Keys to the Study of Chemistry กุญแจในการเรียนวิชาเคมี

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Keys to the Study of Chemistryกุ�ญแจในกุารเร�ยนวิ�ชาเคมี�

Chapter 1


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1.1 Some Fundamental Definitions

1.2 Chemical Arts and the Origins of Modern Chemistry

1.3 The Scientific Approach: Developing a Model

1.4 Chemical Problem Solving

1.5 Measurement in Scientific Study

1.6 Uncertainty in Measurement: Significant Figures

Chapter 1 : Keys to the Study of Chemistry


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Is the study of matter,

its properties,

the changes that matter undergoes,

and

the energy associated with these changes.


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Definitions

Chemical Properties those which the substance shows as it interacts with, or transforms into, other substances such as flammability, corrosiveness

Matter anything that has mass and volume -the “stuff” of the universe: books, planets, trees, professors, students

Composition the types and amounts of simpler substances that make up a sample of matter

Properties the characteristics that give each substance a unique identity

Physical Properties those which the substance shows by itself without interacting with another substance such as color, melting point, boiling point, density


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Figure 1.1

A Physical change B Chemical change

The distinction between physical and chemical change.


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Some Characteristic Properties of CopperTable 1.1

Physical Properties Chemical Properties

reddish brown, metallic luster

easily shaped into sheets(malleable) and wires

(ductile)

good conductor of heatand electricity

density = 8.95 g/cm3

melting point = 10830C

boiling point = 25700C

slowly forms a basic blue-greensulfate in moist air

reacts with nitric acid and sulfuric acid

slowly form a deep-bluesolution in aqueous ammonia


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Figure 1.2 The physical states of matter.


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Sample Problem 1.1 Distinguishing Between Physical and Chemical Change

PROBLEM: Decide whether each of the following process is primarily a physical or a chemical change, and explain briefly:

PLAN: “Does the substance change composition or just change form?”

SOLUTION:

(a) Frost forms as the temperature drops on a humid winter night.

(b) A cornstalk grows from a seed that is watered and fertilized.

(c) Dynamite explodes to form a mixture of gases.

(d) Perspiration evaporates when you relax after jogging.

(e) A silver fork tarnishes slowly in air.

(a) physical change (b) chemical change (c) chemical change

(d) physical change (e) chemical change


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energy due to the position of the object or energy from a chemical reaction

Potential Energy

Kinetic Energy energy due to the motion of the object

Energy is the capacity to do work.

Potential and kinetic energy can be interconverted.


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Energy is the capacity to do work. Figure 1.3 A

less stable

more stable

change in potential energy EQUALSkinetic energy

A gravitational system. The potential energy gained when a lifted weight is converted to kinetic energy as the weight falls.


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Figure 1.3 B

less stable

more stable


A system of two balls attached by a spring. The potential energy gained by a stretched spring is converted to kinetic energy when the moving balls are released.



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Figure 1.3 C

less stable

more stable


A system of oppositely charged particles. The potential energy gained when the charges are separated is converted to kinetic energy as the attraction pulls these charges together.



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Figure 1.3D

less stable

more stable


A system of fuel and exhaust. A fuel is higher in chemical potential energy than the exhaust. As the fuel burns, some of its potential energy is converted to the kinetic energy of the moving car.



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Scientific Approach: Developing a Model

Observations : Natural phenomena and measured events; universally consistent ones can be stated as a natural law.

Hypothesis: Tentative proposal that explains observations.

Experiment: Procedure to test hypothesis; measures one variable at a time.

Model (Theory):Set of conceptual assumptions that explains data from accumulated experiments; predicts related phenomena.

Further Experiment: Tests predictions based on model.

revised if experiments do not support it

altered if predictions do not support it


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Chemical Arts and the Origins of Modern Chemistry

Alchemy

Medicine

Technology

http://www.levity.com/alchemy/brueghel.html

http://azothgallery.com/images/paracelsus.jpg

http://www.siamese-dream.com/page/siam1/CTGY/Hindu-Statues


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A Systematic Approach to Solving Chemistry Problems

•Problem statement

•Plan

Clarify the known and unknown.

Suggest steps from known to unknown.

Prepare a visual summary of steps.

•Solution

•Check

Comment and Follow-up Problem


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Sample Problem 1.2 Converting Units of Length

PROBLEM: To wire your stereo equipment, you need 325 centimeters (cm) of speaker wire that sells for $0.15/ft. What is the price of the wire?

PLAN:Known - length (in cm) of wire and cost per length ($/ft)

We have to convert cm to inches and inches to ft followed by finding the cost for the length in ft.

SOLUTION:length (cm) of wire

length (ft) of wire

length (in) of wire

Price ($) of wire

2.54 cm = 1 in

12 in = 1 ft

1 ft = $0.15

Length (in) = length (cm) x conversion factor

= 325 cm x in2.54 cm

= 128 in

Length (ft) = length (in) x conversion factor

= 128 in x ft12 in

= 10.7 ft

Price ($) = length (ft) x conversion factor

= 10.7 ft x $0.15ft

= $1.60


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Table 1. 2 SI Base Units

Physical Quantity (Dimension) Unit Name

Unit Abbreviation

mass

meter

kg

length

kilogram

m

time second s

temperature kelvin K

electric current ampere A

amount of substance mole mol

luminous intensity candela cd


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Table 1.3

Prefix PrefixSymbol

Word ConventionalNotation

ExponentialNotation

tera T trillion 1,000,000,000,000 1x1012

giga G billion 1,000,000,000 1x109

mega M million 1,000,000 1x106

kilo k thousand 1,000 1x103

hecto h hundred 100 1x102

deka da ten 10 1x101

----- ---- one 1 1x100

deci d tenth 0.1 1x10-1

centi c hundredth 0.01 1x10-2

milli m thousandth 0.001 1x10-3

micro millionth 0.000001 1x10-6

nano n billionth 0.000000001 1x10-9

pico p trillionth 0.000000000001 1x10-12

femto f quadrillionth 0.000000000000001 1x10-15

Common Decimal Prefixes Used with SI Units


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English EquivalentLength

1 kilometer (km)1000 (103) m

0.62 mi

Common SI-English Equivalent QuantitiesTable 1.4

Quantity SI UnitSI Equivalent

English to SI Equivalent

1 kilometer (km) 1000 (103) m

0.62 miles (mi)1 mi = 1.61 km

1 meter (m)100 (102) cm

1.094 yards (yd)1000 (103) mm

39.37 inches (in)1 yd = 0.9144 m

1 foot (ft) = 0.3048 m

1 centimeter (cm)0.01 (10-2) m

0.3937 in1 in = 2.54 cm

(exactly!)


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Volume

1,000,000 (106) cubic

centimeters 35.2 cubic feet (ft3)

1 cubic meter (m3)

1 ft3 = 0.0283 m3

1 cubic decimeter(dm3) 1000 cm3

0.2642 gallon (gal) 1.057 quarts (qt) 1 gal = 3.785 dm3

1 qt = 0.9464 dm3

1 cubic centimeter (cm3)0.001 dm3

0.0338 fluid ounce 1 qt = 946.4 cm3

1 fluid ounce = 29.6 cm3

English Equivalent



Common SI-English Equivalent QuantitiesTable 1.4 (continued)


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English Equivalent



Mass

1 kilogram (kg)

1000 grams

2,205 pounds (lb)1 (lb) = 0.4536 kg

1 gram (g)

1000 milligrams

0.03527 ounce(oz)

1 lb = 453.6 g

1 ounce = 28.35 g

Common SI-English Equivalent QuantitiesTable 1.4 (continued)


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Sample Problem 1.3 Converting Units of Volume

PROBLEM: The volume of an irregularly shaped solid can be determinedfrom the volume of water it displaces. A graduated cylinder contains 19.9 mL of water. When a small piece of galena, an ore of lead, is submerged in the water, the volume increases to24.5 mL. What is the volume of the piece of galena in cm3 and in L?

PLAN: The volume of galena is equal to the change in the water volume before and after submerging the solid.

volume (mL) before and after addition

volume (mL) of galena

volume (cm3) of galena

subtract

volume (L) of galena

1 mL = 1 cm3 1 mL = 10-3 L

SOLUTION:

(24.5 - 19.9) mL = volume of galena

4.6 mL xmL

1 cm3

= 4.6 cm3

4.6 mL xmL

10-3 L = 4.6x10-3 L


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Sample Problem 1.4 Converting Units of Mass

PROBLEM: International computer communications are often carried by optical fibers in cables laid along the ocean floor. If one strand of optical fiber weighs 1.19 x 10-3lbs/m, what is the total mass (in kg) of a cable made of six strands of optical fiber, each long enough to link New York and Paris (8.84 x 103km)?

PLAN: The sequence of steps may vary but essentially you have to find the length of the entire cable and convert it to mass.

length (km) of fiber

mass (lb) of fiber

length (m) of fiber

1 km = 103 m

mass (kg) of cablemass (lb) of cable

6 fibers = 1 cable

1 m = 1.19x10-3 lb

2.205 lb = 1 kg

SOLUTION:

8.84 x 103 km xkm

103 m

8.84 x 106m x m

1.19 x 10 -3 lbs

1.05 x 104 lb xcable

6 fibers

= 1.05 x 104 lb

= 8.84 x 106 m

2.205 lb

1 kgx

6.30 x 104 lb

cable

6.30x 104 lb =cable

2.86 x 104 kgcable

=


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Figure 1. 10

A Length B Volume

Some interesting quantities.

C Mass


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Densities of Some Common Substances*Table 1.5

Substance Physical State Density (g/cm3)

Hydrogen Gas 0.0000899

Oxygen Gas 0.00133

Grain alcohol Liquid 0. 789

Water Liquid 0.998

Table salt Solid 2.16

Aluminum Solid 2.70

Lead Solid 11.3

Gold Solid 19.3

*At room temperature (200C) and normal atmospheric pressure (1 atm).


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Sample Problem 1.5 Calculating Density from Mass and Length

PROBLEM: Lithium (Li) is a soft, gray solid that has the lowest densityof any metal. If a slab of Li weighs 1.49 x 103 mg and has sides that measure 20.9 mm by 11.1 mm by 11.9 mm, whatis the density of Li in g/cm3 ?

PLAN: Density is expressed in g/cm3 so we need the mass in grams and the volume in cm3.

mass (mg) of Li

lengths (mm) of sides

mass (g) of Li

density (g/cm3) of Li

103 mg = 1 g

10 mm = 1 cm

lengths (cm) of sides

volume (cm3)

multiply lengths

SOLUTION:

20.9 mm x

1mg

10-3 g

= 1.49 g1.49 x 103 mg x

10mm1cm

= 2.09 cm

Similarly the other sides will be 1.11 cm and 1.19 cm, respectively.

2.09 x 1.11 x 1.20 = 2.76 cm3

density of Li = 1.49 g

2.76 cm3= 0.540 g/cm3


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Figure 1.11

Some interesting temperatures.


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Figure 1.12 The freezing and boiling points of water.


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Temperature Scales and Interconversions

Kelvin ( K ) - The “Absolute temperature scale” begins at absolute zero and only has positive values.

Celsius ( oC ) - The temperature scale used by science, formally called centigrade, most commonly used scale around the world; water freezes at 0oC, and boils at 100oC.

Fahrenheit ( oF ) - Commonly used scale in the U.S. for our weather reports; water freezes at 32oF and boils at 212oF.

T (in K) = T (in oC) + 273.15 T (in oC) = T (in K) - 273.15

T (in oF) = 9/5 T (in oC) + 32T (in oC) = [ T (in oF) - 32 ] 5/9


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Sample Problem 1.6 Converting Units of Temperature

PROBLEM: A child has a body temperature of 38.70C.

PLAN: We have to convert 0C to 0F to find out if the child has a fever and we use the 0C to kelvin relationship to find the temperature in kelvins.

(a) If normal body temperature is 98.60F, does the child have a fever?(b) What is the child’s temperature in kelvins?

SOLUTION:

(a) Converting from 0C to 0F9

5(38.70C) + 32 = 101.70F

(b) Converting from 0C to K 38.70C + 273.15 = 311.8K


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The number of significant figures in a measurement depends upon the measuring device.

Figure 1.14

32.30C32.330C


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Rules for Determining Which Digits are Significant

All digits are significant

•Make sure that the measured quantity has a decimal point.•Start at the left of the number and move right until you reach the first nonzero digit.•Count that digit and every digit to it’s right as significant.

Numbers such as 5300 L are assumed to only have 2 significant figures. A terminal decimal point is often used to clarify the situation, but scientific notation is the best!

except zeros that are used only to position the decimal point.

Zeros that end a number and lie either after or before the decimal point are significant; thus 1.030 ml has four significant figures, and 5300. L has four significant figures also.


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Sample Problem 1.7 Determining the Number of Significant Figures

PROBLEM: For each of the following quantities, underline the zeros that are significant figures(sf), and determine the number of significant figures in each quantity. For (d) to (f), express each in exponential notation first.

PLAN: Determine the number of sf by counting digits and paying attention to the placement of zeros.

SOLUTION:

(b) 0.1044 g(a) 0.0030 L (c) 53,069 mL

(e) 57,600 s(d) 0.00004715 m (f) 0.0000007160 cm3

(b) 0.1044 g(a) 0.0030 L (c) 53.069 mL

(e) 57,600 s 5 sf(d) 0.00004715 m 4 sf (f) 0.0000007160 cm3 4 sf

2 sf 4 sf 5sf

(d) 4.715 x10-5 m 4 sf

(e) 5.7600x104 s 5 sf (f) 7.160x10-7 cm3 4 sf


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Rules for Significant Figures in Answers

1. For addition and subtraction. The answer has the same number of decimal places as there are in the measurement with the fewest decimal places.

106.78 mL = 106.8 mL

Example: subtracting two volumes

863.0879 mL = 863.1 mL

865.9 mL - 2.8121 mL

Example: adding two volumes 83.5 mL

+ 23.28 mL


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= 23.4225 cm3 = 23 cm39.2 cm x 6.8 cm x 0.3744 cm

2. For multiplication and division. The number

with the least certainty limits the certainty of the

result. Therefore, the answer contains the same

number of significant figures as there are in the

measurement with the fewest significant figures.

Rules for Significant Figures in Answers

Multiply the following numbers:


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Rules for Rounding Off Numbers1. If the digit removed is more than 5, the preceding number increases by 1. 5.379 rounds to 5.38 if three significant figures are retained and to 5.4 if two significant figures are retained.

2. If the digit removed is less than 5, the preceding number is unchanged. 0.2413 rounds to 0.241 if three significant figures are retained and to 0.24 if two significant figures are retained.

3.If the digit removed is 5, the preceding number increases by 1 if it is odd and remains unchanged if it is even.17.75 rounds to 17.8, but 17.65 rounds to 17.6. If the 5 is followed only by zeros, rule 3 is followed; if the 5 is followed by nonzeros, rule 1 is followed: 17.6500 rounds to 17.6, but 17.6513 rounds to 17.7

4. Be sure to carry two or more additional significant figures through a multistep calculation and round off only the final answer.


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Issues Concerning Significant Figures

graduated cylinder < buret ≤ pipet

numbers with no uncertainty

1000 mg = 1 g

60 min = 1 hr

These have as many significant digits as the calculation requires.

be sure to correlate with the problem

FIX function on some calculators

Electronic Calculators

Choice of Measuring Device

Exact Numbers


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Sample Problem 1.8 Significant Figures and Rounding

PROBLEM: Perform the following calculations and round the answer to the correct number of significant figures:

PLAN: In (a) we subtract before we divide; for (b) we are using an exact number.

SOLUTION:

7.085 cm

16.3521 cm2 - 1.448 cm2(a)

11.55 cm3

4.80x104 mg(b)

1 g

1000 mg

7.085 cm

16.3521 cm2 - 1.448 cm2(a) =

7.085 cm

14.904 cm2

= 2.104 cm

11.55 cm3

4.80x104 mg(b)

1 g

1000 mg=

48.0 g

11.55 cm3

= 4.16 g/ cm3


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Precision and Accuracy Errors in Scientific Measurements

Random Error - In the absence of systematic error, some values that are higher and some that are lower than the actual value.

Precision -Refers to reproducibility or how close the measurements are to each other.

Accuracy -Refers to how close a measurement is to the real value.

Systematic error - Values that are either all higher or all lower than the actual value.


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Figure 1.16

precise and accurate

precise but not accurate

Precision and accuracy in the laboratory.


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systematic error

random error

Precision and accuracy in the laboratory.Figure 1.16

continued

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Keys to the Study of Chemistry กุญแจในการเรียนวิชาเคมี