Chemistry: The Study of Change

Chapter 1

 

1.2

The Study of Chemistry

Macroscopic Microscopic

• Matter is anything that occupies space and has mass.

• Chemistry is the study of matter and thechanges it undergoes.

1.4

SugarWater

Gold

What about air? Yes it is matter

MAJOR AREAS OF CHEMISTRY• Organic Chemistry - the study of matter which is carbon-

based.

• Inorganic Chemistry - the study of matter containing all other elements (Inorganic is everything else).

• Analytical Chemistry - analyze matter to determine identity and composition (involves qualitative and quantitative) (eg. Determination of Alcohol in blood)

• Biochemistry - the study of life at the molecular level (DNA sequencing)

• Physical Chemistry - attempts to explain the way matter behaves (eg. Kinetics)

Chemistry Related Other Fields

• Medical Science

• Pharmaceutical Industry

• Oil Industry

• Paints and Coatings

Chemistry: A Science for the 21st Century

• Health and Medicine

• Sanitation systems

• Surgery with anesthesia

• Vaccines and antibiotics

•Energy and the Environment

• Fossil fuels

• Solar energy

• Nuclear energy

Chemistry: A Science for the 21st Century

• Materials and Technology

• Polymers, ceramics

• Food and Agriculture

• Genetically modified crops

• Specialized fertilizers

An element is a substance that cannot be separated into simpler substances by chemical means (Listed in the Periodic Table)

• 114 elements have been identified

• 82 elements occur naturally on Earth

gold, aluminum, lead, oxygen, carbon

• 32 elements have been created by scientists

technetium, americium, seaborgium

•Metals, nonmetals, metalloids

Latin names: aurum (Au), ferrum (Fe),natrium (Na)

Classification of Matter

1. Based on Physical State:

a. Gas: does not have a fixed shape or a fixed volume. Both the volume & the shape of the gas is that of the container.

b. Liquid: Has a fixed volume but variable shape = shape of the container.

c. Solid: Has fixed shape and volume

The Three States of Matter (water in three forms)

solidliquid

gas

Pure Substances

Classification of Pure Substances

1.Elements: These only have 1 type of atoms.

2. Compounds: composed of atoms of two or more elements. eg. H2O chemically united in a fixed proportion – 2 H atoms and 1 O atom. It can be separated only chemically.

Ammonia (NH3) – 1N, 3H (combined at high pressure)

Elements exist in monoatomic or polyatomic form. For eg., Hydrogen and oxygen exist as H2 and O2 , that is diatomic molecules where as Ca exist as monoatomic.

Chemical Symbols: Cobalt – Co, Sodium – Na (natium-Latin)CO is not cobalt, carbon monoxide

A mixture is a combination of two or more substances in which the substances retain their distinct identities.

1. Homogenous mixture – composition of the mixture is the same throughout.

2. Heterogeneous mixture – composition is not uniform throughout.

soft drink, milk, sugar solution

cement, iron filings in sand

Physical means can be used to separate a mixture into its pure components.

magnetdistillation

Physical and Chemical Properties

•Physical Properties: will keep the identity of the substance. Eg. boiling point (B.P), melting point (M.P), density, mass, volume, area. (water changes the physical state by heating and cooling, not the composition, so the M.P and B.P are physical properties)

•Chemical Properties: involves chemical change.eg. rusting, burning of H in O H2O

Needs a chemical change to bring the original substances, not physical like boiling or melting.

An extensive property of a material depends upon how much matter is being considered - additive

An intensive property of a material does not depend upon how much matter is is being considered. e.g., density of water 0.99713g/mL at 25 C

• mass (M)

• length

• volume (V)

• density – M/V

• color

Extensive and Intensive Properties

Hypothesis, Theory, Law

•A hypothesis is an educated guess, based on observation. Needs to prove by experiments to be true. •A scientific theory summarizes a hypothesis or group of hypotheses that have been supported with repeated testing. •A law generalizes a body of observations.

Eg., Newton's Law of Gravity F = Gm1m2/r2

Classifications of Matter

Measurement in Chemistry

• In SI units, not in English units.

• Meaurement is always necessary to collect data

• Mass, volume etc. with proper equipments or glassware

• Always use unit. Mass = 5 for a salt is useless. Needs unit like g or kg.– A measurement is useless without its units.– weight – force that gravity exerts on an object

International System of Units (SI)Revised Metric System

• Truly systematic

- 1 meter = 10 decimeters = 100 centimeters

Basic Units of the Metric System

Mass gram gLength meter mvolume liter L

Matter - anything that occupies space and has mass.

mass – measure of the quantity of matter

SI unit of mass is the kilogram (kg)

1 kg = 1000 g = 1 x 103 g

In Chemistry, the smaller g is more convenient.

weight – force that gravity exerts on an object

Volume – SI derived unit for volume is cubic meter (m3)

1 cm3 = (1 x 10-2 m)3 = 1 x 10-6 m3

1 dm3 = (1 x 10-1 m)3 = 1 x 10-3 m3

1 L = 1000 mL = 1000 cm3 = 1 dm3

1 mL = 1 cm3

Density – SI derived unit for density is kg/m3

1 g/cm3 = 1 g/mL = 1000 kg/m3

Chemical Application: g/cm3 or g/mL

density = mass

volume d = mV

A piece of platinum metal with a density of 21.5 g/cm3 has a volume of 4.49 cm3. What is its mass?

d = mV

m = d x V = 21.5 g/cm3 x 4.49 cm3 = 96.5 g

Derived Units in SI1. Area: (length x width); unit = m x m = m2

2. Volume: length x width x height; unit = m x m x m = m3

3. Density: mass/volume; unit = kg/m3

4. Speed: distance/time; unit = m/s5. Acceleration: speed/time; unit = m/s/s = m/s2

6. Force: mass x acceleration; unit = kg m/s2 = N (newton)7. Energy: force x distance; unit = kg m/s2 x m = kg m2/s2 = J

(joule)8. Pressure: force/Area; unit = kg m/s2 ÷ m2 = kg/ms2 = Pa

(Pascal)

K = 0C + 273.15

0F = x 0C + 3295

273 K = 0 0C 373 K = 100 0C

32 0F = 0 0C 212 0F = 100 0C

Conversions between Fahrenheit and Celsius

1.8

32-FC

oo

32C)(1.8F oo

1. Convert 75oC to oF.

2. Convert -10oF to oC.

1. Ans. 167 oF 2. Ans. -23oC

Convert 172.9 0F to degrees Celsius.

0F = x 0C + 3295

0F – 32 = x 0C95

x (0F – 32) = 0C95

0C = x (0F – 32)95

0C = x (172.9 – 32) = 78.395

Chemistry In Action

On 9/23/99, $125,000,000 Mars Climate Orbiter entered Mar’s atmosphere 100 km (62 miles) lower than planned and was destroyed by heat (read p.17, ch.1)

1 lb = 1 N

1 lb = 4.45 NF=ma=4.45Newton

“This is going to be the cautionary tale that will be embedded into introduction to the metric system in elementary school, high school, and college science courses till the end of time.”

Scientific NotationN x 10n

The number of atoms in 12 g of carbon:602,200,000,000,000,000,000,000

6.022 x 1023

The mass of a single carbon atom in grams:

0.0000000000000000000000199

1.99 x 10-23

N x 10n

N is a number between 1 and 10

n is a positive or negative integer

Scientific Notation568.762

n > 0

568.762 = 5.68762 x 102

move decimal left

0.00000772

n < 0

0.00000772 = 7.72 x 10-6

move decimal right

Addition or Subtraction

1. Write each quantity with the same exponent n

2. Combine N1 and N2 3. The exponent, n, remains

the same

4.31 x 104 + 3.9 x 103 =

4.31 x 104 + 0.39 x 104 =

4.70 x 104

Scientific Notation

Multiplication

1. Multiply N1 and N2

2. Add exponents n1 and n2

(4.0 x 10-5) x (7.0 x 103) =(4.0 x 7.0) x (10-5+3) =

28 x 10-2 =2.8 x 10-1

Division

1. Divide N1 and N2

2. Subtract exponents n1 and n2

8.5 x 104 ÷ 5.0 x 109 =(8.5 ÷ 5.0) x 104-9 =

1.7 x 10-5

Scientific Notation

• The measuring device determines the number of significant figures a measurement has.

• In this section you will learn – to determine the correct number of significant

figures (sig figs) to record in a measurement– to count the number of sig figs in a recorded value– to determine the number of sig figs that should be

retained in a calculation.

Copyright © 2001 The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

Significant FiguresFigure TA 1.2

Significant figures - all digits in a number representing data or results that are known with

certainty plus one uncertain digit.

Significant Figures

• Any digit that is not zero is significant

1.234 kg 4 significant figures

• Zeros between nonzero digits are significant

606 m 3 significant figures

• Zeros to the left of the first nonzero digit are not significant

0.08 L 1 significant figure

• If a number is greater than 1, then all zeros to the right of the decimal point are significant

2.0 mg 2 significant figures

• If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant

0.00420 g 3 significant figures

How many significant figures are in each of the following measurements?

24 mL 2 significant figures

3001 g 4 significant figures

0.0320 m3 3 significant figures

6.4 x 104 molecules 2 significant figures

560 kg (5.6 x 102 ) 2 significant figures

Significant Figures

Addition or Subtraction

The answer cannot have more digits to the right of the decimalpoint than any of the original numbers.

89.3321.1+

90.432 round off to 90.4

one digit after decimal point

3.70-2.91330.7867

two digit after decimal point

round off to 0.79

Scientific Notation

1.8

568.762

n > 0

568.762 = 5.68762 x 102

move decimal left

0.00000772

n < 0

0.00000772 = 7.72 x 10-6

move decimal right

Addition or Subtraction

1. Write each quantity with the same exponent n

2. Combine N1 and N2 3. The exponent, n, remains

the same

4.31 x 104 + 3.9 x 103 =

4.31 x 104 + 0.39 x 104 =

4.70 x 104

Significant Figures

Multiplication or Division

The number of significant figures in the result is set by the original number that has the smallest number of significant figures

4.51 x 3.6666 = 16.536366 = 16.5

3 sig figs round to3 sig figs

6.8 ÷ 112.04 = 0.0606926

2 sig figs round to2 sig figs

= 0.061

Significant Figures

Numbers from definitions or numbers of objects are consideredto have an infinite number of significant figures

The average of three measured lengths; 6.64, 6.68 and 6.70?

6.64 + 6.68 + 6.703

= 6.67333 = 6.67 = 7

Accuracy – how close a measurement is to the true value

Precision – how close a set of measurements are to each other

accurate&

precise

precisebut

not accurate

not accurate&

not precise

UNIT CONVERSION

• You need to be able to convert between units

- within the metric system

- between the English and metric system

• The method used for conversion is called the Factor-Label Method or Dimensional Analysis

ALWAYS NEEDED

Dimensional Analysis Method of Solving Problems

1. Determine which unit conversion factor(s) are needed

2. Carry units through calculation

3. If all units cancel except for the desired unit(s), then the problem was solved correctly.

given quantity x conversion factor = desired quantity

given unit x = desired unitdesired unit

given unit

• We use these two mathematical facts to do the factor label method

– a number divided by itself = 1– Example: 2/2 = 1

– any number times one gives that number back

– Example: 2 x 1 = 2• Example: How many donuts are in 3.5 dozen?

• You can probably do this in your head but let’s see how to do it using the Factor-Label Method.

Start with the given information...

3.5 dozen

Then set up your unit factor...

dozen 1

donuts 12

See that the units cancel...

Then multiply and divide all numbers...

= 42 donuts

Dimensional Analysis Method of Solving Problems

Conversion Unit 1 L = 1000 mL

1L

1000 mL1.63 L x = 1630 mL

1L1000 mL

1.63 L x = 0.001630L2

mL

How many mL are in 1.63 L?

The speed of sound in air is about 343 m/s. What is this speed in miles per hour?

1 mi = 1609 m 1 min = 60 s 1 hour = 60 min

343ms

x1 mi

1609 m

60 s

1 minx

60 min

1 hourx = 767

mihour

meters to miles

seconds to hours

conversion units

Common Relationships Used in the English System

A. Weight 1 pound = 16 ounces 1 ton = 2000 pounds

B. Length 1 foot = 12 inches 1 yard = 3 feet

1 mile = 5280 feetC. Volume 1 gallon = 4 quarts

1 quart = 2 pints 1 quart = 32 fluid ounces

Commonly Used “Bridging” Units for Intersystem conversion

Quantity English MetricMass 1 pound = 454 grams

2.2 pounds = 1 kilogramLength 1 inch = 2.54 centimeters

1 yard = 0.91 meterVolume 1 quart = 0.946 liter

1 gallon = 3.78 liters

1.3

Mea

sure

me

nt

in

Ch

em

istr

y 1. Convert 5.5 inches to millimeters

(139.7 mm)

2. Convert 50.0 milliliters to pints

(0.1057 pint)

3. Convert 1.8 in2 to cm2

(11.613 cm2)

1.7

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; cm3 for solids; and L for gases.

Derived Units

• The density of an object is its mass per unit volume,

Vm

d

A Density Example

• A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm3. What is the density of galena?

A Density Example

• A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm3. What is the density of galena?

Density = mass

volume

A Density Example

• A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm3. What is the density of galena?

Density = mass

volume=

12.4 g

1.64 cm3

A Density Example

• A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm3. What is the density of galena?

Density = mass

volume=

12.4 g

1.64 cm3= 7.5609 = 7.56 g/cm3

liquid mercury

brass nut

water

cork

What is the density of a sample of bone with mass of 12.0 grams and volume of 5.9 cm3?

A sinker of lead has a volume of 0.25 cm3. Calculate the mass in grams. The density of

lead is 11.3 g/cm3.

What is the volume of air in liters (density = 0.00129 g/mL) occupied by 1.0 grams.

WORKED EXAMPLES

Worked Example 1.1

Worked Example 1.2

Worked Example 1.4

Homework Problems

1.9,1.11,1.12,1.14,1.15,1.16,1.18,1.21,1.22,1.23,1.29,1.33,1.38,1.39,1.40,1.51,1.52,1.64,1.74,1.80