CHEMISTRY 1307General Chemistry I

Instructor: Mrs. Anna Mkrtchyan-Antonyan

Classroom: 219 Phone: EX. 209

E-Mail: amkrtchyan@agbumds.org

Text: Chemistry: Principles and Reactions Masterton • Hurley 5th Edition

AP CHEMISTRY

CHEMISTRY:CHEMISTRY:

Chemistry is the study of the Chemistry is the study of the properties, compositionproperties, composition,,and and structure structure of of mattermatter, the , the physicalphysical and and chemicalchemicalchangeschanges it undergoes, and the it undergoes, and the energyenergy liberated or liberated orabsorbed during those changes.absorbed during those changes.

Why Study Chemistry?

1. To better understand the world: what it is made of1. To better understand the world: what it is made ofand how it works.and how it works.

2. Because it is the most practical and relevant of the2. Because it is the most practical and relevant of thesciences - chemistry is the study of sciences - chemistry is the study of EVERYTHING!EVERYTHING!

3. It is the 3. It is the “Central Science”“Central Science” - All other sciences - All other sciencesintersect at and depend on chemistry.intersect at and depend on chemistry.

4. It is essential to the national 4. It is essential to the national and localand local economies. economies.

Chapter 1: Matter and Measurements

Contents

Physical properties and states of matter Système International Units Uncertainty and significant figures Dimensional analysis

Properties of Matter

Matter: Occupies space, has mass and inertiaExamples:Examples: chairs chairs gasoline gasoline clothes clothes batteriesbatteries people people the earth the earth paint paint paper paper oxygenoxygen water water salt salt aluminumaluminum air air rocksrocks

Composition: Parts or componentsand their relative proportionsex. H2O, 11.9% H and 88.81% O H2O2, 5.93% H and 94.07% O

Properties: Distinguishing features physical and chemical properties

Matter and Change

• Physical Property- A property displayed without a change in composition

E.g. color.

• Chemical Property – Ability or inability of a sample of matter to undergo a change in composition under stated conditions.

Matter and Change

• Physical Change - A change in which each substance involved in the change retains its original identity and no new elements or compounds are formed.

H2O (s) H2O (l)

MeltingMelting“ice”

Matter and Change

• Chemical Change (Chemical Reaction)-

• A change that involves a change in composition. – One or more kinds of matter are converted to

new kinds of matter with different compositions.

• 2 H2 (g) + O2 (g) 2 H2O (l)

• AgNO3 (aq) + HCl (aq) AgCl (s) + HNO3 (aq)

““Reacting”Reacting”

Classification of Matter

• AtomAtom: Matter is made up of very tiny units called atoms. There are 115 different atoms. (The basis of elements)

• ElementElement: A substance made up of only a single type of atom. There are 115 elements ( 90 of them are from natural sources)

– A substance that cannot be broken down (decomposed) into simpler substances by chemical reactions.

• CompoundCompound: A combination of two or more different elements.

– A substance composed of two or more elements chemically combined in fixed ratios by mass.

Water - H2O Carbon dioxide - CO2

Sodium Chloride - NaCl Iron(II) sulfide - FeS

• MoleculeMolecule: The smallest entity having the same elemental combination as the compound.

A molecule of water consists of three atoms: two hydrogen atoms joined to a single oxygen atom. A molecule of hydrogen peroxide has two hydrogen atoms and two oxygen atoms; the two oxygen atoms are joined together and one hydrogen atom is attached to each oxygen atom. By contrast, a molecule of the blood protein gamma globulin is made up of 19,996 atoms altogether, but they

are of just four types: carbon, hydrogen, oxygen, and nitrogen.

• SubstanceSubstance: Pure elements and compounds are callled substances

• MixtureMixture: Combination of elements and compounds.

1_15

Matter(materials)

Substances Mixtures

Elements CompoundsHomogeneous

mixtures(solutions)

Heterogeneousmixtures

Physical processes

Chemical

reactions

Classification of Matter

States of Matter

•Phase Phase - A sample of matter that is uniform in composition and physical state and is separated from other phases by a

definite boundary.

Atomic and Molecular Concepts

NucleiNuclei ElectronsElectrons

PlasmaPlasma

Atoms or MoleculesAtoms or Molecules

GasGas

LiquidLiquid

Atoms or MoleculesAtoms or Molecules

Crystalline SolidCrystalline Solid

Tem

per

atu

reT

emp

erat

ure

Elements that exist as gases at 250C and 1 atmosphere

Energy Involved in Phase Changes

Solid

Liquid

Gas

Boiling Condensation

Melting Freezing

RequiresEnergy

LiberatesEnergy

Measurement

• Chemistry is an Observational science.

• Chemistry is a Quantitative science.

• Measurement - A quantitative observation.

Measurement

All measurements have three parts:All measurements have three parts:

1.1. A valueA value

26.97626.97622 gg2. Units2. Units

3.3. An UncertaintyAn Uncertainty

Examples:Examples: 33.2 mL33.2 mL 72.36 mm72.36 mm426 kg426 kg 31 people31 people

MeasurementSystems of Units - Standards of Measurement

1. The Need for Standards

2. The English System (What a pain!!!)

12 in/ft 3 ft/yd 12 in/ft 3 ft/yd 5280 ft/mi5280 ft/mi

16 fl.oz/pt 2 pts/qt 16 fl.oz/pt 2 pts/qt 4 qt/gal4 qt/gal

16 oz/lb 2000 lb/ton16 oz/lb 2000 lb/ton

3. The Metric System - A decimal system

meter (m) - meter (m) - Length Length

liter (L) - liter (L) - Volume Volume

gram (g) - gram (g) - MassMass

Measurement

Metric Examples: 1 m = 1000 mm1 m = 1000 mm 1 kg = 1000 g = 1 000 000 mg1 kg = 1000 g = 1 000 000 mg 10 cm = 0.01 m = 0.000 01 km10 cm = 0.01 m = 0.000 01 km 23 kL = 23 000 000 000 23 kL = 23 000 000 000 LL 1 mL = 0.001 L1 mL = 0.001 L

4. The SI System - Système International d’Unitès

A complete system of units adequate forA complete system of units adequate for the the entire realm of physical science.entire realm of physical science.

SI System of Measurement

Rules for Using the SI SystemsRules for Using the SI Systems

1. Use only singular form of units and do NOT use a period after the symbol for the unit.

2. Use a dot on the base line for the decimal point.

23.6 m23.6 m notnot 23,6 m23,6 m

3. Group digits in threes around the decimal point and do NOT use commas.

1 000 000.000 003 km1 000 000.000 003 km

SI System of Measurement

4. Do NOT use spaces for four-digit

measurements.

1645 mL or 0.2367 g

5. Do NOT use the degree sign (o) for

temperature recorded for the Kelvin

temperature scale.

78.6 K not 78.6 o K

Units

S.I. Units

Length metre, m

Mass Kilogram, kg

Time second, s

Temperature Kelvin, K

Quantity Mole, 6.022×1023 mol-1

Derived Quantities

Force Newton, kg m s-2

Pressure Pascal, kg m-1 s-2

Eenergy Joule, kg m2 s-2

Other Common Units

Length Angstrom, Å, 10-8 cm

Volume Litre, L, 10-3 m3

Energy Calorie, cal, 4.184 J

Pressure

1 Atm = 1.064 x 102 kPa

1 Atm = 760 mm Hg

Common SI-English Equivalent Quantities

Quantity English to SI Equivalent

Length 1 mile = 1.61 km 1 yard = 0.9144 m 1 foot (ft) = 0.3048 m 1 inch = 2.54 cm (exactly!)

Volume 1 cubic foot = 0.0283 m3

1 gallon = 3.785 dm3

1 quart = 0.9464 dm3 (Lt.) 1 quart = 946.4 cm3

1 fluid ounce = 29.6 cm3

Mass 1 pound (lb) = 0.4536 kg 1 pound (lb) = 453.6 g 1 ounce = 28.35 g

Measurement6. Uncertainty in Measurements -6. Uncertainty in Measurements -

Exact Measurements: Exact Measurements: Measured values determinedMeasured values determinedby by countingcounting or when a value is or when a value is defined defined..

Examples:Examples: 31 people31 people 27 rocks 27 rocks2.54 cm = 1 in2.54 cm = 1 in 10 106 6 L = 1LL = 1L

The uncertainty in these measurements = 0The uncertainty in these measurements = 0

Non-exact Measurements: Non-exact Measurements: All other measurements.All other measurements.The last digit recorded is The last digit recorded is uncertainuncertain; it is ; it is estimated!!estimated!!

Examples:Examples: 27.5 g27.5 g 32.7 mm32.7 mm 12 467 km12 467 km1.156 x 101.156 x 1022 mL mL

Accuracy, Precision, & SensitivityAccuracy, Precision, & Sensitivity

Accuracy - Accuracy - The degree to which a measuredvalue agrees with the true or “accepted” value.

Precision - Precision - The reproducibilityreproducibility of a measuredvalue.

Sensitivity - Sensitivity - The “fineness” of a measuredvalue; the number of significant figures it has.

23.5673 g23.5673 g is a more sensitivesensitive measurement than 23.57 g23.57 g.

MeasurementSignificant Figures: Significant Figures: Each digit obtained as a resultEach digit obtained as a resultof a measurement. This includes all of the of a measurement. This includes all of the certaincertaindigits and the first digits and the first uncertain uncertain digit. The number ofdigit. The number ofsignificant figures in a measurement is an indicatorsignificant figures in a measurement is an indicatorof the of the SENSITIVITYSENSITIVITY of the measurement.of the measurement.

How many significant figures are in the following:How many significant figures are in the following:

65 mL65 mL 173.4 g173.4 g 12.2 m12.2 m 1 x 101 x 109 9 nsns

22 4 4 3 3 1 1

MeasurementThe Problem with Zero:The Problem with Zero:

22007.1 mm7.1 mm 00..00002 36 mm2 36 mm 262600.1 mm.1 mm

00.123 .123 0000 mm mm 22004400..00 mm mm 36360000 mm mm

Rules for Significant Figures:Rules for Significant Figures: All non-zero digits are significant. 25.7925.79 km 2727 mL

A zero between other significant figures is significant. 207.9 207.9 nm 100.7 100.7 mL

Measurement Initial zeros are NOT significant.

0.001 231 23 cm3

Final zeros after the decimal point ARE

significant. 23.100 23.100 ps

Final zeros in a measurement with no decimal point may not be significant.

323200 cm Exact measurements have an infinite number of significant figures. (They are (They are CERTAIN!!CERTAIN!!))

MeasurementSignificant Figures in CalculationsSignificant Figures in Calculations: : In aIn ameasurement, the last significant figure is assumed tomeasurement, the last significant figure is assumed tobe be uncertainuncertain..

The result of a calculation involving measured valuesThe result of a calculation involving measured valuescan be no more can be no more certaincertain than the least certain than the least certainmeasurement.measurement.

The number of significant figures in a result dependsThe number of significant figures in a result dependson the number of significant figures in the measure-on the number of significant figures in the measure-mentment andand on the mathematical operation beingon the mathematical operation beingperformed.performed.

MeasurementSignificant Figures in CalculationsSignificant Figures in Calculations::

Addition and Subtraction - Addition and Subtraction - A sum or a dif- ference of two or more measurements has thesame number of decimal placesdecimal places as the measure-ment with the leastleast number of decimal placesdecimal places.

35.2 mL + 0.34 mL = 35.2 mL + 0.34 mL = 35.5 35.5 mLmL

1.007 94 amu+ 1.007 94amu+ 15.9994 amu1.007 94 amu+ 1.007 94amu+ 15.9994 amu= = 18.0153 amu 18.0153 amu

amuamu = atomic mass units = atomic mass units

Measurement Multiplication and Division - Multiplication and Division - A product or

quotient of two or more measurements has thesame number of significant figuressignificant figures as the measure-ment with the leastleast number of significant figuressignificant figures.

density = (9.5760 g)/(12.2density = (9.5760 g)/(12.2 mL)mL)== 0.785 0.785 g/mLg/mL

Round-off Rules - Round-off Rules - For digits 0 - 40 - 4, do not round up.For digits 5 - 95 - 9, round up.

Measurement Round-off the following to two decimal places:Round-off the following to two decimal places:

23.044 39 =23.044 39 =23.04 g23.04 g 65.891 mL65.891 mL = =

45.106 ms45.106 ms = = 30.1149 kg30.1149 kg = =45.11 ms45.11 ms

65.89 mL65.89 mL

30.11 kg30.11 kg

6. Dimensional Analysis - 6. Dimensional Analysis - An extremely useful toolAn extremely useful tool to help you solve mathematical problems. It isto help you solve mathematical problems. It isbased on the fact that when doing calculationsbased on the fact that when doing calculationsinvolving measured quantities, theinvolving measured quantities, the units units must be must beadded, subtracted, divided, or multiplied just likeadded, subtracted, divided, or multiplied just likethe the value value of the measurements.of the measurements.

37.995 ng = 37.995 ng = 38.00 ng38.00 ng

Dimensional Analysis

How many meters are in each of the following?How many meters are in each of the following?

21 km21 km 1023 570 1023 570 mm

(21 km)(1 x 103 m) = 21 x 103 m =km

(1023 570 m)( 1 m ) = (106 m)

1.023 570 m

2.1 x 104 m

Measurement

5. Conversion Factors - 5. Conversion Factors - A fraction whoseA fraction whose numerator and denominator contain the samenumerator and denominator contain the samequantity expressed in quantity expressed in differentdifferent units. units.

1 mile = 5280 ft1 mile = 5280 ft 1 mile5280 ft

= 5280 ft1 mile

= 11

1 cm = 0.01 m1 cm = 0.01 m1 cm 0.01 m

= 0.01 m1 cm

= 11

1 in = 2.54 cm1 in = 2.54 cm2.54 cm 1 in

= 1 in 2.54 cm

= 11

Dimensional Analysis How many mL are in 3.0 ftHow many mL are in 3.0 ft33??

1 ft = 12 in1 ft = 12 in 1 in = 2.54 cm1 in = 2.54 cm 1 cm1 cm33 = 1 mL = 1 mL

(3.0 ft3)(12 in)(12 in)(12 in)(2.54 cm)(2.54 cm)(2.54 cm)(1 mL) (1 ft) (1 ft) (1 ft) (1 in) (1 in) (1 in) (1 cm3)

= = 8.5 x 108.5 x 1044 mL

How many ns are in 23.8 s?How many ns are in 23.8 s?

(23.8 s)(109 ns) (1 s)

= = 23.8 x 1023.8 x 109 9 ns ns = = 2.38 x 102.38 x 1010 10 nsns

Classification of Properties of Matter

• Properties can be classified as:– Physical or Chemical Properties– Intensive or Extensive Properties

Prentice-Hall © 2002

Properties of Matter

• Physical Properties - Properties that do NOT involve substances changing into other substances.

Melting Point Boiling Point

Temperature Density

Mass Volume

• Chemical Properties - Properties that involve substances changing into other substances.

Chemical Reactivity Reduction Potential

Flammability Oxidation Potential

Prentice-Hall © 2002

Properties of Matter

• Extensive Properties - Properties that depend on the amount of matter present in a sample.

Mass Volume Heat Capacity

• Intensive Properties - Properties that do NOT depend on the amount of matter present in a sample.

Color Temperature Density

Melting Point Specific Heat Boiling Point

Mass and WeightMass and Weight

Mass: Mass: the measure of the quantity or amount of the measure of the quantity or amount of matter in an object. The mass of an object does notmatter in an object. The mass of an object does notchange as its position changes.change as its position changes.

Weight: Weight: A measure of the gravitational attraction ofA measure of the gravitational attraction ofthe earth for an object. The weight of an objectthe earth for an object. The weight of an objectchanges with its distance from the center of the earth.changes with its distance from the center of the earth.

Mass is measured using a Mass is measured using a BALANCEBALANCE..

Weight is measured using Weight is measured using SCALESSCALES..

Sample Calculations Involving MassesSample Calculations Involving Masses

1.1 How many 1.1 How many g are in 2.56 kg?g are in 2.56 kg?

(2.56 kg)(10(2.56 kg)(1033 g)(10 g)(1066g)g) (1 kg) ( 1 g)(1 kg) ( 1 g) = = 2.56 x 102.56 x 1099 gg

1.2 How many g are in 2.578 x 101.2 How many g are in 2.578 x 101212 ng? ng?

(2.578 x 10(2.578 x 101212 ng) (1 g) ng) (1 g) (10(1099 ng) ng) = = 2578 g2578 g

Volume

Sample Calculations Involving VolumesSample Calculations Involving Volumes

1.3 How many mL are in 3.456 L?1.3 How many mL are in 3.456 L?

(3.456 L)((3.456 L)(1000 mL1000 mL)) LL

= = 3456 mL3456 mL

1.4 How many 1.4 How many L are in 23.7 cmL are in 23.7 cm33??

(23.7 cm(23.7 cm33)()( 1 mL 1 mL )()( 1 L_ _ 1 L_ _)()(101066 L)L) (1 cm(1 cm33)(1000 mL)( 1L ))(1000 mL)( 1L )

= = 2.37 x 10 2.37 x 10 44 LL= = 23 700 23 700 LL

DensityDensity - Density - The mass of a unit volume of a material.The mass of a unit volume of a material.

density = mass/volumedensity = mass/volume

1.5 What is the density of a cubic block of wood that is 1.5 What is the density of a cubic block of wood that is 2.4 cm on each side and has a mass of 9.57 g? 2.4 cm on each side and has a mass of 9.57 g?

volume = [2.4 cm x 2.4 cm x 2.4 cm]volume = [2.4 cm x 2.4 cm x 2.4 cm]

density = (9.57 g)/(13.density = (9.57 g)/(13.88 cmcm33))

= 0.69 g/cm= 0.69 g/cm33 = 0.69 g/mL= 0.69 g/mLNote that 1 cmNote that 1 cm33 = 1 mL = 1 mL

Conversion

What is the mass of a cube of osmium that is 1.25 inches on each side?

Have volume, need density = 22.48g/cm3

Temperature and Thermal Energy

Temperature: Temperature: A measure of the “hotness”and “cold-A measure of the “hotness”and “cold-ness” of an object; a measure of the average kineticness” of an object; a measure of the average kineticenergy of the atoms and molecules of the object.energy of the atoms and molecules of the object.The higher the temperature, the more kinetic energyThe higher the temperature, the more kinetic energythe atoms and/or molecules have. This is anthe atoms and/or molecules have. This is anINTENSIVEINTENSIVE property.property.

Thermal Energy: Thermal Energy: Often called Often called “heat”“heat”, it is the form, it is the formof energy toward which all other forms tend to go.of energy toward which all other forms tend to go.

Temperature

F = (5/9) (C + 32)

Relative Temperatures

K = T + 273.15

Sample Calculations Involving Temperatures

1.6 Convert 73.61.6 Convert 73.6ooF to Celsius and Kelvin temperatures.F to Celsius and Kelvin temperatures.

ooC = (5/9)(73.6C = (5/9)(73.6ooF - 32) = (5/9)(41.6)F - 32) = (5/9)(41.6)

ooC = (5/9)(C = (5/9)(ooF - 32)F - 32) K = K = ooC + 273.15C + 273.15

= = 23.123.1ooCC

K = 23.1K = 23.1ooC + 273.15 = C + 273.15 = 296.3 K296.3 K

MemorizeMemorize

End of Chapter Questions