Daniel L. Reger Scott R. Goode David W. Ball Chapter 1 Introduction to Chemistry.

Daniel L. RegerScott R. GoodeDavid W. Ball

www.cengage.com/chemistry/reger

Chapter 1 Introduction to Chemistry

• Definitions• Science: Study of the natural universe

• Specifically, knowledge acquired by experience• Science is both an activity and the result of the

activity.• Chemistry: the study of matter and its interactions

with other matter and with energy.• Chemistry is how matter is

organized/reorganized/changed at the molecular level

The Nature of Science and Chemistry

Chemistry: The Central Science

• Chemistry is often called the central science because it is an essential component of the natural and life sciences.

Chemistry and Astronomy

• Elemental composition of stars can be determined by different wavelengths of visible light emitted.

• When starlight passes through a planets atmosphere, certain frequencies of light disappear because they are absorbed by compounds in the atmosphere.

http://cs.fit.edu/~wds/classes/cse5255/cse5255/davis/text.html

Chemistry and Geology

• Geochemistry: Study of the chemical composition of the earth• Chemical transformations in solids

• Ex. Polymorphism.• How limestone becomes marble.

http://geology.com/rocks/limestone.shtml; http://www.italartworld.com/

Chemistry and BiologyChemistry and Biology

• You reach a certain level in biology where processes can only be understood in terms of chemistry.

• Chemistry in biology explains:• Why your adrenaline levels increase when you are afraid or

excited • Why a body fails to produce insulin (diabetes)• Why cells become cancerous• Neurotransmitter (e.g., dopamine, norepinephrine) imbalances

that can produce:• Euphoria when you have a few beers or fall in love• Depression

http://cs.fit.edu/~wds/classes/cse5255/cse5255/davis/text.html

• Scientific method: investigations that are guided by theory and earlier experiments.

• Hypothesis: a possible explanation for an event.

• Law: a statement that summarizes a large number of observations.

• Theory: an explanation of the laws of nature.• In the realm of science, theory has a much narrower

and more rigorous meaning than in general.

The Scientific Method

• Matter: anything that has mass and occupies space.• Mass: the quantity of matter in an object.• Weight: the force of attraction between an object and other

objects.

Matter, Mass and Weight

Mass on moon and earth is the same.

Weight on moon and earth is the different.

• Matter can be described by different properties• Property: anything observed or measured about a

sample of matter.• Extensive property: depends on the size of the sample.

• mass, volume• Intensive property: independent of sample size.

• density, color, melting or boiling point• 2 samples with same intensive properties may be the

same material

Properties of Matter

• Physical properties: can be measured without changing the composition of the sample.

• mass, density, color, MP, BP, solubility• Physical change: a change that occurs without changing the

composition of the material.• freezing, melting, crushing a brick into powder

Physical Properties and Changes

Physical Change Chemical Change

• Chemical properties: describe the reactivity of a material.– Flammability (whether something is ignitable, flash point <

100 °F)– Combustibility (whether something will burn, flash point >

100 ° F)• Iron rusts

• Chemical change: at least part of the material is changed into a different kind of matter.

• Digestion of sugar is a chemical change• Acid/base neutralization

Chemical Properties

• Substances - Material that is chemically the same throughout.• cannot be separated into component parts by physical methods

• Two types of substances• Elements cannot be broken into simpler substances by

chemical methods• Table 1.1 (p. 09): Memorize these!!!

• Compounds can be separated into simpler substances (or elements) by chemical methods

• Always contain same elements in same proportions (H2O is always 11.2% H and 88.8% O)

Classification of Matter - Substances

• Mixture: matter that can be separated into simpler materials by physical methods.

• Heterogeneous mixture: composition of the mixture changes from one part to another.

• Chocolate chip cookies• Italian dressing

• Homogeneous mixture or solution: composition of the mixture is uniform throughout.

• Chocolate pudding• Sugar dissolves in water

• Alloy: a solution of a metal and another material (usually another metal).

Classification of Matter - Mixtures

What’s the Matter?

• Modern chemistry is largely based on experimental measurements. The confidence in measurements involves:

• Accuracy: agreement of a measurement with the true value.• Precision: agreement among repeated measurements of the

same quantity.

Accuracy and Precision

accurate and precise

precise but not accurate

accurate but not precise

neither accurate nor precise

Significant Figures• Significant figures are the numbers in a measurement that represent the

certainty of the measurement, plus one number representing an estimate.

Q: When is a number NOT significant?A: Look at the zerosLeading zeros are NOT significant. 0.00123Confined zeros ARE significant. 0.00103Trailing zeros ARE significant, when decimal visible 0.0012300But NOT significant if no decimal 12300

Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

Calculations with Significant Figures

Rules for Rounding•If the first nonsignificant figure to drop from your answer is ≥ 5, all nonsignificant figures dropped, last significant figure increased by 1.•If the first nonsignificant figure to drop from your answer is < 5, all nonsignificant figures dropped, last significant figure stays the same.

Exact Numbers

Numbers with no uncertainty, or are known values. Exact numbers do not change.

Ex. 1 foot is always = 12 inches. It will never be = 12.5 inches.

Not used to determine sig figs

•Na = 1 mol = 6.02 x 1023

•π= 3.142•1m = 1000 mm

• How many significant figures are present in each of the measured quantities?

• 0.00121062006900.01.00120.001060

Significant Figures

• How many significant figures are present in each of the measured quantities?

• 0.0012 2106 32006 4900.0 41.0012 50.001060 4

Significant Figures

• Trailing zeros in numbers without a decimal point may not be significant. Avoid ambiguity by using scientific notation.

• 100 1, 2 or 31 x 102 11.0 x 102 21.00 x 102 3

Significant Figures

•Determine the number of significant figures:

100. 100.030505 437,000125,904,000 4.80 x 10-3

4.800 x 10-3 0.0048

Test Your Skill

•Determine the number of significant figures

100. 100.030505 437,000125,904,000 4.80 x 10-3

4.800 x 10-3 0.0048•Answer:

100. 3 100.0 430505 5 437,000 3-6125,904,000 6-9 4.80 x 10-3 34.800 x 10-3 4 0.0048 2

Test Your Skill

Scientific Notation

• Scientific notation provides a convenient way to express very large or very small numbers.

• Numbers written in scientific notation consist of a product of two parts in the form M x 10n, where M is a number between 1 and 10 (but not equal to 10) and n is a positive or negative whole number.

• The number M is written with the decimal in the standard position.

Scientific Notation (continued)• STANDARD DECIMAL POSITION

• The standard position for a decimal is to the right of the first nonzero digit in the number M.

• SIGNIFICANCE OF THE EXPONENT n

• A positive n value indicates the number of places to the right of the standard position that the original decimal position is located.

• A negative n value indicates the number of places to the left of the standard position that the original decimal position is located.

Scientific ↔ Standard Notation• Converting from scientific notation to standard numbers

1.1 x 102 = 1.1 x 10 x 10 = 1.1 x 100 = 110 Decimal

1.1 x 10-2 = 1.1/ (10 x 10) = 1.1/100 = 0.011 Decimal

Converting ExponentsEx. 1 Ex. 2 0.67 x 10-5 0.067 x 10-4 0.0067 x 10-3 0.00067 x 10-2 0.000067 x 10-1 0.0000067 x 100 0.00000067 x 101 0.000000067 x 102

1.2 x 10-3 0.12 x 10-2 0.012 x 10-1 0.0012 x 100 0.00012 x 101

• When you move the decimal (l or r), the exponent will be equal to number of places you moved the decimal.

Standard to Scientific Notation

60023.5 345.233 -345.233 0.00345 0.10345 1.42

6.00235 × 104

3.45233 × 102

-3.45233 × 102

3.45 × 10-3

1.0345 × 10-1

1.42 × 100


• Sum (addition) or difference (subtraction) must contain the same number of places to the right of the decimal (prd) as the quantity in the calculation with the fewest number of places to the right of the decimal (i.e., the least accurate number).


prd 1 prd 1prd 3

8.10825.10 5.5 325.5

prd 1 prd 1prd 3

2.0175.0 5.5 325.5

Ex. 01

Ex. 02


Product (multiplication) or quotient (division) must have same number of sig figs as value with the fewest number of sig figs (least accurate number).


SF 2 SF 2SF 4

.194625.19 5.4 325.4

SF 2 SF 2SF 4

96.0169.0 5.4 325.4

Ex. 01

Ex. 02

• Determine accuracy in the same order as you complete the mathematical operations, # of significant digits are in red.

• density = 3.7 g/mL

3

2

2.79 g

8.34 mL - 7.58 mLv

m 2.79 g

0.76mL=

33

3

2

=

Mixed Operations

• Evaluate each expression to the correct number of significant figures.

(a) 4.184 x 100.620 x (25.27 - 24.16)

(b)

(c)

8.925 - 8.904x 100%

8.925

9.6 x 100.65

8.321+ 4.026

Test Your Skill

Answers: (a) 467; (b) 0.24%; (c) 1.2 x 102

Calculate each to the correct number of significant figures .0.1654 + 2.07 - 2.114 8.27 x (4.987 - 4.962)

9.5 + 4.1 + 2.8 + 3.175

4 (4 is exact)

x 100%9.025 - 9.024

9.025

Test Your Skill

Calculate each to the correct number of significant figures .0.1654 + 2.07 - 2.114 = 0.128.27 x (4.987 - 4.962) = 0.21

9.5 + 4.1 + 2.8 + 3.175

4 (4 is exact)

x 100%9.025 - 9.024

9.025

Test Your Skill

= 4.89

= 0.1%

Math Operations with Scientific Notation

Multiplication

Division

)10)(()10)(10( 2424 baba

)10()10(

)10( 242

4

b

a

b

a

Addition/Subtraction Convert numbers to the same exponents

(5.00 x 102) + (6.01 x 103) =(0.500 x 103) + (6.01 x 103) = (5.00 x 102) + (60.10 x 102) = (5.00 + 60.10) x 102 = (65.10 x 102) = 6510

(6.01 x 103) - (5.00 x 102) =(6.01 x 103) - (0.500 x 103) = (60.10 x 102) - (5.00 x 102) = (60.10 - 5.00) x 102 = (55.10 x 102) = 5510

Examples of Math Operations

Multiplicationa. (8.2 X 10-3)(1.1 X 10-2) = (8.2 X 1.1)(10(-3+(-2))) = 9.0 X 10-5 b. (2.7 X 102)(5.1 X 104) = (2.7 X 5.1)(102+4) = 13.77 X 106 Now change to Scientific Notation 1.4 X 107

Division a. 3.1 X 10-3 = (3.1/1.2)(10-3-2) = 2.6 X 10-5

1.2 X 102

b. 7.9 X 104 = (7.9/3.6)(104-2) = 2.2 X 102

3.6 X 102

Adding/Subtracting 3.05 X 103 + 2.95 X 103 = (3.05 + 2.95)(103) = 6.0 X 103

Quantity Unit Abbreviation

Length meter m

Mass kilogram kg

Time second s

Temperature kelvin K

Amount mole mol

Electric current ampere A

Luminous intensity candela cd

Base Units in the SI

Prefix Abbreviation Meaning

mega- M 106

kilo- k 103

centi- c 10-2

milli- m 10-3

micro- 10-6

nano- n 10-9

pico- p 10-12

Common Prefixes Used With SI Units

• Unit conversion factor: a fraction in which the numerator is a quantity equal or equivalent to the quantity in the denominator, but expressed in different units

• The relationship 1 kg = 1000 g• Generates two unit conversion factors:

1kg

g 1000 and

g 1000

kg 1

Unit Conversion Factors

What would be some other examples?

• Volume is the product of three lengths.• The standard unit of volume is the cubic meter

(m3).100 cm = 1 m(100 cm)3 = (1 m)3

106 cm3 = 1 m3

• Two important non-SI units of volume are the liter and milliliter.1 liter (L) = 1000 mL = 1000 cm3

1 mL = 1 cm3

Conversion Among Derived Units

Volumes can be expressed in different units depending on the size of the object.

1 m3 contains1000 L

1 L contains1000 mL

1 mL = 1 cm3 or 1 cc

Volume

• Express a volume of 1.250 L in mL, cm3, and m3

33-36

3

33

m10 1,250cm 10

m 1 L 1.250

cm 1,250L 1cm 1000

L 1.250

mL 1,250L 1mL 1000

L 1.250

Using Unit Conversions

Factor Unit Method Examples

• A length of rope is measured to be 1834 cm. How many meters is this?

• Solution: • Write down known quantity (1834 cm). • Set known quantity = units of the unknown quantity (meters). • Use factor (100 cm = 1 m), to cancel units of known quantity

(cm) and generate units of the unknown quantity (m). • Do the math.

m 34.18cm 100

m 1cm 1834

m cm 1834


Factor Unit Method Example

Q: If an arrow shot from a bow travels 30 yards in 1 second, many cm does it travel in 4 seconds?

•Time = 4 s•Rate = 30 yards/sec•1 yard = 3 feet•1 foot = 12 in.•1 in = 2.54 cm

cm ?in 1

cm 2,54

foot 1

in 12

yard 1

feet 3

s 1

yards 30 s 4


• Density: mass per unit volume

•

• Density, in SI base units, is kg/m3 (kg m-3).

• Most commonly used density units are g/cm3 (g cm-

3 or g/mL) for solids and liquids, and g/L for gases.

Vm

d

Density

• The density of Ti is 4.50 g/cm3 or 4.50 g = 1 cm3

• Calculate the volume of 7.20 g Ti.

Conversions Between Equivalent Units

• The density of Ti is 4.50 g/cm3 or 4.50 g = 1 cm3.

• Calculate the volume of 7.20 g Ti.

• Ti cm 1.60g 4.50

cm 1 Ti g 7.20 3

3

Conversions Between Equivalent Units

Percentage

• The word percentage means per one hundred. It is the number of items in a group of 100 such items.

• PERCENTAGE CALCULATIONS• Percentages are calculated using the equation:

• In this equation, part represents the number of specific items included in the total number of items.

100whole

part%


Percentage Calculation

• A student counts the money she has left until pay day and finds she has $36.48. Before payday, she has to pay an outstanding bill of $15.67. What percentage of her money must be used to pay the bill?

• Solution: Her total amount of money is $36.48, and the part is what she has to pay or $15.67. The percentage of her total is calculated as follows:

%96.4210048.36

67.15100

whole

part%

Density Calculation

• A 20.00 mL sample of liquid is put into an empty beaker that had a mass of 31.447 g. The beaker and contained liquid were weighed and had a mass of 55.891 g. Calculate the density of the liquid in g/mL.

• Solution: The mass of the liquid is the difference between the mass of the beaker with contained liquid, and the mass of the empty beaker or 55.891g -31.447 g = 24.444 g. The density of the liquid is calculated as follows:

mL

g222.1

mL 20.00

g 444.24

v

md

Energy CalculationsQ: In order to lose 1 lb/week, you need to cut 500.0 Cal from your diet each day, or use an equivalent number of joules by working each day. How many equivalent joules would you have to spend on work to achieve this each day? How many joules would you have to expend to achieve this over 7 days?

To answer this, you need to first know the following:• 1Cal = 1 kcal = 1000 scientific calories or 1 nutritional calorie• 1 scientific calorie = 1 cal = 4.184 J

Temperature Scales

• The three most commonly-used temperature scales are the Fahrenheit, Celsius and Kelvin scales.

• The Celsius and Kelvin scales are used in scientific work.


Temperature Conversions• Readings on one temperature scale can be converted to the

other scales by using mathematical equations.• Converting Fahrenheit to Celsius.

• Converting Celsius to Fahrenheit.

• Converting Kelvin to Celsius.

• Converting Celsius to Kelvin.

32F9

5C

32C5

9F

273KC

273CK


• Express 17.5°C in °F and in K.

Test Your Skill

• Express 17.5°C in °F and in K.

• Answer:TF = 63.5° F; TK = 290.6 K

Test Your Skill

Useful Conversions/Constants

• Na = 1 mol = 6.02 x 1023

• π= (Circumference of circle/diameter of circle) = 3.142

• 1 scientific calorie = 4.184 J• 1Cal = 1 kcal = 1000 scientific calories or

1 nutritional calorie• 1 N = (kg)(m)/s2

• g = 9.81 m/s2 (earth’s gravity)

Chapter 1 Visual Summary

Daniel L. Reger Scott R. Goode David W. Ball Chapter 1 Introduction to Chemistry.

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