Matter And Measurement Chapter 1: Introduction - Matter and Measurement John D. Bookstaver St....

MatterAnd

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Chapter 1:Introduction -Matter and

MeasurementJohn D. Bookstaver

St. Charles Community College

St. Peters, MO

2006, Prentice Hall

Chemistry, The Central Science, 10th editionTheodore L. Brown; H. Eugene LeMay, Jr.;

and Bruce E. Bursten

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1.1 ----------------------------------

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The Study of Chemistry

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Let the journey Begin…

• Chemistry - the study of matter and the changes it undergoes.

• Matter – anything that has mass and takes up space

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Overview of

Matter

• Atoms are the building blocks of matter.• Each element is made of the same kind of atom.• A compound is made from atoms of two or more

of elements that are chemically bonded together

Atoms

Elements

Compounds

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Classification of Matter

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States of Matter

• Solid - has definite shape and volume• Liquid - definite volume; takes shape of its

container• Gas - no definite volume or shape; takes the

shape of its container and is easily compressed/ expanded

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• Matter that does not vary from sample to sample

• Have distinct or characteristic properties

• Have a constant composition

• All pure substances are either:– Elements– Compounds

Pure Substances:

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Elements:

• Made up of only one type of atom

• Cannot be decomposed into smaller substances

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Compounds:

• Composed of two or more elements

• Ex: Water (H & O), salt (Na & Cl),

• Can be broken down into their elemental particles only by chemical means

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Law of Constant Composition:

• A pure compound is always made of the exact same elemental composition (by mass)

• aka - Law of Definite Proportions

• Ex: Water (H2O) is always in the ratio of

1 oxygen : 2 hydrogen

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• Physical combination of many substances• Composition can vary (variable)• Each substance in a mixture retains its

own chemical identity & properties― allows the mixture to be separated by

physical means into its component substances

• 2 types:― Heterogeneous― Homogenous

Mixtures:

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• Heterogeneous Mixture― Not uniform in appearance, composition,

or properties ― Appears as two or more distinct phases― Ex: sand, chocolate chip cookies

• Homogenous Mixture (aka: solution)― Uniform throughout― appears as one phase―Ex: air, saltwater, iced tea, Kool-Aid

Mixtures: (con’t)…

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Classification of Matter

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Propertiesof Matter

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Properties of Matter

• Physical Properties:– Observed without changing the substance

into a different substance.• Boiling point, density, mass, color, volume, etc.

• Chemical Properties:– Only observed when the substance is

changed into a different substance.• Flammability, corrosiveness, reactivity with

acid, etc.

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Properties of Matter (con’t)

• Intensive Properties:– Independent of the amount/quantity of

matter present.– Only depends on the identity of the

substance• Density, boiling point, color, state of matter, etc.

• Extensive Properties:– Dependent upon the amount of matter

present.• Mass, volume, energy, etc.

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Changes of Matter

• Physical Changes:– Changes in matter that do NOT change the

composition of a substance.• Changes of state, temperature, volume, etc.

• Chemical Changes (aka: reaction = rxn):– Changes where the original substances

(reactants) are converted to all new & different substances (products). • Combustion, oxidation, decomposition, etc.

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Chemical vs. Physical

• Ask, “Can I somehow get the original substance back?”– Yes = phys. change– No = chem. change

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Separation of Mixtures:

• Components of a mixture retain their individual properties―Ex) magnetism, color, density

• A mixture can be physically separated based on these properties― Each component remains chemically

unchanged

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• Physical combination of 2 or more substances

• Composition is variable• Each substance in a mixture retains its own

chemical identity & properties― allows the mixture to be separated by physical

means into its component substances (by exploiting differences in properties)

• 2 types of mixtures:― Heterogeneous = 2+ phases visible― Homogenous = uniform (appears as 1 phase)

Mixtures:

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3 Methods to Separate a Mixture:

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Filtration:

• Separates solid substances from liquids (and solutions)

• Method is based on separation due to differences in the size of component particles in the mixture

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Distillation:

Separates a homogeneous mixture on the basis of differences in boiling point.

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Chromatography:

Separates on the basis of differences in the solubility of the components in the mixture

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The Scientific Method

A systematic approach to solving problems

(textbook pg. 13)

Collect data

Analyze data

Tentative explanation that can be TESTED by

experimentation

Generally explains causes of phenomena

and can be used to make

predictions

Theory of Evolution

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Units of Measurement

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Qualitative vs. Quantitative• Qualitative:

– Observations which reflect a physical description of data

– Ex) color, odor, texture, state of matter

• Quantitative:– Measurements that include numerical

quantities & data– Mass, volume, density, length, time

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SI Units

• Système International d’Unités• Base unit for each quantity

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Example: SI UNITS

Density = =

Derived Units:

massvolume

kgm3

• How would you determine the density of a liquid? … A solid?

• Two or more SI base units combined

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• Format : M x 10n

– M is a # between 1-9– n is an integer

• Insert an understood decimal point• Decide where the decimal point needs to end up• Count the # of places that the decimal point moved (n)

• #s SMALLER than 1 = negative exponent, – 0.00000765 =

• #s BIGGER than 1 = positive exponent, – 2, 340,000, 000 =

Scientific Notation

2.5 x 109

7.65 x 10-6

2.34 x 109

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Temperature:

• Two Definitions:

1) A measure of the average kinetic energy of particles in a sample

2) A measure of how much heat is in an object

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Temperature Conversions

• SI base unit = Kelvin

K = C + 273.15

C = (F − 32) 1.8

F = (1.8)(C) + 32

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Metric System

• Prefixes convert the base units into units that are appropriate for the item being measured.

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• Prefixes convert the base units into units that are appropriate for the item being measured.

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Metric Conversions

• The is a way to SHOW your work & units

Ex #1) Covert 945 mL to dL

945 mL x 1 dL =

100 mL

Ex #2) Convert 0.54 km to nm

0.54 km x 1000 m x 109 nm =

1 km 1 m

9.45 dL

5.4 x 1011 nm

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Density

• Density = mass

volume

Volume• Most common units:

L , mL, cm3

• 1 mL = 1 cm3

• 1 L = 1 dm3

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Uncertainty in

Measurement

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• Accuracy - how close, a measurement is to the true or known value.

• Precision - several measurements with values very close to one other.

Accuracy vs. Precision

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% error = experimental – theoretical x 100

theoretical

• Experimental = the value determined from lab data

• Theoretical = the standard, true, or accepted value

Percent Error Calculation

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Exact vs. Inexact• Exact #s:

– conversion factors & counting Ex) There are exactly 5280 ft in 1 mile

“ “ “ 12 eggs in 1 dozen

“ “ “ 2.54 cm in 1inch

• Inexact #s:– Measurements from equipment w/limitations

Ex) Mass using balance, volume using graduated cylinder

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Uncertainty in MeasurementsDifferent pieces of equipment have different uses and different degrees of accuracy.

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Uncertainty in Measurements

What is the unit of volume here that can be measured with absolute certainty?

All measurements must be reported with the last digit being uncertain (1 degree of uncertainty)

43 mL … (last volume level marked on equipment)

Report the measurement as 43.0 mL (or 43.1 mL)

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Significant Figures

• The term significant figures refers to digits that were “measured” or inexact #s.

• We pay attention to sig. figs. so we do not overstate the accuracy of our calculations

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Rules for Sig. Figs

• All non-zero digits (1-9) are significant

• Zeroes between any non-zero digits are significant (ex: 101)

• Zeroes at the beginning of a number, before any non-zero digits, are NEVER significant (ex: 0.082)

• Zeroes at the end of a number are significant ONLY if the number contains a decimal point (ex: 35.0 vs. 350 vs. 350. )

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Practice with Sig. Figs

1) 0.0034567 g

2) 2.97 mL

3) 6.700 x 103 m

4) 598.300 K

5) 78,000 mg

6) 9020 cm

5 sig figs

3 sig figs

4 sig figs

6 sig figs

2 sig figs

3 sig figs

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Calculations with Significant Figures

• Addition or Subtraction:– Round the answer off to the least number

of decimal places

Ex: 20.42 g + 7.9764 g

+ 102.3 g

130.6964 g

The answer you report is = 130.7 g (1 dec. place)

(2 dec. places)(4 dec. places)(1 dec. place)

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Practice +/- with Sig. Figs

1) 15.002 cm + 24.1104 cm

2) 22.35 kg – 0.154 kg

3) 100.0 g – 23.73 g

4) 2.030 mL – 1.870 mL

39.112 cm3 dec. places

22.20 kg2 dec. places

76.3 g1 dec. places

0.160 mL3 dec. places

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• Multiplication or Division:– Round the answer off to the least number

of significant figures

Ex: 6.221 cm x

5.2 cm

32.3492 cm2

The answer you report is = 32 cm2 (2 sig. figs)

(4 sig. figs)

(2 sig. figs)

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Practice x /÷ with Sig. Figs

1) 75.246 g / 6.33 mL

2) 16.00 cm x 2.5 cm x 3.66 cm

3) 3.24 m x 7.00 m

4) 710 m / 3.0 s

11.9 g/mL3 sig. figs

150 cm3

2 sig. figs

22.7 m2

3 sig. figs

240 m/s or 2.4 x 102 m/s 2 sig. figs

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• Mixed Operations– Always follow PEMDAS– Count sig. figs as you go to get it correct

Ex: (44.739 g – 36.21 g) = =

1.3 mL

The answer you report is = 6.6 g/mL (2 sig. figs)

(8.53 g)

1.3 mL

6.561 g/mL

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Dimensional Analysis

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Dimensional Analysis (D.A.)

• It’s all about converting & matching up units!

• D.A. provides a way to check your solution method by following the units or “dimensions” of the problem

• When a problem is solved correctly, the unwanted units will cancel with one another. This will leave behind only the desired units!

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Dimensional Analysis

• Units are…– multiplied together– divided into one another– or canceled completely

• The KEY is…– Finding the correct conversion relationship– Setting up the relationships to cancel out the units

you don’t want

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D.A. – Relationships as Ratios• How many dozens are there in 48 eggs?• HOW did you solve that problem?

• What are the relationships/conversions you needed to know?– Relationships:

1 dozen eggs = 12 eggs

If we write that relationship as a ratio…

1 dozen - or - 12 eggs .

12 eggs 1 dozen

• How did you know to multiply or divide?– D.A. will figure that out for you!!!

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D.A. - Problem Solving Steps1) Identify the unknown (what are you solving for).

2) Write down what is given or known

3) Look for relationships between the known and unknown.

*Note: it may take more than one relationship to get from the known to the unknown*

4) Arrange the relationships into a series of ratios so that the given unit will cancel out and the desired unit will be left over.

5) Multiply numbers across the top of the expression and divide numbers on the bottom.

given unit desired unit = desired unit

given unit

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D.A. - Intro Example• How many large pizzas will I need for a

birthday party with 22 guests if I allow each person to eat 4 slices?

Unknown/Solving for: # of pizzas

Given: 22 guests, 4 slices each

Relationships: 1 pizza = 8 slices, 1 guest = 4 slices

*** Must set up the relationships like ratios (which are equal to 1)

1 pizza8 slices

8 slices1 pizza

1 guest4 slices

4 slices1 guest

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Example Continued

22 guests x 4 slices x

1 guest

1 pizza =

8 slices

11

pizzas

Given in problem

All the unwanted units should cancel (top cancels w/matching bottom) leaving only the desired quantity and units

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TURTLE ISLAND• On an island with 375 inhabitants, the people

are growing very concerned with the declining number of turtles

• Your mission is to calculate how many turtles are on the island.

• Use the clues the inhabitants found to get an accurate count of the turtle population

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Need-to-Know Conversions

Length• 1 inch = 2.54 cm• 1 mile = 5280 ft• 1 ft = 12 in

Volume• 1 cm3 = 1 mL

• All the metric prefixes

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1) How many decimeters are equal to 32.74 yards?

2) How many milliliters are in 2.35 gal of water?

299.37456 dm (w/sig.fig = 299.4 dm)

8892.4 mL (w/sig.fig = 8890 mL)

Practice the following using dimensional analysis and sig figs:

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Derived Units as Ratios

• All derived units can be set-up as ratios for use in D.A.

– Density Example:

1.35 g/mL =

1.35 g

1 mL

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Practice the following using dimensional analysis and sig figs:

4) What is the mass of gold, which has a density of 19.3 g/cm3 and a volume of 2.0 in3?

632.54 g (w/sig. fig = 630 g)

3) The speed of light is 3.0 x108 m/s. What is the speed of light in mi/hr (mph)?

6.7108 x108 mi/hr (w/sig. figs = 6.7 x 108 mph)

Matter And Measurement Chapter 1: Introduction - Matter and Measurement John D. Bookstaver St....

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