Chapter 1 Jan12

7/29/2019 Chapter 1 Jan12

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1.1 Introduction

1.2 Classification of Matter1.3 Properties of Matter

1.4 Units of Measurement

1.5 Uncertainty in Measurement

1.6 Dimensional Analysis


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Able to differentiate between the states

of .Able to distinguish between ,

and .

Able to distinguish between and

properties.Able to use and convert different of

measurement.


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is the study of.

Central role in science and technology.Has a high impact on our daily living, e.g.

health and medicine, energy and environment,

materials and technology and food andagriculture.

Able to contribute to problem solving analysis.


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Physical material - anything that has and

.

Matter can be classified according to its:Physical state (solid, liquid or gas)

Composition (element, compound or mixture)


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Classification

of Matter

Physical State Composition

Gas

Liquid

Solid

Puresubstance

Mixture

Element

Compound

Homogeneous

Heterogeneous


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no fixed volume/shape

easy to compress/expandmolecules are far apart

move at high speed

often collide


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volume independent of

containerslightly compressible

molecules closer than gas

move rapidly but can slide over each other


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defined volume & shape

Incompressiblemolecules packed closely in

definite arrangement/rigid

shape


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Matter with fixed composition and distinct

properties, E.g H2O , NaCl

(i)

- simplest form of matter

- cannot be decomposed into simplersubstances by chemical means i.e only

of element

- can exist as atoms or molecules


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114 elements identified

Each given a unique name organized in a PeriodicTable


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(ii)

- substance composed of atoms of two or moreelements in fixed proportions

- can be separated only by chemical means

- exist as molecules (H2O, CO2)

- properties are different from the

elemental properties


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Combination of two or more substances, in whicheach substance retains its own chemical identity.

:

components uniformly mixed

(one phase) e.g. air also called solutions (gaseous,

liquid, solid solutions)


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(ii) :

components are not distributed uniformly (morethan one phase)

e.g. sand & rocks

sugar & sand

Separating Mixtures (by physical means):

basic techniques: filtration, floatation, crystallization,distillation, extraction and chromatography.


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Properties of matter can be grouped into twocategories:

: measured and observedwithout changing the composition or identity of a

substance. e.g. color, odor, density, melting point,boiling point.

: describe how substancesreact or change to form different substances. e.g.hydrogen burning in oxygen.


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Properties of substance can be divided into twoadditional categories:

Do not depend on the amount of the samplepresent. e.g. temperature, melting point, density.

Depends on quantity present. e.g. mass, volume.


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Changes in matter can involve either

chemical or physical changes.

: substance changes physicalappearance but not composition. e.g. changes ofstate :

liquid gas solid liquid

: substance transform into achemically different substance i.e. identifychanges. e.g. decomposition of water.


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SI Units

1960 : All scientific units use Systme International

dUnits (SI Units).Seven base units :

Physical Quantity Name of Unit Abbreviation

Mass Kilogram Kg

Length Meter mTime Second s (sec)

Electric current Ampere ATemperature Kelvin K Luminous intensity Candela cd

Amount of substance Mole mol


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SI base unit of length : meter (m)

1 m = 1.0936 yards

Mass :A measure of the amount of material in anobject.

SI base unit of mass : kilogram (kg)

1 kg = 2.2 pounds


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Temperature is a measure of hotness or coldnessof an object

3 temperature scales are currently in use:

(i) OF (degrees Fahrenheit)

(ii) OC (degrees Celsius)

(iii) K (Kelvin)

Scientific studies commonly used

Celsius and Kelvin scales


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Based on properties of gases

0 K is the lowest temperature that can beattained theoretically (absolute zero)

0 K = -273.15C


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Temperaturescale

Properties of water at sea level

Freezing point Boiling point

Fahrenheit, F 32 212

Celcius, C 0 100

Kelvin, K 273.15 373.15


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K = 0C + 273.15

C = K - 273.15

( )

( )329

5

32

5

9

=

+=

FC

CF


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SI unit of volume = (unit of length)3 = m3

Generally, chemists work with much smaller

volumes:cm3 , mL or cc

1 cm3 = 1 mL = 1 10 -6 m3

1000 cm3

= 1 L*Note: liter (L) is not an SI unit

1 dm 3 = 1 10 -3 m3


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Widely used to characterize substances.

Defined as mass divided by volume, d = mass (m)

volume (V)

Unit : g/cm3

Varies with temperature because volume changes

with temperature.

Can be used as a conversion factor to change massto volume and vice versa.

Common units :

g/mL for liquid, g/cm3 for solid, g/L for

gas.


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i. Determine the number of significant figures ina measured quantity.

ii. Express the result of a calculation with theproper number of significant figures.


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(i) - those that have defined

values or integers resulting from countingnumbers of objects. e.g. exactly 1000g in akilogram, exactly 2.54 in an inch.

(ii) - those that obtained

from measurements and require judgement.Uncertainties exist in their values.

Note : Uncertainties always exist in measuredquantities.


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- how well measured quantities agree

with each other.

- how well measured quantities agree

with the true value.


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Good precision

Good accuracy

Good precisionPoor accuracy

Poor precision

Good accuracy

Poor precisionPoor accuracy


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The standard deviation,s is a precision estimatebased on the area score where:

xi - i-th measurement

is the average measurementN is the number of measurements

N

xx

si

i =

2)(


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Measured quantities (inexact) are generallyreported in such a way that the last digit is thefirst uncertain digit. (2.2405g)

All certain digits and the first uncertain digit arereferred to as significant figures.

(i) Non-zero numbers are always significant

e.g. 2.86 : has three significant figures.


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(ii) Zeros between non-zero numbers are alwayssignificant. E.g. 1205 has four significant figures.

(iii) Zeros before the first non-zero digit are notsignificant. E.g. 0.003 : has one significant figure.

(iv) Zeros at the end of a number after a decimalplace are significant.. E.g. 0.0020 : has twosignificant figures.


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(v) Zeros at the end of a number before a decimalplace are ambiguous.

100: has one significant number unlessotherwise stated. If it is determined from countingobjects, it has three significant figures.

Method- Scientific notation removes the ambiguity of

knowing how many significant figures a number

possesses.


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Example:

(i) 225, 2.25

102

:

(ii) 10.004, 1.0004 104 :

(iii) 0.0025, 2.5 10-3 :

(iv) 0.002500, 2.500 10-3 : (v) 14 100.0, 1.41000 x 104 :

(vi) 14100, 1.4100 104, 1.41 104, 1.410 104 :could have , or . - need knowledge.


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1.5.3.1

Result must be reported to the least number of

.E.g. 20.4 g - 3.322 g = 17.1 g

Other Examples:The final answer should have thesame uncertainty, with the greatest uncertainty.

(i) 325.24 (uncertainty = 0.01)

21.4 (uncertainty = 0.1)

+ 145 (uncertainty = 1)

491.64 Answer :


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Cont:1.5.3.1 Addition (+)and Subtraction (-)

Other Examples:

(ii) 12.25 + 1.32 + 1.2 = 14.77 1.2 has the greatest uncertainty ( 0.1) the answer must be rounded to onedigit to the right of the decimal point.

Answer : 14.8

(iii) 13.7325 - 14.21 = -0.4775, Answer:

-0.48


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1.5.3.2 Multiplication ( ) andDivision ( )

Result must be to the least number ofsignificant figures.

E.g. 6.221 cm 5.2 cm = 32 cm2

To round off the final calculated answer sothat it has the same number of significantfigures as the least certain number.

Other Example:

(i) 1.256 2.42 = 3.03952

The least certain/precise number is 2.42 3

significant figures(s.f.). The answer must be

rounded to the 3 s.f.: 3.04


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Cont: 1.5.3.2 Multiplication( ) and Division ( )

Other Examples:

(ii) 16.231 2.20750 = 7.352661The least precise number is 16.231 (5 s.f.).

Answer is 5 s.f. : 7.3527

(iii) (1.1)(2.62)(13.5278) 2.650 =14.712121

The least precise number is 1.1 (2 s.f.).

Answer must be rounded to 2 s.f. : 15


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1.5.3.3 Rules for RoundingOff Numbers

(i) When the figures immediately following thelast digit to be retained is less than 5, the lastdigit unchanged.

e.g. 6.4362 to be rounded off to four significantfigures : 6.436

(ii) When the figure immediately following the last

digit to be retained is greater than 5, increase thelast retained figure by 1.

e.g 6.4366 to be rounded off to four significantfigures : 6.437


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Cont: 1.5.3.3 Rules forRounding Off Numbers

(iii) When the figure immediately following the lastdigit to be retained is 5, the last figure to beretained is increased by 1, whether it is odd or

even. e.g. 2.145 becomes 2.15 if three significant figures

are to be retained.

(iv) When a calculation involves an intermediateanswer, retain at least one additional digit past thenumber of significant figures.


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To be able to convert differentmeasurement units by using dimensional analysis.

is the algebraic process of

changing from one system of units to another. Conversion factors are used.

A conversion factor is a fraction whose numerator

and denominator are the same quantity expressedin different units.

Given units are being multiplied and divided togive the desired units.


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Desired unit = given unit conversion factor

conversion factor

In dimensional analysis, always ask threequestions:

(i) What data are given?

(ii) What quantity do we need?

(iii) What conversion factors are available totake us from what are given to what we need?

)unitgiven(

)unitdesired(


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Quantity 1 in. = 2.54 cm yields twoconversion factors

2.54 cm and 1 in. 1 in. 2.54 cm

Convert 5.08 cm to in. and 4.00 in. to cm

5.08 cm 1 in. = 2.00 in.

2.54 cm

4.00 in. 2.54 cm = 10.2 cm

1 in.


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Convert 6.23 ft3 to the appropriate SI unit.

ft3 to m3 and 3.272 ft = 1m

(1 ft )3 = (1m)3

(3.272ft)3

6.23 ft3 = 6.23 ft3 (1m)3 =0.178 m3

(3.272ft)3


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A persons average daily intake of glucoseis 0.0833 pound. What is this mass inmilligrams?

( 1 lb = 453.6 g)

lb1

g6.453

Answer: 3.78 x 10

-4

mg

lb g mg

0.0833 lb x x =g1

mg1000


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Chapter 1 Jan12