Unit 1 Measurement and data processing chemistry 2014-2015 Mr.Trimble 1.

Unit 1

Measurement

and data

proce

ssing

chemist

ry 2014-2015

Mr .Trim

ble

1

UNCERTAINTIES IN MEASUREMENTS 2

3

UNCERTAINTY IN MEASUREMENTS

Different measuring devices have different uses and different degrees of precision

4

ESTIMATING UNCERTAINTY

• To estimate a measuring devices uncertainty, read to ½ the smallest scale division.

INDICATING UNCERTAINTY IN A MEASUREMENT• Indicate using + after the recorded

measurement stating the ½ beyond the smallest scale division followed by the units

• Example:

5

For digital devices, uncertainty is always +/- .1 of the smallest scale

6

INDICATING UNCERTAINTY IN A MEASUREMENT

SIGNIFI

CANT FIG

URES IN

MEASUREMENTS

7

Significant figures - Counting• Counting Significant Figures – Atlantic Pacific Method

• Decimal point absent (Atlantic) – begin on right side of number and cross out all zero digits until first non – zero digit, remaining digits are significant

• Decimal point present (Pacific) - begin on left side of number and cross out all zero digits until first non – zero digit, remaining digits are significant

• Exact numbers – have infinite number of significant figures

8

Examples: Counting Sig. Figs• 234 cm• 67000 cm• 45000 cm• 560. cm• 0.5630 cm• 1.0034 cm• 0.00467 cm

9

10

SCIENTIFIC NOTATION

• Used to shorten large or small numbers

• Standard FormBase - one number to the left of the decimal place, with all the significant figures shown

Exponent – is the number of places the decimal point must be shifted to give the number in standard formNegative Exponent – value less than 1Positive Exponent – value greater than 1

11

EXAMPLES: SCIENTIFIC NOTATION

124300 =

0.00362 =

1300000 =

1.23E-5 =

5.61 x 10^6 =

UNCERTAIN

TIES IN

CALCULA

TED R

ESULTS

12

SIGNIFICANT FIGURES IN MATHEMATICAL OPERATIONS• Addition/Subtraction – the answer will

have the same number of decimal places as the measurement with the fewest decimal places

• Multiplication/Division – the answer will have the same number of significant figures as the measurement with the fewest significant figures

• Mixed operations – follow the order of operations, determining s.f. for each step, do not round until the end!

13

14

EXAMPLE OF SIG. FIGS IN CALCULATIONS33.5 cm + 7.88 cm + 0.977 cm =

23000 km + 8.7 km =

67.23 cm x 9.22 cm =

(200 cm x 3.333) + (300 x 1.35) cm =

15

CALCULATORS AND SCIENTIFIC NOTATION• Calculators handle scientific

notation by only inputting the exponent, using an EXP or EE keyenter the base as you would a regular number, then press EXP or EE, then enter the exponent

• Display – the calculator used E to show exponent ( E means x 10 )

EXPERIM

ENTAL E

RRORS

16

RANDOM ERRORS - PRECISION

• Random errors - Precision•A random error makes the measured value both smaller and larger than the true value (this happens by chance alone)

•Reduce random errors- repeat the experiment

17

SYSTEMATIC ERRORS

• Systematic error - Accuracy • Errors due to "incorrect" use of

equipment or poor experimental design•Makes the measured value always smaller or larger than the true value, but not both.

•An experiment may involve more than one systematic error and these errors may nullify one another

18

CATEGORIES OF SYSTEMATIC ERRORS

• Personal errors – the result of ignorance, carelessness, prejudices, or physical limitations on the experimenter.

• Instrumental Errors - attributed to imperfections in the tools with which the analyst works.

• Method Errors - results when you do not consider how to control an experiment.

19

DETERMINATION OF ERROR

• Accuracy - how close a measurement is to the true value of a quantity.

• Precision - how close several measurements are to each other.

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PRECISION VS. ACCURACY

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EVALUATING ACCURACY AND PRECISION IN DATA• When evaluating whether data is

accurate or precise you could look at Accuracy Precision

Single Data Point

A single data point is accurate if it is close the literature value

A single data point is precise if it has many decimal places/significant figures. Meaning, you used a very precise tool or method for measuring the value.

Set of Data Points

A set of data points are accurate if the average or mean is close to the literature value. This is the reason we do multiple trials. Any random errors cancel out when the average is taken, thus random error is reduced.

A set of data points are precise if they are all very close together. This definition refers to the consistency of the data.

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PERCENTAGE UNCERTAINTIES AND PROPAGATION OF ERRORS

23

Percent Uncertainties and errors• Percent Error

• (accepted value – experimental value) x 100• accepted value

24

CH. 3 - MEASUREMENT

Temperature Conversions

A. Temperature• Temperature

• measure of the average KE of the particles in a sample of matter

273.15Kelvin Co

32Fahrenheit Co 5

9

32)Celsius F( 9

5 o

A. Temperature• Convert these temperatures:

1) 25oC = ______________K

2) -15oF = ______________ K

3) 315K = ______________ oC

4) 288K = ______________ oF

Dimensional AnalysisConversion Factors Problems

CH. 3 - MEASUREMENT

A. Problem-Solving Steps

1. Analyze

2. Plan

3. Compute

4. Evaluate

B. Dimensional Analysis• Dimensional Analysis

• A tool often used in science for converting units within a measurement system

• Conversion Factor• A numerical factor by which a quantity expressed in

one system of units may be converted to another system

3

3

cm

gcm

B. Dimensional Analysis• The “Factor-Label” Method

• Units, or “labels” are canceled, or “factored” out

g

B. Dimensional Analysis

• Steps to solving problems:

1. Identify starting & ending units.

2. Line up conversion factors so units cancel.

3. Multiply all top numbers & divide by each bottom number.

4. Check units & answer.

C. Conversion FactorsFractions in which the numerator and denominator are EQUAL quantities expressed in different units

Example: 1 in. = 2.54 cm

Factors: 1 in. and 2.54 cm 2.54 cm 1 in.

How many minutes are in 2.5 hours?

Conversion factor

cancel

By using dimensional analysis / factor-label method, the UNITS ensure that you have the conversion right side up, and the UNITS are calculated as well as the

numbers!

2.5 hr

1 x 60 min

1 hr = 150 min

C. Conversion Factors

Write conversion factors that relate each of the following pairs of units:

1. Liters and mL2. Hours and minutes3. Meters and kilometers

Learning Check:

E. Dimensional Analysis Practice

• You have $7.25 in your pocket in quarters. How many quarters do you have?

X$7.25 1 1 dollar

4 quarters


How many seconds are in 1.4 days?

= 12000 s

1.4 days

24 hr 60 min

60 s

1 day

1 hr 1 min


• How many milliliters are in 1.00 quart of milk?


• You have 1.5 pounds of gold. Find its volume in cm3 if the density of gold is 19.3 g/cm3.

Your European hairdresser wants to cut your hair 8.0 cm shorter. How many inches will he be cutting off?


Roswell football needs 550 cm for a 1st down. How many yards is this?


A piece of wire is 1.3 m long. How many 1.5-cm pieces can be cut from this wire?



• How many liters of water would fill a container that measures 75.0 in3?

Base Units

• In the SI (Le Système International d'Unités) system of measurement there are seven base units

• These units are used in the measurement of different quantities and independent of each other

Base UnitQuantity Unit Symbol

Mass Grams g

Length Meter m

Amount of Substance

Mole mol

Time Second s

Electric Current Ampere A

Temperature Kelvin K

Luminous Intensity candela cd

D. SI Prefix Conversions

•Memorize the following chart. (next slide)

•Find the conversion factor(s).

•Insert the conversion factor(s) to get to the correct units.

•When converting to or from a base unit, there will only be one step. To convert to or from any other units, there will be two steps.

mega- M 106

deci- d 10-1

centi- c 10-2

milli- m 10-3

Prefix Symbol Factor

micro- 10-6

nano- n 10-9

kilo- k 103

BASE UNIT --- 100

giga- G 109

deka- da 101

hecto- h 102

tera- T 1012

mo

ve le

ft

mo

ve r

igh

tA. SI Prefix Conversions

pico- p 10-12

D. SI Prefix Conversions1 T(base) = 1 000 000 000 000(base) = 1012 (base)

1 G(base) = 1 000 000 000 (base) = 109 (base)

1 M(base) = 1 000 000 (base) = 106 (base)

1 k(base) = 1 000 (base) = 103 (base)

1 h(base) = 100 (base) = 102 (base)

1 da(base) = 101 (base)

1 (base) = 1 (base)

10 d(base) = 1(base)

100 c(base) = 1 (base)

1000 m (base) = 1(base)

1 (base) = 1 000 000 µ = 10-6(base)

1 (base) = 1 000 000 000 n = 10-9(base)

1 (base) = 1 000 000 000 000 p = 10-12(base)

Tera-

Giga-

Mega-

Kilo-

Hecto-

Deka-

Base

Deci-

Centi-

Milli-

Micro-

Nano-

Pico-


a. cm to m

b. m to µm

c. ns to s

d. kg to g


1) 20 cm = ______________ m

2) 0.032 L = ______________ mL

3) 45 m = ______________ m


4) 805 Tb = ______________ b

Terabytes bytes


1) 400. g = ______________ kg

1) 57 Mm = ______________ nm

• 1 cm3 = _______________ dm3

• 1 kL2 = _____________ L2

53

DENSITY FUN

Objective•Student will demonstrate an understanding of density by using it to perform different unit conversions.

A. Derived Units• Combination of base units• Volume – length length length

1 cm3 = 1 mL 1 dm3 = 1 L• Density – mass per unit volume (g/cm3)

D = MV D

M

V

Broken Heart

Density• Density is a physical property of matter, as each element and

compound has a unique density associated with it.

• The density of any sample of a substance at the same temperature will always be the same.

Density• If a substance is more dense than a liquid it will sink in it, if it is

less dense than it will float

• The density of water is 1.00 g/mL, however this varies slightly with temperature

B. Density Calculations• An object has a volume of 825 cm3 and a density of 13.6

g/cm3. Find its mass.

GIVEN:

V = 825 cm3

D = 13.6 g/cm3

M = ?

WORK:

M = DV

M = (13.6 g/cm3)(825cm3)

M = 11,220 g

M = 11,200 gDM

V

B. Density CalculationsA liquid has a density of 0.87 g/mL. What volume is

occupied by 25 g of the liquid?

GIVEN:

D = 0.87 g/mL

V = ?

M = 25 g

DM

V

WORK:

V = M D

V = 25 g

0.87 g/mL

V = 28.7 mL = 29 mL

B. Density CalculationsYou have a sample with a mass of 620 g and a volume of 753

cm3. Find its density.

GIVEN:

M = 620 g

V = 753 cm3

D = ?

DM

V

WORK:

D = M V

D = 620 g

753 cm3

D = 0.82 g/cm3

C. Density Calculations with DA

• Used when units do not agree• Conversions must be made before using formula

D = MV

D = g

cm3

• You have 3.10 pounds of gold. Find its volume in cm3 if the density of gold is 19.3 g/cm3.

lb cm3

3.10 lb 1 kg

2.2 lb= 73.0 cm3

1000 g

1 kg

1 cm3

19.3 g


• You have 0.500 L of water. Find its mass in ounces if the density of water is 1.00 g/cm3.

L oz0.500 L 1000 mL

1L

= 17.6 oz

1.00g

1 cm3


1 cm3

1 mL1 kg

1000 g16 oz

1lb

2.2 lbs

1kg

• I threw a plastic ball in the pool for my dog to fetch. The mass of the ball was 125 grams.

• What must the volume be to have a density of 0.500 g/mL. ( I want it to float of course!)

• A little aluminum boat (mass of 14.50 g) has a volume of 450.00 cm3. The boat is place in a small pool of water and carefully filled with pennies. If each penny has a mass of 2.50 g, how many pennies can be added to the boat before it sinks?

Matter• Matter is defined as anything that has mass and takes up

space.• Volume is the amount of space matter occupies.• The smallest building block of matter is the atom.

States of Matter• Solid- definite

volume and definite shape; vibrate at fixed points

• Liquid- has a definite volume but an indefinite shape; particles can move past one another and will take the shape of its container

• Gas- Neither a definite volume nor a definite shape; particles move rapidly and will take the shape of its container

• Plasma- high temperature state in which atoms lose electrons (Sun and stars)

Pure Substances

• Element- The simplest type of a pure substance made of only one kind of atom. The smallest particle of an element is an atom.

Elements are list on the periodic table• Compounds- Pure substances that are made of

two or more elements that are chemically bonded. Examples are sugar, carbon dioxide, ammonia, baking soda, and vinegar.

Identify the following as an Element or Compound• Na• MgO2

• As• Cl2

• H2O

• Fe(NO3)2

70

Elements

ElementsCompound

Elements

Compound

Compound

Compounds

• Compounds can be broken down into simpler substances while elements cannot. Heat and electricity are often used to break apart compounds.

• The properties of a compound are very different from the elements that make up the compound. Ex.- Salt is a compound made from sodium ( a silvery metal) and chlorine ( a greenish poisonous gas)

Compounds

• The smallest particle of a compound is a molecule. Molecules of a compound are all alike. Water in the ocean is like water that comes from the faucet.

• Atoms are represented by symbols and molecules are represented by formulas.

• Every element and compound has chemical properties that are used for classification.

Classifying Matter

• Matter is classified into one of two groups: Pure Substances or Mixtures

• Pure substances include compounds and elements and they have a definite composition.

• Mixtures contain more than one substance; therefore, the composition varies.

Mixtures• A mixture is a blend of two or more kinds of

matter which retains its own identity and properties.

• Mixtures that have a uniform composition are called homogeneous mixtures or solutions.

• Mixtures that do not have uniform compositions are called heterogeneous mixtures.

• In Chemistry, the amount of substance present in a mixture is referred to by percent by mass of the substance.

Properties• Physical properties

can be observed without changing the identity of the substance.• Melting Point• Boiling Point• Freezing Point• Color• Mass

• Chemical Properties can be observed when new substances are produced.• Ability to Burn• Ability to Tarnish• Ability to Rust

Types of Physical Properties

• Extensive properties depend on the amount of matter present.• Volume• Mass• Amount of Energy

• Intensive properties do not depend on the amount of matter present.• Melting Point• Boiling Point• Density• Conductivity

Physical Change

• A change is a substance that does not involve a change in the identity of the substance

• Examples include• Breaking• Tearing• Changes of state• Grinding• Cutting

Phase Changes

• When a substance undergoes a phase change it is a physical change

• Melting: Solid to liquid• Freezing: Liquid to solid• Vaporizing: Liquid to gas• Condensing: Gas to liquid• Sublimation: Solid to gas• Deposition: Gas to Solid 80

Chemical Change

• Chemical Reactions- changing of a substance into something else

• Examples include • Burning, Rusting, Tarnishing, Reacting

• Reactants- substances that react• Products- Substances that are produced

Chemical or Physical Change

• Dissolving salt in water• Burning oil• Melting solid iron• Sugar ferments • Copper sulfate crystals are broken• Paper is torn in half• Iron rusts• An egg is cooked

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Chemical ChangePhysical Change

Physical ChangeChemical Change

Physical Change

Physical ChangeChemical ChangeChemical Change

Unit 1 Measurement and data processing chemistry 2014-2015 Mr.Trimble 1.

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Transcript of Unit 1 Measurement and data processing chemistry 2014-2015 Mr.Trimble 1.