asc 0301 01 chemistry matter and measurement 54 slides 1 2

Inorganic Chemistry ACS 0301

Dr. Mahnaz M. AbdiE- mail: [email protected] First semester 2012Lecture1: Chemistry and Matter1

COURSE TITLE: CHEMISTRY COURSE CODE: CREDIT: LECTURE: LAB: TUTORIAL:

INORGANIC ASC 0301 2+1 2 HOURS/WEEKS 3 HOURS/2 WEEKS 2 HOURS/2 WEEKS

ASSESSMENT TEST 1: WEEK 5 15% TEST 2: WEEK 10 20% LAB REPORT: 15% PRESENTATION (TUTORIAL): 10% FINAL: 40% REFERENCE BOOK: Kenneth W.Whitten, Raymond Davis, Larry Peck, George G.Stanley, General chemistry, 7th. Ed., THOMSON, Brooks/ColeLecture1: Chemistry and Matter2

MINGGU PENGAJARAN

ASC 0301 TOPIK Chemistry and Matter Unit and Measurement The Structure of Atoms Atomic Mass Mole Concept Concentrations of Solutions

MINGGU PENGAJARAN 8

ASC 0301 TOPIK Periodic Table

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Periodicity Chemical bonding Ionic Bond Covalent Bond Polar Bond VSPER Theory Valence Bond Theory Intermolecular Solubility Forces and

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3

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Chemical ReactionsStoichiometry of Reactions Electromagnetic radiation Atomic Model Introduction to Quantum Theory Electronic configuration

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Important Elements in Firtilizer Important elements in Industry (PRESENTATION)3

Lecture1: Chemistry and Matter

Chemistry and Matter

Chemistry Classification of Matter Atoms Molecules Elements, substances and compounds Mixtures States and Properties of Matter Changes of matter4


Learning OutcomesClassify samples of matter as pure substances, homogeneous mixtures, heterogeneous mixtures, compounds, and elements. Recognize various form of matters. Relate names to formulas and

charges of simple ions.Combine simple ions to write formulas and names of some ionic compounds.5

ChemistryWhat is Science ? SCIENCE is the qualitative and quantitative study of nature and the natural laws What is Chemistry ? CHEMISTRY is a science that examines the composition, properties and changes of matter. What is Matter ? Anything that have mass and occupy space, and can generally be perceived by our senses (color, texture, odor, hardness, and taste)

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Disciplines of Chemistry Major traditional chemistry disciplines :

ChemistryInorganic Physical Organic Analytical

Study ofNon-carbon-containing substance (ASC 0301) Structure and changes of matter (ASC 0302) Substance containing carbon (ASC 0303) Composition of matter

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Disciplines of ChemistryA vast area of study & research; can be divided into traditional, modern, applied chemistry etc


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Higher TLecture1: Chemistry and Matter Lecture1: Chemistry and Matter

Lower T9

Properties of Matter Matter exhibits two types of properties : Physical property is displayed by

a sample of matter without undergoing any change in its composition It includes mass, color, temperature volume, density, melting point, etc. Chemical property is displayed by a

sample of matter as it undergoes a change in its composition. This includes flammability, toxicity, reactivity, acidity, corrosiveness, etc.

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Changes of Matter In a physical change, there is no change in composition; no

new substances are formed (without changes in chemical identity). Examples: evaporation, melting, solubility, crystallization In a chemical change or chemical

reaction, matter undergoes a change in composition; new substance are formed. Examples: corrosion of metals, polymerization of alkene, fermentation of food.

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Types of MatterMatter (may be solid, liquid, Or gas): anything that occupies space and has mass

Heterogeneous matter: Non uniform composition

Physically separable into

Homogeneous matter: uniform composition throughout

Substances: fixed composition; cannot be further purified

Physically separable into

Solutions: homogeneous mixtures; uniform compositions that may vary widely

Compounds: elements united in fixed ratios

Chemically separable into

Elements: cannot be subdivided by chemical or physical changes12

Combine chemically Lecture1: Chemistry and Matter to form

Element, Substance and Compound Substance is matter that has a definite composition and do

not vary from one sample to another. Substances are either elements or compounds Element is matter that cannot be broken down into simpler

substances by chemical reaction. Compound is made up of two or more elements, chemically

bonded in definite proportions and can be broken down into simpler substances. Terms Substance and Compound are often used

synonymouslyLecture1: Chemistry and Matter13

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MixtureA mixture can contain two or more different components in varying proportions. A homogeneous mixture has the same composition

throughout a sample of matter. components are indistinguishable. Examples: alloy, air, salt water. A heterogeneous mixture varies in composition from one part of the mixture to another. Components are distinguishable. Examples: Mixture of salt and charcoal.15

Homogeneous

Heterogeneous


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PURE SUBSTANCE In chemistry, a pure substance has characteristic physical

and chemical (phys/chem) properties. Cannot be broken down by physical means. Pure substances are elements or compounds.


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COMPOUND A chemical substance made up of two or more

elements in fixed ratio. More than 20 million compounds are known. Cpds have characteristic phys/chem properties. Cannot be resolved into its constituent elements via physical means.


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ELEMENT Fundamental type of matter that cannot be chemically broken down. There are 116 known elements with 90 occurring

naturally; shown on the Periodic Table Building blocks for all substances. Classified as metals (75%), nonmetals (20%) or metalloids (5%). Elements have characteristic phys/chem properties


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AtomsThe building blocks of matter; the smallest unit of matter that participates in a chemical reaction.


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Molecules

A molecule is the smallest particle of an element that can have a stable independent existence. Usually have two or more atoms bonded together Examples of molecules: H2 O2 S8 NaCl H2O CH4

C2H5OH

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Dmitri Ivanovich Mendeleev 1869


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Noble gases Halogens

Alkaline earth metalsAlkali Metals

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Chemical FormulasChemical formula shows the chemical composition of the substance. Ratio of the elements present in

the molecule or compound.Monatomic elements: He, Au, Na

Diatomic elements: O2, H2, Cl2 More complex elements: O3, S4, P8 Compounds: H2O, NaCl, C12H22O11 Substance consists of two or more elementsLecture1: Chemistry and Matter24

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Ions Cations consist of one atomGroup 1Li+ , lithium ion Na+ , sodium ion

Group 2Be2+ , beryllium ion Mg2+ , magnesium ion

Group 3Al3+ , aluminium ion

K+ , potassium ionRb+ , rubidium ion Cs+ , cesium ion

Ca2+ , calcium ionSr2+ , strontium ion Ba2+ , barium ion


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Transition Metals and Post-Transition metalsCr2+, chromium(II) ion Cr3+ , chromium(III) ion Mn2+, manganese (II) ion Mn3+ , manganese (III) ion Fe2+, iron(II) ion Fe3+, iron(III) ion Cu+, copper(I) ion Cu2+ , copper(II) ion Sn2+, tin(II) ion Sn4+ , tin(IV) ion Pb2+, lead(II) ion Pb4+ , lead(IV) ion

Anions consist of one atomH-, hydrideC-, carbide

N3-, nitrideP3-, phosphide

O2-, oxideS2-, sulfide

F-, fluorideCl-, chloride

Si4-, silicide

As3-, arsenide

Se2-, selenide

Br-, bromide26


Polyatomic ions contain more than one atom.NH4+, ammonium H3O+, hydronium OH-, hydroxide CN-, cyanide NO2-, nitrite CO32-, carbonate HCO32-, hydrogen carbonate SO32-, sulfite HSO32-, hydrogen sulfite SO42-, sulfate

NO3-, nitrateClO-, hypochlorite ClO2-, chlorite ClO3-, chlorate ClO4-, perchlorate

CrO42-, chromateCr2 O72-, dichromate PO43-, phosphate HPO42-, mono hydrogen phosphate H2PO42-, dihydrogen phosphate


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Ionic CompoundsPotassium chloride: KCl = K+ + ClCalcium sulphate: CaSO4 = Ca2+ + SO42Sodium carbonate: Na2CO3 = 2Na+ + CO32Barium nitrate: Ba(NO3)2 = Ba2+ + 2NO3Ammonium phosphate: (NH4)3PO4 = 3NH4+ + PO43Chromium(III) chloride: Cr(Cl)3 = Cr3+ + 3ClIron(II) sulphite: FeSO3 = Fe2+ + SO32-


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Unit and Measurement

Imperial System Metric System System International (SI) Accuracy and Precision Significant Figures


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Modern Chemistry As of 1850s onwards, chemists used weighing balance more

systematically to measure mass (a quantity of matter) in their research Since then, better equipments and devices were progressively

invented and used; then chemistry involves doing all sorts of measurements (simple and complex; visual and aided) So, what is it that chemists measure? All sort of properties of

matter - distance, vol., temp., time, energy, area, density, velocity, frequency, luminescence, light absorbance ... etcLecture1: Chemistry and Matter30

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Scientific measurement and unitAll measurements contains 2 essential pieces of information:1) a number (the quantitative piece) 2) a unit (the qualitative piece)

Example:The number 50 is somewhat meaningless without Units. Consider this: Daily wage = RM50/day wt of an egg = 50g/egg size of a square paper = 50cm x 50 cmLecture1: Chemistry and Matter32

Properties of Matter Intensive A quality of a pure substance which is independent of quantity color, melting point, hardness Extensive A quality of a pure substance which is dependent on quantity mass, volume

What are other intensive or extensive properties?Lecture1: Chemistry and Matter 33

Scale, Unit and Standard Results of a measurement are always expressed in some kind of a

scale that is defined in terms of a particular kind of unit

Example : Scale of length British Imperial system French Colonial system Inch = base unit Millimetre = base unit 12 inch = 1 foot*** 10 mm = 1 cm 3 feet = 1 yard 100 cm = 1 metre*** 1760 yards = 1 mile 1000 m = 1 km For comparison and communication purposes,

the scale must be defined by a standard (a reference value for comparison) The scales of length are related & interconvertible by

a conversion factor 2.54 cm = 1 inch 1 cm = 0.3937 inch

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Imperial System of Unit Still being used in Britain and USA

Weight ounce, pound, ton 16 oz = 1 lb; 14 lb = 1 stone; 2240 lb = 1 tonLength inch, foot, yard, mile 12 inches = 1 foot; 3 feet = 1 yd; 5280 feet = 1760 yds = 1 mile Volume cup, pint, quart, gallon 2 cups = 1 pint; 2 pints = 1 quart; 4 quart = 1 gal Not employed in scientific work since 1960s

Confusing and have to memorize many conversion tables

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Metric System of Measurement Basic/fundamental units/types of measurements are :

Type Mass Length Volume Time Energy

Name gram metre litre second joule

Symbol g m L s J

Simple and easy system of measurement Units of measurement can be multiplied or divided by

a factor 10 Can use prefixes to change size of unit36

Metric System Prefixes Base units in the Metric System can be converted into units that

are more appropriate for the quantity being measured by adding a prefix to the name of the base unit. Common metric prefixes are:Prefix Symbol femtof picop nanon micro millim centic decid kilok megaM gigaG teraT Meaning x 1/1,000,000,000,000,000 (10-15) x 1/1,000,000,000,000 (10-12) x 1/1,000,000,000 (10-9) x 1/1,000,000 (10-6) x 1/1,000 (10-3) x 1/100 (10-2) x l/10 (10-1) x 1,000 (103) x 1,000,000 (106) x 1,000,000,000 (109) x 1,000,000,000,000 (1012)


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Systeme Internationale (SI Units of Measure)Before 1960, Science uses the Metric System of Units After 1960, the International System of Units (SI = Le Systeme Internationale) was proposed.A modern version of the Metric System of Units. It is used globally in science, industry and commerce as a standard/reference measurement units. The seven base/standard units of the SI system are:

length mass time Electric current Temperature Amount of substance Luminous intensity

meter kilogram second s ampere kelvin mole candela

m kg s A K mol cd38

Units outside SI system honorary unitsSI has 7 base/fundamental units 3 more are included as honorary members

Base unit 1.Litre 2. Metric tonne 3. Unified atomic mass unit

SymbolL t amu (u)

Dimension1 L = 1 dm3 = 103 ml 1 t = 103 kg 1 u = 1.66054 X 10-27 kg

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Derived SI units [dimension] Other units of measurement can be derived by combining one or more of

the 10 SI base units. Some of the common non-SI or derived units used in chemistry are : Physical Quantity Name of derived unit density electric charge coulomb electric potential volt energy joule force newton frequency hertz pressure pascal velocity meters/second volume cubic meter Symbol kg/m3 C (A.s) V (J/C) J (kg-m2/s2) N (kg-m/s2) Hz (s-1) Pa (N/m2) m/s m3

** Symbol of a derived unit is termed the DIMENSION/UNIT of the physical quantityLecture1: Chemistry and Matter40

Exact numbersExact numbers = numbers that are known/exact by definition They are called exact values because they are measured in complete units and are not divided into smaller parts. Examples of exact values: 17 people; 28 cars Conversion factors have an infinite (never ending) number of significant figures.Examples: 60 seconds in one minute 1000 meters in one kilometer 7 days in one week 1 inch = 25.4 mm 1 yard = 36 inchesLecture1: Chemistry and Matter41

Measured numbers Measured numbers = an actual estimated reading:

(a) as shown by meters/detectors of instruments (b) by visual account from calibrated/graduated equipment (burette, pipette) (c) from mathematical calculations (average, standard deviation) Measurements by these methods have inherent uncertainty

or experimental errors


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Measurement Chemistry is an experimental science; it involves measurements of

properties of matter Measuring properties of matter involves 3 parameters:

(a) magnitude/quantity (b) base unit or dimension (c) uncertainty/error Example: weighing a sample of salt

(a) average magnitude = 5.7076 (b) base unit = gram (c) uncertainty (using 4 different balances) 5.6785 g; 5.702 g; 5.65 g; 5.8 g So, what is the true/real weight of the salt sample?Lecture1: Chemistry and Matter43

Types of experimental errors Systematic

Instruments not calibrated properly Reagents are not correctly and properly prepared Imperfections/limitations of instrument and glassware can be improved by proper calibration, adherence to strict preparation procedures, use more sensitive and reliable instruments, run control/blanks ... etc Random

Tests done at different times, different room T ... etc Personal mistakes and carelessness Can be improved by tight control of environmental factors, replication of tests, applying appropriate statistical analyses on data obtained ...etc44

Uncertainty in measurement Measurement of a property of matter involves 2 parameters(A) (B) PRECISION ACCURACY

(A) PRECISION = degree of reproducibility i.e.

(a) how close are the measurements to each other (b) how well they agree with each other (B) ACCURACY = how close measurements are to the

true/accepted value AIM = experimental measurements must be as precise

and accurate as possible; i.e. least possible errorsLecture1: Chemistry and Matter45

Accuracy vs PrecisionAccuracy refers to the proximity ofa measurement to the true (accepted) value of a quantity.

Precision refers to the proximityof several measurements to each other.


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Accuracy vs PrecisionIn the real world, we never know whether the

measurement we make is accurate. Why not?We make repeated measurements; calculate

the mean, x and the standard deviation, s,We hope (but not always correctly) that good

precision implies good accuracy.


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Uncertainty in measurementsDifferent measuring devices have different uses and different degrees of accuracy.

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Uncertainty in measurementsThe smallest division of this graduated cylinder is 1 mL. Therefore, our reading error will be 0.1 mL or 1/10 of the smallest division. An appropriate reading of the volume is 36.5 0.1 mL. An equally precise value would be 36.6 mL or 36.4 mL. The smallest division in this buret is 0.1 mL. Therefore, our reading error is 0.01 mL. A good volume reading is 20.38 0.01 mL. An equally precise answer would be 20.39 mL or 20.37 mL.49

Significant figures Topics Determining the number of significant figures Scientific notation

Addition & Substraction Multiplication & Division Rounding off


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Significant figures When we measure quantities, esp. for the very first time, it is

difficult to know if measurement is accurate. To know if a measurement is accurate, you need to know its

true value. If you knew its true value, then, you dont need to do the measurement! But, it is easy to know if your measurements are precise.

One indicator of the preciseness of a measurement is the number of significant figures in a measured number. Significant figures relate to certainty of the measurement.

As the number of significant figures increases, the more certain is the measurement.Lecture1: Chemistry and Matter51

Number of significant figures Significant figures = number of digits that expresses the result to

the true measured precision Example: Consider a number

92.154

(5 sig. Figures)

Start from the left and count the digits. There is

5 significant figures in this number DO NOT worry about the decimal point !!


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Zeroes in significant figures Zeroes are special digits Sometimes they count as a sig. digit,

sometimes they dont. It depends on their position. Example: 0.092060 has 5 significant figures First two zeroes are not significant, they

are just position holders So, when counting significant digits, always

start from the left and don't start counting until you meet the first non-zero digit. Then everything after that counts, including the zeroes !!

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Rules for counting significant figures are summarized below Leading zeros are not significant. 0.421 has three significant figures. Zeros within a number are always significant. Both 4012 and 40.05 contain four significant figures. Trailing zeros in the whole number are not significant. Thus, 470,000 has two significant figures. Trailing zeros that aren't needed to hold the decimal point

are significant. For example, 114.20 has five significant figures.54

Scientific notation A good way to avoid problems of significant and non-significant

zeroes is to use scientific notation Example

920000 = 9.2 x 105 0.092067 = 9.2067 X 10-2 0.092 = 9.2 X 10-2 0.0920 = 9.20 X 10-2

Has 2 sig. figs. Has 5 sig. figs. Has 2 sig. figs. Has 3 sig. figs.

When using scientific notation, all digits including zeroes are

significant

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Rounding offChemistry is an inexact science All physical measurements have some error Thus, there is some inexactness in the last digit of any number Use what ever round-off procedures you so choose Reasonably close answers are acceptableLecture1: Chemistry and Matter56

Rounding significant figures A salt sample is weighed 3 times:

12.4330 g

12.4334 g 12.4335 g

6 sig. figures

Calculate the average:

(12.4337 + 12.4334 + 12.4335) / 3 = 12.43353333 (calculator display) Makes no sense to write down all the numbers; average value

cannot be more precise than that of the measurements. Need to round off the number to the correct number of sig. figures. In this case Insignificant digits 3333 are dropped 2.43353333 is rounded off to 6 sig. figures; i.e.12.4335Lecture1: Chemistry and Matter57

Rounding off numbers When insignificant digits are dropped from a number, the last digit should be rounded for best accuracy Example :

To determine how the last digit should be rounded, convert

the digits to be dropped to a decimal fraction ( .333333); then follow the rules for rounding numbers


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Rules for rounding numbers to the correct number of significant figures1.

If the decimal fraction is greater than , add one to the last digit retained If the decimal fsraction is less than , do not change the last retained digit If the decimal fraction is exactly , add one to the last digit retained if it is odd59

2.

3.


Examples Round off 9.473, 9.437, 9.450 & 9.750 to 2 sig. figures.For 9.473 the last digit retained is 4; decimal fraction is .73. So Rule #1 applies; 9.473 is rounded to 9.5 For 9.437 the last digit retained is 4; the decimal fraction is .37. So Rule #2 applies; 9.437 rounded to 9.4 For 9.450 the last digit retained is 4; the decimal fraction is .50. So Rule #3 applies; 450 is rounded to 9.4 For 9.750 the last digit retained is 7; the decimal fraction is .50. So Rule #3 applies; 9.750 is rounded to 9.8Lecture1: Chemistry and Matter

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asc 0301 01 chemistry matter and measurement 54 slides 1 2

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