mm = 1 cm
1000 um = 1 mm	100 mm = 1 dm
1000 nm = 1 um
1000 pm = 1 nm
English to Metric		Miscellaneous
1 in. = 2.54 cm		1 ft = 12 in.
1 gal = 3.785 L		1 mi = 5280 ft
1 cal = 4.184 J		1 min = 60 s
1 atm = 101,325 Pa		1 hr = 60 min
		1 atm = 760 mm Hg
*These unit equalities follow from the definitions of the metric prefixes and they apply to any metric unit which would use these prefixes.

Conversion Factors

The next step in solving a problem is to write conversion factors from the unit equality. 

The unit equality 1 in. = 2.54 cm  can be written

1 in	or	2.54 cm
2.54 cm		1 in

In a conversion factor the unit on the bottom is the unit you are starting with, and the unit on top is the unit you are ending with.  The first conversion factor will convert centimeters to inches and the second will convert inches to centimeters.

 

Solving Problems

When solving a problem make a list of what you know and what you need to know.

It is best to start with what you are given in the problem.

Dimensional analysis is used to cancel out units not numbers. You must always write the number and the unit when working the problem.

When solving a problem using dimensional analysis, you need to start with a grid:

 

The horizontal lines are used to separate numbers from the conversion factor and represent the mathematical operation division.  The vertical lines are used to separate different conversion factors and represent the mathematical operation multiplication.

Example 1:

1. A tile was measured to be 45.56 cm. What is the measurement in inches?
2. What do we know: tile is 45.56 cm, 1 inch = 2.54 cm.
3. Draw a grid. In the first box write 45.56 cm. The second grid to the right is used to convert from cm to inches. Since cm is on the top in the first grid, we must put cm on the bottom in the second grid and inches on top. We can cancel cm and we are left with inches.
4. The answer is rounded to 4 significant figures because the number from the problem has four, while the conversion factor is not counted because it is a definition.

45.56 cm	1 in.	= 17.94 in.
	2.54 cm

Example 2:

1. The mass of a lead block was measured to be 45.6 kg. What is the mass in micrograms?
2. What do we know: The mass is 45.6 kg, 1 kg = 1000 g, 1 g = 1000 mg, 1 mg = 1000 ug.
3. Start with what is given, then arrange the units so they cancel.

45.6 kg	1000 g	1000 mg	1000 ug	= 4.56 x 1010 ug
	1 kg	1 g	1 mg

Example 3:

1. If the speed limit is 65 miles/hr, what is the speed limit in kilometers/hr.
2. What do we know: 1 mi = 1760 yd 1.094 yd = 1 m 1000 m = 1 km
3. Start with 65 miles = 1 hour and then arrange conversion factors to change to kilometers/hour.

65 miles	1760 yd	1 m	1 km	= 88 km/hr
1 hour	1 mile	1.094 yd	1000 m

*** Do not let mixed units such as miles/hour or g/mL or g/cm3 cause you problems.  Using this method you can work with each unit separately, but use a clear organized approach.

Example 4:

1. A car is advertised to get 15 km/L, what is it in miles/gallon.
2. What do we know: 1 km = 1000 m 1 m = 1.094 yd 1760 yd = 1 mi 1 L = 1.06 qt 4 qt = 1 gal
3. Start with what is given in the problem (15 km/L) and then arrange conversion factors so that the units will cancel.  Convert the km to miles and then convert liters to gallons.

15 km	1000 m	1.094 yd	1 mi	1 L	4 qt	35 miles/gallon
1 L	1 km	1 m	1760 yd	1.06 gt	1 gal

 

A Four-Step Problem-Solving Strategy

1. Analyze:  Read the entire problem carefully.  Identify the unknown quantity in the problem.  Organize your given information and try to determine what is important and what is simply extra information.
2. Plan:  What is your approach to solving the problem going to be?  Is there an equation that you will need to use or can you simply use dimensional analysis?  How is this problem similar to other problems that you have solved?
3. Solve:  Perform the mathematical steps outlined in your plan.  Check units and significant figures.
4. Evaluate:  Ask yourself, Does the answer make sense?  Can you follow your steps.  Do the units on top match units on the bottom?  If you are going from a small unit to a large unit your answer should be smaller.  If you are going from a large unit to a small unit, then your answer should be larger.  If it doesn't make sense to you then it probably won't make sense to me either.

Graphing

Chemists use graphs to transform rows and columns of data into a picture.  With a graph we can determine the general trend, our accuracy and/or precision and any patterns without having to examine the individual numbers.  There is an old saying "a picture is worth a thousand words", well a graph is also worth a thousand words.

Graphs in chemistry do not have to involve all four quadrants.  Most graphs that we will use are called scatter plots or xy graphs.  XY Plots are most common in science when we want to see if there is a direct relationship between the independent variable (x) and the dependent variable (y).

Given the data below, the table is difficult to view, but the graph is very clear.  We can see from the graph that there is a direct relationship between temperature and volume and we can also see where the data points do not agree with the line of best fit.

Effect of Changes in Temperature on the Volume of Air in a Balloon

All of the topics covered in this Chapter are basic skills that you will need to be able to use ALL YEAR.  Make sure that you learn them well now because you will see them again in future chapters and labs.