Chemistry: The Study of ChangeChemistry: The Study of Change

Chapter OneChapter One

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What is Chemistry?The science of studying matter, its composition, properties, the changes it undergoes, and the energy associate with these changes.

Chemistry is called the central science, because a basic knowledge of chemistry is essential for students of biology, physics, geology…..and many other subjects

Why study ChemistryWhy study Chemistry•Explain the natural worldExplain the natural world ,….Why? ,….Why? e.g.: -how the battery of your car works, e.g.: -how the battery of your car works, - why the color of leaves of trees change in autumn….ect.- why the color of leaves of trees change in autumn….ect.

you can answer these questions and others by understanding some you can answer these questions and others by understanding some chemistrychemistry

you are able to read and understand this sentence because chemical you are able to read and understand this sentence because chemical reactionsreactions occur in your brain.occur in your brain.

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The food you ate is furnishing energy through The food you ate is furnishing energy through chemical reactionschemical reactions Trees and grass grow because of chemical reactionsTrees and grass grow because of chemical reactions•Prepare for a careerPrepare for a career

IndirectlyIndirectly - problem solving and thinking skills- problem solving and thinking skills•AgricultureAgriculture

ProductionProduction - fertilizers, soil tests- fertilizers, soil testsProtectionProtection – pesticide, herbicide – pesticide, herbicide

•MedicineMedicineDrugsDrugsMaterials- hips, artificial skinMaterials- hips, artificial skinBiotechnology- using organisms as a Biotechnology- using organisms as a means of means of

productionproduction

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•Environment- PollutionEnvironment- PollutionEliminate sourcesEliminate sourcesTreatment once pollutedTreatment once polluted

•AstronomyAstronomy Remote analysis of stars from their Remote analysis of stars from their lightlight Analysis of extraterrestrial samplesAnalysis of extraterrestrial samples

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Scientific method science: is a framework of gaining and organizing knowledge

science: - a set of facts - a plan of action ( a procedure for processing and understanding certain type of information)

Scientific thinking is useful in all aspect of life. The process that lies at the center of scientific inquiry is called the scientific method.

The scientific method: A way of solving problems and answering question, or a systematic approach to research

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Steps of scientific methods:1. Observations- what is seen or measured

it can be

qualitative quantitative (general observations (comprising numbers obtained about the system; by various measurements of system; water is liquid, sky is blue) water is boil at 100°C)

2. Hypothesis- a possible explanation of observation, based on research and previous knowledge.

3. Experiment- designed to test hypothesisonly two possible answers: hypothesis is right…….hypothesis is wrong…..

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Experiment always produce new observations and this bring the process back to the beginning again…generate new data observations from experiment modify hypothesis repeat the cycle.The cycle repeat many times by you and by others, hypothesis gets more and more certain, becomes a theory.Theory: a set of tested hypotheses that gives an overall explanation of some natural phenomenon( explain why things behave a certain way). Theories have predictive value.The true test of a theory is if it can predict new behaviors.If the prediction is wrong, the theory must be changed.Theory can never be proven, it is the best explanation.

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if the same observations applies to many different systems…scientific law is developed.Law: describe how things behaves, summaries of observation (usually mathematical equations between phenomena that is always the under the same conditions)…equation of how things changeSo law…how theory….why

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Observations

Hypothesis

Experiment

Law

Theory(Model)

Prediction

Experiment

Modify

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Classifications of Matter

Matter: anything that occupies space and has mass.

There are some fundamental ways in which matter is classified, the most important of these are:

1) according to its physical state:

All substances can exist in three states, which are liquid, solid, and

gas.

gas state:gas state: - no fixed volume and shape

- take the shape and volume of the container

- can be compressed

- their molecules far apart and move fast

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liquid state: - has distinct volume

- no specific shape (take the shape of the container)

- can not be compressed

- close to each other but are not held so rigidly in position and can move past one another

solid state: - has definite shape and volume

- rigid in shape.

- can not be compressed

- held close to each other with little freedom of motion

The following figure shows the three states of matter, it can be interconvert without changing its composition

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Solid Liquid Gas

Melt Evaporate

CondenseFreeze

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2) According to its composition:

Pure substance mixture

elements compoundsHomogenous

mixtureHeterogeneous

mixture

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Elements:

- simplest kind of matter, cannot be separated into simpler substances by chemical means.

- building blocks for all complex

- about 117 elements are known. Most of them occur naturally on earth. The others have been created by scientists via nuclear process.

- building blocks of elements are atoms

- the same type of elements combine to for molecules

e.g.: O2 , Cl2 , N2 …ect

Compounds:

- Forms due to interaction of two or more elements chemically united in fixed proportions. e.g.: H2O, NaCl , Na2CO3 ….ect

- Unlike mixtures, compounds can be separated only by chemical means into their pure components

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mixture:� A combination of two or more substances.� Variable composition.� each substance retain its own chemical identity and its own

properties. Examples: air, soft drink, milk …

a) Heterogeneous mixture

- mixture is not the same from place to place.

- consist of two phases - e.g.: soil, oil and water

b) Homogeneous mixture (solution) - same composition throughout. therefore uniform properties

- Every part keeps its properties.

- consist of a single phase. - e.g.: air, salt dissolved in H2O

Any mixture can be separated by physical means into pure components without changing the identities of the components.

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Physical and Chemical Properties of matterEvery substance has a unique set of properties that allow us

to recognize it and distinguish it from other substances.

e.g: water: state: liquid, m.p.=0°C , b.p=100°C , density=1g/ml , flammability: no

Properties of matter can be grouped into two categories:

1) physical properties

2) chemical properties

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Physical properties: It is the one that can be measured or observed without changing

the composition or identity of substance. e.g.: color, m.p., b.p., state of matter, viscosity …ect these physical can be divided to: Extensive Properties - only depends on the amount of matter,

e.g: mass, volume, length as the amount of substance increase, these properties increaseValues of the same extensive property can be added together. Intensive Properties - only depends on the type of matter, not the amount. e.g.: state of matter, color, melting point, boiling point . … do not depend on the amount

of substance.Values of the same intensive property are not additive

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The ratio of two extensive properties will give an intensive properties.

example: the density (d), which is the ratio of mass to volume

d =

� The SI- derived unit for density is kg/m3. This unit is large for most chemical applications. Therefore, (g/cm3) and its equivalent (g/mL) are used for solid and liquid densities. Because gas densities are very low, we express them in (g/L)

� Useful for identifying a compound� Useful for predicting weight

mV

Extensive property

Extensive property

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1.0 g/cm3 = 1.0 g/mL = 1000kg/m3

1.0 g/ L = 10-3 g/mL

The following table lists the densities of several substances:

substance Density (g/mL(Air .0 001

Ethanol .0 79Water .1 00

Mercury .13 6Table salt .2 2

Iron .7 9Gold .19 3

Osmium .22 6

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Example:

� carbon tetrachloride (density 1.53 g/cm3 ) weighs 161.9 g. What is the volume of the carbon tetrachloride ?

d = m / V V = m / d

V = 161.9 g / 1.53 (g/cm3)

V = 105.82 cm3

Example:

� 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

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physical changes: A change that changes appearances, without changing the

composition. e.g: water change from solid to liquid to vapor

Chemical properties: is the tendency of a substance to undergo a particular chemical reaction.

e.g: flammability, reactivity toward acids

Chemical changes A change where a new form of matter is formed. Also called

chemical reaction. Alter not only the physical appearance but also the chemical makeup

as well

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Units of measurements

Making observations is fundamental of all science.

Quantitative observation or measurement, always consists of two parts:

- Number - Scale (unit)

Any measured value must contain both parts to be meaningful. Two major systems of measurements are adopted in different parts of the world: 1)English system: used in U.S.A. length…yard, feet, inch. mass….tons, pounds, ounces volume…gallon, quarts, pints2)International System of Units (abbreviated SI)

(le System International) based on the metric system. used by most of the rest of the world and in science.

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The SI system specify a set of seven basic units which are:

Physical quantity Name of unit symbol

Mass kilogram kg

Length meter m

Time second s

Temperature Kelvin K

Electric current ampere AAmount of substance mole molLuminous intensity candela cd

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Derived units

The SI units seems to be very limited, there are many quantities such as area, volume, speed whose units don’t appear in the previous table.

In the SI system, units of such quantities are obtained by appropriate combination of the base unit and called the derived unit.

Examples:Area = length x width = m (base unit) x m (base unit) = m2

speed = distance/time = m (base unit)/s (base unit) = m/s force = mass x acceleration

= mass x [distance/(time)2 ] = kg x m/s2

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Because the fundamental units are not always convenient (e.g: expressing the mass of a pen in kg), prefixes are used to change the size of the unit.

Some of these prefixes a listed in the following table:

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prefix symbol meaning Exponential notation

mega M 1,000,000 106

kilo k 1,000 103

hecto h 100 102

deka da 10 101

- - 1 100

deci d 0.1 10-1

centi c 0.01 10-2

milli m 0.001 10-3

micro µ 0.000001 10-6

nano n 0.000000001 10-9

pico p 0.000000000001 10-12

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prefix factor example

kilo 1000 (103) 1.0kilometer (km)= 103m1.0 kilogram (kg) = 103g

milli 1/1000 (10-3) 1.0millisecond(ms)=10-3s1.0milligram (mg) = 10-3g

centi 1/100 (10-2) 12.5 cm = 12.5x10-2m

pico 10-12 7.91x109 pg =7.91x109x10-12g =7.91x10-3g

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Uncertainty in measurementThe number associated with a measurement is obtained using some measuring device.

For example, consider the measurement of the length of an object.

21 3 4 5

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Note that the length of the object occurs at about 4.55 cm.note that we must estimate the last number by interpolating between the 0.1 cm marks. Since the last number is estimated, its value may be different if another person makes the same measurement. If several persons make the same measurement, they obtain the following values: 4.54 cm, 4.56 cm,..ect.These results show that the first two numbers (4.5) remain the same ..called certain digits The digit to the right of the (5) must be estimated and therefore varies; it is called an uncertain digit.

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Any measured value contain : certain digits + uncertain digit. These numbers are called the significant figures of a measurement.Measurement always has some degree of uncertainty,

which depends on the precision of the measuring device. Example: consider the measuring the same object by using the following meter… its length will be 4.6 or4.5 or 4.7 or 4.4 …( 4 is certain), (6, 5, 7, 4 uncertain)

21 3 4 5

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The convention of sig. fig. automatically indicate something about the uncertainty in measurement. The uncertainty in the last number is usually assumed to be ± (smallest marks /2)In the previous examples: ± 0.1/2 =0.05

± 1/2 = 0.5 example: what is the length of the following object

100 200150 25050

142 (± 5) cm

certain uncertain

uncertainty

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Significant figures and calculations

Rules for counting sig. fig.:1) Nonzero integers: always count as sig. fig.

e.g: 234 cm 3 sig. fig., 4.742 g 4 sig. fig.2) Zeros: there are three classes of zeros

a) leading zeros: zeros that precede the nonzero digits, do not count as sig. fig., their purpose is to indicate the placement of decimal point e.g.: 0.23 2 sig. fig.; 0.00432 3 sig. fig.b) captive zeros: zeros between nonzero digits

count as sig. fig. e.g. 1.08 3 sig. fig. ; 3.0502 5 sig. fig.

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c) trailing zeros: zeros at the right end of the number.

They are significant only if the number contains a

decimal point e.g.: 2.0 mg 2 sig. fig.

For numbers that do not contain decimal point, (zeros after the last nonzero digit) may or may not be significant,

400. 3 sig. fig. , 400 1 sig. fig.

3) Exact numbers: determine by

- counting: 3 apples, 10 students…

- definition: 1 inch = 2.54 cm

they do not considered in as sig. fig.

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Scientific notation:

If the length of an object is 150 cm, is the zero consider as sig. fig. or not ?

It depends on the measuring device used as shown in the following example:

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100 200150 25050

100 200

140 cm

certain uncertain

(3 sig. fig.)

140 cm

certainuncertain

(2 sig. fig.)

Uncertainty ±5

Uncertainty ±50

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Note the number 140 can be written in exponential notation :

the first measuring device : 1.40 x 102 (3 sig. fig.) the second measuring device: 1.4 x 102 (2 sig. fig.)This type of notation has at least two advantages - indicate the no. of sig. fig. easily - fewer zeros are needed to write a very large and a very

small number.e.g.: 0.000060 represent as 6.0 x 10-5

660000 represent as 6.6 x 105 ( 2 sig. fig.) 6.60 x 105 (3 sig. fig.)

6.600 x 105 (4 sig. fig.)

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Problem500 is only 1 significant figure.

if it really has three, how can I write it?

Three ways to write it:� 500.� 500 (± 5) � In scientific notation: 5.00 x 102

now the zero counts.

If it contain two sig. fig.:

5.0 x102 or 500 (± 20)

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Example(1): How many sig figs in the following measurements?

� 458 g (3)� 4085 g (4)� 4850 g (3)� 0.0485 g (3)� 0.004085 g (4)� 40.004085 g (8)

� 405.0 g (4)� 4050. g (4)� 0.450 g (3)� 4050.05 g (6)� 0.0500060 g (6)

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Example (2): write the following measurements in scientific notation to the correct number of sig. fig.

- 67000 (to 3 sig. fig.) = 6.70 x 104

- 67000 (to 2 sig. fig.) = 6.7 x 104

- 0.00009010 = 9.010 x 105

Example (3): (3): how many sig. fig. in the following number

- 896000 (± 100), it contains (4 sig. fig.)

uncertaintyScientific notation= 8.960 x 105

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Rules for sig. fig. in mathematical operations

� For addition and subtraction:

The result has the same number of decimal places as the least precise measurement used in calculation. For example:

12.11

18.0 one decimal places

1.013

31.123 correct 31.1

e.g: 27.93 + 6.4 = 34.3

45.86 – 32.467 = 13.393 = 13.40

+

+

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� For multiplication and division:

The number of sig.fig. in the result is the same as the number in the least precise measurement used in the calculation. For example:

4.56 x 1.4 = 6.38 correct 6.4

e.g : 6.56 x 8.215 = 53.8904 = 54.0

(3.97 x 5.870) / 2.1 = 0.332057… = 0.33

3 sig. fig. 2 sig. fig. 2 sig. fig.The product should have only two sig. fig.

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Practice� 4.8 + 6.8765 = 11.7 � 520 + 94.98 = 615� 0.0045 + 2.113= 2.118� 500 -126 = 374� 6.0 x 103 - 3.8 x 102 = 5.6 x 103

� 6.0 x 10-2 - 3.8 x 10-3 = 5.6 x 10-2

� 5.33 x 1022 - 3.8 x 1021 = 5.0 x 1022

� 4.5 / 6.245 = 0.72� 4.5 x 6.245 = 28� 9.8764 x 0.043 = 229.6837… = 230 or 2.3 x 102

� 3.876 / 1980 = 0.001957575… = 0.001958

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Practice� 4.184 x 100.62 x (25.27 – 24.16) =

4.184 x 100.62 x 1.11 = 467.3034288

= 467

• 9.2 x 100.658.321 + 4.026

= 74.996355.. = 75

• 1.00866 – 1.00728

6.02205 x 1023= = 2.29 x 10-270.00138

6.02205 x 1023

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Rounding rules

� Look at the number behind the one you’re rounding.

� If it is 0 to 4 don’t change it.� If it is 5 to 9 make it one bigger.� Round 45.462 to four sig figs.:� to three sig figs.:� to two sig figs.:� to one sig figs.:

45.46

45.5

4550

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Units for lab. measurements

In chemistry, it is necessary to measure mass, volume, length, and temperature.

Length: The SI unit is meter (m), we use smaller units for length in lab. which are cm and mm.

1 cm = 10-2 m or 1 m = 100 cm 1mm = 10-3 m or 1m = 1000mm

1 cm = 10 mm

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Mass and Weight� Mass: amount of matter in an object.

Mass is a measure of the resistance of an object to change in its state of motion

� Weight: the response of mass to gravity.

weight measure the force with which the object of a given mass is attracted by gravity.

� Sometimes used interchangeably� Mass can’t change, weight can� Your mass on the earth is the same as your mass on the moon, but

your weight will be differ.� SI unit for mass is kg (the only base unit that contain prefix, kilo).� In the lab. We use the unit of gram, (g) (1.0 kg = 103 g)

we use the balance to measure the mass.

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� Volume: it is not a fundamental SI unit, it is a derived unit from the length.

A cube that measure 1 m of each edge has a volume = 1 m3

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1 m = 10 dm

1 m3 = (10)3 dm3 = 103 dm3

Another common unit of volume is the Liter (L): which is the volume occupied by 1.0 dm3

1 L = 1 dm3 = 10-3 m3

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

in the same way:

1dm = 10 cm 1 dm3 = 103 cm3

cm3 = ml so 1 dm3 = 103 ml

also 1 L = 103 ml

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•TemperatureIt determine the direction of heat flow spontaneously. Heat

always flows spontaneously from a substance of higher temp. to one of lower temp.

Temp. is measured by a device called thermometer.

Three temp. scale employed in scientific studies:1) Celsius scale: (known as centigrade) (°C)

two reference temp.’s are chosen to make marking on the scale of thermometer which are the m.p.(0°C) and the b.p.(100°C) of water

difference between m.p and b.p. = 100

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2) Fahrenheit scale (°F): used in U.S.A.

m.p. of water = 32°F b.p = 212°F

difference between m.p and b.p. = 180

100°C = 180°F 1°C =(180/100)°F

1°C = (5/9)°F

So:

T(°C) = (T(°F) – 32) (5/9)

T(°F) = T(°C) x (9/5) + 32

Celsius degree is nearly twice as larger as degree in Fahrenheit scale

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3) Kelvin scale (K): it is the SI base unit for temp.

at Kelvin scale m.p. for water = 273.15K

b.p. for water = 373.15K

difference between m.p and b.p. = 100

T(K) = T(°C) + 273.15

in K-scale all temp. have positive values

Zero point in Kelvin scale (0K) is called the absolute zero 0 K = - 273.15 °C

0 K is lowest temp. that can be reached

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Fahrenheit Celsius Kelvin

- 32°F - 0°C - 273.15K

- 212°F - 100°C - 373.15 K

Melting point

of water

180°F

Boiling point of water

100°C 100 K

- - 40°F - - 40°C - 233.15 K

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Example: a) Convert 172.9 0F to degrees Celsius.0F = x 0C + 329

50F – 32 = x 0C9

5x (0F – 32) = 0C

95

0C = x (0F – 32)95

0C = x (172.9 – 32) = 78.3ºC95

b) Convert 172.9 0F to Kelvin

T (K) = 78.3ºC + 273 = 315.3 K

Dimensional Analysis

(Using the units to solve problems)

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Dimensional analysis

Dimension: unitAnalyze: solveSo dimensional analysis is use the units to solve problems. Use the conversion factors to change the unit.Conversion factor: a ratio of equivalent measurements.Start with two things that are the same

1 cm = 10-2 m

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Dived each side to come up with two ways of writing the number 1.0

= = 1

= = 1

So : = 1 =

1 cm10-2 m

10-2 m10-2 m

10-2 m1 cm

10-2 m10-2 m

1 cm10-2 m

10-2 m1 cm

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Choose the conversion factor that get rid of the unit you do not want.

Conversion factors = 1

1 kilogram = 103 gram (equivalent statement)

= 1 =

� There are 2 conversion factors.� Multiply by the one that will give you the correct unit in your

answer.

1 kg103 g

103 g1 kg

� given quantity x conversion factor = desired quantity

desired unitgiven unit

given unit x = desired unit

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� Examples:Perform the following conversions

1. 172 mm to m.

1 mm = 10-3 m

conversion factors: or

172 mm x = 172 x 10-3 mm

1 mm10-3 m

10-3 m1 mm

10-3 m1 mm

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2. 172 cm to dm.

cm m dm

1 cm = 10-2 m 1 dm = 10-1 m

conversion factors (cm to m ): or

conversion factors (m to dm): or

172 cm x x = 172 x 10-1 dm

1 cm10-2 m

10-2 m1 cm

1 dm10-1 m

10-1 m1 dm

10-2 m1 cm

1 dm10-1 m

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3. 16.85 feet to inches1 foot = 12 inches (equivalent statements)

= 1 = (2 conversion factors)

16.85 feet x = 202.2 inch.

4. 2.85 cm to inches

2.85 cm x = 1.12 inch.

1 foot

12 inch12 inches

1 foot

12 inch.1 foot

1 inch. = 2.54 cm

1 inch.2.54 cm

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Example:A race is 1.25 miles. How long is the race in meters, and kilometers?

mile yard meter km

race in (m) = 1.25 mile x x =2011 m

race in (km) = 2011 m x =

1 m = 1.094 yd 1 mile = 1760 yd 1 km = 103 m

1 m1.o94 yd1 mile

1760 yd

1 km103 m

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Units to a Power� How many cm2 is 15 m2?

1 cm = 10-2 m

1 cm x 1 cm = 10-2 m x 10-2 m

1 cm2 = 10-4 m2

� How many m3 is 1500 cm3?

1 cm = 10-2 m

1 cm x 1 cm x 1 cm= 10-2m x 10-2 m x 10-2 m

1 cm3 = 10-6 m3

15 m2 x 1 cm2

10-4 m2= 15 x 104 cm2

1500 cm3 x 10-6 m3

1 cm3 = 1500 x 10-6 m3 1.500 x 10-3 m3=

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� how many mm3 are in 36 cm3 ?

1cm = 10-2 m 1cm3 = 10-6 m3

1 mm = 10-3 m 1 mm3 = 10-9 m3

36 cm3 x x =1 cm3

10-6 m3 1 mm3

10-9 m336 x 103 mm3

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Multiple unitsExample (1)The speed limit is 65 mi/hr. What is this in

m/s?1 mile = 1760 yds 1 meter = 1.094 yds

65 mihr

1760 yd1 mi 1.094 yd

1 m 1 hr60 min

1 min60 s

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Example (2): A particle has a velocity of 6.51 x105 cm/h what is its velocity in mile/s.

velocity (mile/s) = 6.51x105 x x x

x x

1 m = 1.094 yd 1 mile = 1760 yd 1 cm = 10-2m

1 h = 60 min. 1 min. = 60 s

cmh

10-2 mcm 1 m

1.094 yd

1760 yd1 mile 1 h

60 min.1 min.60 s

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