Chapter 1 July08

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CHAPTER 1CHAPTER 1

INTRODUCTION TO MATTER AND

MEASUREMENT

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CONTENTSCONTENTS

1.1 Introduction1.2 Classification of Matter1.3 Properties of Matter1.4 Units of Measurement1.5 Uncertainty in Measurement1.6 Dimensional Analysis

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Learning outcomes:

Able to differentiate between the three states of matter.

Able to distinguish between elements, compounds and mixtures.

Able to distinguish between physical and chemical properties.

Able to use and convert different units of measurement.

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1.1 Introduction

Chemistry is the study of properties of materials and changes they undergo.

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 and agriculture.

Able to contribute to problem solving analysis.

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1.2 Classification of Matter

MatterPhysical material - anything that has mass

and occupies space.

Classifications of MattersMatter 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

Pure substance

Mixture

Element

Compound

Homogeneous

Heterogeneous

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1.2.1 State of Matter (Physical State)

Gasno fixed volume/shapeeasy to compress/expandmolecules are far apartmove at high speedoften collide

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Con’t: 1.2.1 State of Matter (Physical State)

Liquidvolume independent of containerslightly compressiblemolecules closer than gasmove rapidly but can slide over each other

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Con’t: 1.2.1 State of Matter (Physical State)

Soliddefined volume & shapeIncompressiblemolecules packed closely in definite arrangement/rigid shape

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1.2.2 Composition

Pure SubstanceMatter with fixed composition and distinct

properties, E.g H2O , NaCl

(i) Elements - simplest form of matter- cannot be decomposed into simpler substances by chemical means i.e only one kind of element- can exist as atoms or molecules

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Con’t: 1.2.2 Composition

114 elements identified Each given a unique name organized in a

Periodic Table

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Con’t: 1.2.2 Composition

(ii) Compounds - substance composed of atoms of two or more

elements 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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Con’t: 1.2.2 Composition

MixtureCombination of two or more substances, in

which each substance retains its own chemical identity.

(i) A Homogeneous mixture:– components uniformly mixed

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

liquid, solid solutions)

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Con’t: 1.2.2 Composition

(ii) A Heterogeneous mixture:– components are not distributed uniformly (more

than 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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1.3 Properties of Matter

Properties of matter can be grouped into two categories:

Physical properties : measured and observed without changing the composition or identity of a substance. e.g. color, odor, density, melting point, boiling point.

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

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Con’t: 1.3 Properties of Matter

Properties of substance can be divided into two additional categories:

Intensive propertiesDo not depend on the amount of the sample present. e.g. temperature, melting point, density.

Extensive propertiesDepends on quantity present. e.g. mass, volume.

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1.3.1 Physical and Chemical Changes

Changes in matter can involve either chemical or physical changes.

Physical change : substance changes physical appearance but not composition. e.g. changes of state : liquid gas solid liquid

Chemical change : substance transform into a chemically different substance i.e. identify changes. e.g. decomposition of water.

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1.4 Units of Measurement

SI Units1960 : All scientific units use Système International

d’Unités (SI Units).Seven base units :

Physical Quantity Name of Unit AbbreviationMass Kilogram KgLength Meter mTime Second s (sec)Electric current Ampere ATemperature Kelvin KLuminous intensity Candela cdAmount of substance Mole mol

1.4 Units of Measurement

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1.4.1 Length and Mass

SI base unit of length : meter (m)1 m = 1.0936 yards

Mass :A measure of the amount of material in an object.SI base unit of mass : kilogram (kg)

1 kg = 2.2 pounds

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1.4.2 Temperature

Temperature is a measure of hotness or coldness of an object

3 temperature scales are currently in use: (i) OF (degrees Fahrenheit) (ii) OC (degrees Celsius) (iii) K (Kelvin)

Scientific studies commonly usedCelsius and Kelvin scales

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Con’t: 1.4.2 Temperature

Kelvin (SI Unit)Based on properties of gases0 K is the lowest temperature that can be

attained theoretically (absolute zero)0 K = -273.15C

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Temperature scale

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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Con’t: 1.4.2 Temperature

Temperature conversions K = 0C + 273.15C = K - 273.15

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5

325

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FC

CF

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1.4.3 Volume

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 unit1 dm 3 = 1 10 -3 m3

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1.4.4 Density

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 mass to volume and vice versa.

Common units : g/mL for liquid, g/cm3 for solid, g/L for gas.

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1.5 Uncertainty in Measurement

Objectives

i. Determine the number of significant figures in a measured quantity.

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

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Con’t: 1.5 Uncertainty in Measurement

Two types of numbers: (i) Exact numbers - those that have

defined values or integers resulting from counting numbers of objects. e.g. exactly 1000g in a kilogram, exactly 2.54 in an inch.

(ii) Inexact numbers - those that obtained from measurements and require judgement. Uncertainties exist in their values.

Note : Uncertainties always exist in measured quantities.

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1.5.1 Precision and Accuracy

Precision - how well measured quantities

agree with each other.

Accuracy - how well measured quantities agree with the “true value”.

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Con’t: 1.5.1 Precision and Accuracy

Good precisionGood accuracy

Good precisionPoor accuracy

Poor precisionGood accuracy

Poor precisionPoor accuracy

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Con’t: 1.5.1 Precision and Accuracy

• The standard deviation,s is a precision estimate based on the area score where:

• xi - i-th measurement is the average measurementN is the number of measurements

N

xxs i

i

2)(

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1.5.2 Significant Figures

Measured quantities (inexact) are generally reported in such a way that the last digit is the first uncertain digit. (2.2405g)

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

Rules:(i) Non-zero numbers are always significant

e.g. 2.86 : has three significant figures.

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Con’t: 1.5.2 Significant Figures

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

(iii) Zeros before the first non-zero digit are not significant. E.g. 0.003 : has one significant figure.

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

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Con’t: 1.5.2 Significant Figures

(v) Zeros at the end of a number before a decimal place are ambiguous.E.g. 100: has one significant number unless otherwise stated. If it is determined from counting objects, it has three significant figures.

Method - Scientific notation removes the ambiguity of knowing how many significant figures a number possesses.

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Con’t: 1.5.2 Significant Figures

Example:

(i) 225, 2.25 102 : three significant figures (s.f.).

(ii) 10.004, 1.0004 104 : five s.f. (iii) 0.0025, 2.5 10-3 : two s.f. (iv) 0.002500, 2.500 10-3 : four s.f. (v) 14 100.0, 1.41000 x 104 : six s.f. (vi) 14100, 1.4100 104, 1.41 104, 1.410 104 :

could have three, four or five s.f. - need knowledge.

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1.5.3 Significant Figures in Calculations

1.5.3.1 Addition (+) and Subtraction (-) Result must be reported to the least number of

decimal places.E.g. 20.4 g - 3.322 g = 17.1 g Other Examples:The final answer should have the

same uncertainty, with the greatest uncertainty.(i) 325.24 (uncertainty = 0.01) 21.4 (uncertainty = 0.1) + 145 (uncertainty = 1) 491.64 Answer : 492

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Con’t: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 one digit 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 () and Division ()

Result must be to the least number of significant figures.

E.g. 6.221 cm 5.2 cm = 32 cm2

To round off the final calculated answer so that it has the same number of significant figures 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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Con’t: 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 = 14.712121The 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 Rounding Off Numbers

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

e.g. 6.4362 to be rounded off to four significant figures : 6.436

(ii) When the figure immediately following the last digit to be retained is greater than 5, increase the last retained figure by 1.

e.g 6.4366 to be rounded off to four significant figures : 6.437

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Con’t: 1.5.3.3 Rules for Rounding Off Numbers

(iii) When the figure immediately following the last digit to be retained is 5, the last figure to be retained 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 intermediate answer, retain at least one additional digit past the number of significant figures.

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1.6 Dimensional Analysis

Objective: To be able to convert different measurement units by using dimensional analysis.

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 expressed in different units.

Given units are being multiplied and divided to give the desired units.

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Con’t: 1.6 Dimensional Analysis

Desired unit = given unit conversion factor conversion factor

In dimensional analysis, always ask three questions:(i) What data are given?(ii) What quantity do we need?(iii) What conversion factors are available to

take us from what are given to what we need?

)unitgiven(

)unitdesired(

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Example 1

Quantity 1 in. = 2.54 cm yields two conversion 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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Example 2

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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Exercise 1.1

A person’s average daily intake of glucose is 0.0833 pound. What is this mass in milligrams?

( 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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END of CHAPTER 1

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