Lecture Presentation Chapter 1

Introduction to Matter and Measurement

Dr. Rajani Srinivasan

Tarleton State University

The Text book

Our journey together in CHEM 105

Total chapters covered 12.5 ?????

Chapters 1-------12

Chapter 13 up till section 13.4

Present Chapter Contents

• Matter• Classification of Matter• Properties of Matter• Units of Measurements• Uncertainty in Measurement• Dimensional analysis

Chemistry

• Study of matter, its properties and behavior• Matter is the central to everything so

chemistry is the Central Science

Chemistry

Health Sciences

AgricultureGeology

Engineering

Matter

“Anything that has mass and occupies space”

Example : anything and everything around you is matter What is the basic building block of Matter ?

ATOMS

Matter

When two or more atoms combined together they are called

We will study about these later……………..Later

molecules

Classification of Matter

Matter

States of Matter Composition

Solid Liquid

gas

Element Compound

Mixture

States of Matter

• Solid • Liquid • Gas

States of Matter

SOLIDS LIQUIDS GAS

Have definite shapeDefinite volumeNon compressible

No definite shapeDefinite volumeNon compressible

No definite shapeNo Definite volume compressible

States of Matter

SUBSTANCES

Pure Impure

Element Compound Mixture

Element

Simple definition:Substances that cannot be further decomposed into simpler substances

Molecular levelSubstances made up of one kind of atom.

Example : Oxygen, Hydrogen, Iron etc.

Important facts about Elements

• A total of 118 elements are known• 90% of earths Crust is made up of only five

elements ( Oxygen, Silicon, Aluminum, Iron and Calcium)

• 90% of Human body is made up of (Oxygen , hydrogen and Carbon.

• Please refer to table 1.2 in the book which includes the names and symbols of common element.

Compounds

• Pure substances contains two or more substances

Molecular Level • Made up of two or more types of atoms

Example: Water (H2O), HCl etc.

Important facts

• In a compound each atom is present in a definite composition

• Example: irrespective of the source in water 11% of Hydrogen combines with 89% of Oxygen

“ This is called Law of constant composition or Law of definite proportion”

Mixtures

• Made of two or more substances but each substance retains its identity

• Example AIR – It’s a mixture of all the gasses but all the gasses retain their identity.

• Each substance that forms a mixture is called

“Component” of the mixture.

Types of Mixture

Mixture

Homogenous Heterogeneous

Air , salt , sugar etc. Rocks and woods

Having same composition, properties and appearance

Having different composition, properties and appearance

Atomic Level

Properties of Matter

• Physical• Chemical • Intensive • Extensive

Physical

• Can be observed without Changing the composition and identity of the substance

Example: color. Odor , melting point , boiling point

Evaporation of water

Such changes are called Physical Changes

Chemical Properties

• These properties describes the way substances completely change or react.

Example : Burning

“Chemical Change is a change in which substances changes into a completely new substances.”

Example : decomposition of water into Hydrogen and Oxygen , formation of water

Chemical change

Intensive Properties

• Properties of Matter which does not depend upon the amount of the examined substance.

• Can be used to identify a substance

Example: Color, odor, melting point etc.

Extensive Properties

• Properties which depend upon the amount of the substance examined

Example: Mass, length, volume etc.

Classification of Matter

Separation of Mixtures

• Based on the type of Mixture different methods are used for their separation

• Separation is based on the individual properties of the components in the mixture.

Gravity separation (filtration)

Magnetic Separation distillation

Chromatography

Heterogeneous mixture Homogenous mixture

filtration

Principle: difference in the states of matterUsed for : Heterogenous mixtures

Distillation

Used for separation homogenous mixture Principle : difference in boiling points between the components

ChromatographyHomogenous mixture Principle : Difference in the solubility in the solvents

Units of measurement

• Every quantitative measurement is represented by a number and a “Unit” depending upon the substance.

• Most commonly used System of measurements “ Metric” developed on France in the 18th century.

• Other type of system used is “English system” commonly used in United states .

SI Units

• “System International” Units derived from Système International d’Unités

• In 1960 an international agreement was reached to use same type of measurement system.

• A total of seven base units are used .

SI units

Prefixes used with the Metric systems

SI units

• Length – meter- English unit = yard• Mass – kilogram English unit = Pound ( 1Kg=

2.2lb)• Temperature – SI unit is Kelvin English unit =

Fahrenheit

Temperature

• Degree of hotness or coldness is called temperature

• Always flows from temperature tohigher

lower

Temperature scales

Celsius or centigrade Kelvin Fahrenheit

Based on boiling (100 degree centigrade) and the freezing point of water (0 degree centigrade)

• It’s the SI unit • Zero kelvin is the

lowest attainable temperature =

- 273.15◦ C• Also called “Absolute

Zero”

• Commonly used in The USA

• Boiling point = 212◦ F• Freezing point = 32◦ F

K = ◦ C + 273.15 ◦ C = 5/9(◦ F-32) or ◦ F = 9/5 ◦ C + 32

Celsius or centigrade Fahrenheit

Comparison

Derived units

• Those units which are derived from basic SI units or made of more than one basic SI units

1) Speed- – Ratio of distance travelled per unit time – SI units used distance= meter; time= seconds– Represented by

2) Volume- Length Cubed (L*L*L)

-SI unit used – m3, cm3,1L, etc.

ms

Derived units

3) Density

– mass of a substance per unit volume

- SI units used mass= g and Volume = cm3 or 1ml

- Represented byg

cm3

Uncertainty in measurement

Exact Numbers Inexact numbers

Numbers which can easily be counted Example = 12 dozens , 1000g in 1Kg

Numbers obtained by measurements usually include several errors like equipment errors, manual errors . Example: measure the length of a given object using ruler by many students

Precision accuracy

Uncertainty in measurements

Uncertainty in measurement

• Precision – Measure of how closely the individual measurements agree.

Average of several measured values gives us precision • Accuracy – how closely individual

measurement agree with the true value

Properly calibrated instrument gives us the accurate measurement

Significant figures

• All the digits of the measured quantity including the uncertain ones are called “Significant figures”.

• Usually there is a uncertainty in the last digit reported for any measured quantity.

Example: No. of significant figures in the reported measured mass

2.2g = 2

0.4g = 1

1.04g = 3

Rules 1. All nonzero digits are significant

613 has three sig figs

123456 has six sig figs

2. Zeroes between two significant figures are themselves significant.

5004 has four sig figs

602 has three sig figs

6000000000000002 has 16 sig figs!

Rules 3. Zeroes at the end of a number are significant if a decimal point is written in the number• 5.640 has four sig figs• 120000. has six sig figs• 120000 has two sig figs – unless you’re given additional information

in the problem

4.Zeroes at the beginning of a number are never significant.

• 0.000456 has three sig figs• 0.052 has two sig figs• 0.000000000000000000000000000000000052 also has two sig figs!• 2.30 x 10¯5 = 3 sigfig• 4.500 x 1012 = 4 sigfig

Rules

Exact numbers

Exact numbers, such as the number of people in a room, have an infinite number of significant figures.

There are exactly 12 inches in one foot. Therefore, if a number is exact, it DOES NOT affect the accuracy of a calculation nor the precision of the expression.

Example = 100 years in a century

Scientific notation

• Depending upon the scenario or type of experiment significant digits vary .

• To express the number of significant digits we use scientific notation

Example : 200 based on the question

Suppose

we need 2 sig figs = 2.0 * 10 2

We need 3 sig figs = 2.00 * 10 2

Rules

5) When addition or subtraction is performed, answers are rounded to the least significant decimal place.7.939 + 6.26 + 11.1 = 25.299 (this is what your calculator shows) your final answer is limited to one sig fig to the right of the decimal or 25.3 (rounded up).

Rules

6)When multiplication or division is performed, answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation

(27.2 x 15.63) 1.846 = 230.3011918 (this is what you calculator shows)

your final answer is limited to three sig figs, the answer is 230. (rounded down)

Dimensional analysis

• Used to convert one unit to other• Arrive at a proper unit at the end of the

problem• Correct use of conversion factors• Provides systematic way to solve the problem

and detect errors

Conversion factor

• It is a fraction whose numerator and denominator are same quantities expressed in different units (e.g., 1 in. = 2.54 cm)

1 in.

2.54 cm

2.54 cm

1 in.or

Formula

Given unit desired unitdesired unit

given unit

Conversion factor

Example

• For example, to convert 8.00 m to inches,– convert m to cm– convert cm to in.

8.00 m 100 cm

1 m

1 in.

2.54 cm 315 in.

Conversion using Volume

• density of gold = 19.3 g/cm3

• Volume 2 in3

Find = mass in grams of gold conversion factor required = in3 to cm3

We know 1in = 2.54 cm Therefore 13in3= (2.54)3 cm3 = 16.39 cm3

Contd…

• Conversion factor used • 1in3/16.39 cm3 or 16.39cm3/1in3

2 in3

1in3

19.3g. 633 g

16.39cm3

1cm3

in3 cm3 g