Chapter 2



Chemistry:

Matter and Measurements

Objectives: -To understand the importance of learning chemistry

-To define chemistry

-To understand scientific thinking

-To describe the method scientists use to study nature

-To develop successful strategies for learning chemistry

-To show how very large or very small numbers can be expressed

as the product of a number between 1 and 10 and a power of 10.

- To learn the English, metric, and SI systems of measurements

-To understand the metric system for measuring length, volume,

and mass

-To understand how uncertainty in a measurement arises

-To learn to indicate the measurement's uncertainty by using

significant figures

-To learn to determine the number of significant figures in a

calculated result

-To learn how dimensional analysis can be used to solve various

types of problems

-To learn the three temperature scales

- To learn to convert from one scale to another; continue to

develop problem solving skills

-To define density and its units

-To learn about matter and its three states

-To learn to distinguish between physical and chemical properties

-To learn to distinguish between physical and chemical changes

-To understand the definitions of elements and compounds

-To learn to distinguish between mixtures and pure substances

-To learn two methods of separating mixtures

-To understand energy and its effects on mater

Introduction:

What is chemistry? Chemistry is the study of composition, properties and reaction changes of properties of substances. The science of chemicals/chemical reactions.

Chemistry is central to all other sciences.

There are 5 Subdivisions of chemistry:

Biochemistry Inorganic Chemistry Analytical Chemistry

Organic Chemistry Physical Chemistry

Start to think like a chemist(Chemist study problem solving methods

Why Learn Chemistry

Like any language, you will learn the language and vocabulary for you can communicate effectively in Chemistry.

There are no disasters or mistakes in experimentation as long as you learn from it.

Questioning, Problems, and Memorization are important as you explore Chemistry.

Solving Problems Using a Scientific Approach

The Scientific Method- a procedure for solving problems. Used to gather information

Ex. Michael Faraday ( electric motors; today = refrigerators, car starters, hair

dryers, ect.

Scientific Theory- is a set of tested hypotheses that give an overall explanation

- a logical and time tested explanation of a phenomenon

Scientific Law (Natural Law)-

Natural Science-

Fact-

Theory-

Hypothesis-

1.2 Measurements

Predictions are made when forming a hypothesis. This is why we study science.

The numbers of scientific measurements are very large or very small.

Ex. distance Earth is from the Sun, size of a cell

To make these numbers more compact, we use Scientific Notation

Scientific Notation - a product of a number between 1 and 10 AND the appropriate power of 10

Ex.

125 1.25 x 10

1700 1.7 x 10

93,000,000,000 9.3 x 10

0.010 1.0 x 10

0.000167 1.67 x 10

Using Scientific Notation

1. Any number can be represented as the product of number between 1 and 10 and a power of 10 (either positive or negative)

2. The power of 10 depends on the number of places the decimal point is moved and in which direction. The number of places the decimal point is moved determine the power of 10. The direction of the move determines whether the power of 10 is positive or negative. If the decimal point is moved to the left, the power of 10 is negative. If the decimal point is moved to the right, the power of 10 is positive.

Example 2.1

Write in scientific notation.

A) 238,000 B) 15,000,000

Example 2.2

Write in scientific notation

A) 0.00043 B) 0.089

Example 2.3

Take the number out of scientific notation

A) 6.443 x 103

B) 5.991 x 10-5

C) 2.001 x 104

D) 1.997 x 10-3

E) 7.871 x 10-1

F) 1.001 x 101

Unit-

Need for a scale or standard.

English vs. Metric vs. International

Units of Measurement (SI System)- International System of Units

Table 1

A) Base/Fundamental Units (6)

Physical Quality Name Abbrev.

1.Mass Kilogram kg

2.Length Meter m

3.Time Seconds s

4.Temperature Kelvin k

5.Amount of Substance Mole mol

6.Electric Current Ampere

Table 1.2 (p7)

B) SI Prefixes

Prefix Abbrev. Factor

1.Mega M 1,000,000

2. k

3.deci

4. 1x10-2 or 0.010

5. m

6.micro

7. n

8. 1x10-12 or 1/Trillion

C) Derived Units - combination of 2 or more base units

Volume = length x width x height

V= lwh

Liters is a unit of volume Table 1.3 (p12)

1l is defined as 1 dm3

1ml is 1 cm3

D) Weight vs. Mass & Weighing

Mass is measure of resistance of an object to some force. (any object that takes up space has mass)

Weight is a measure of the pull of gravity of some object that has mass.

Balance ---- mass

Scales ------ weight

Example elementary problem: "Which is heavier a pound of lead or a pound of feathers?"

- here students are thinking about density, not mass.

E) Density - amount of matter present in a given volume of substance

**Density = mass

volume

Example 2.4

Suppose a student finds that 23.50 mL, of a certain liquid weighs 35.062 g. What is the density of the liquid?

Example 2.5

At a local pawn shop a student finds a medallion that he shop owner insists is pure platinum. However, the student suspects that the medallion may actually be silver and thus much less valuable. The student buys the medallion only after the shop owner agrees to refund the price if the medallion is returned within two days. The student, a chemistry major, then takes the medallion to her lab and measures its density as follows. She first weighs the medallion and finds its mass to be 55.64g. She then places some water in a graduated cylinder and read the volume to be 75.2 mL. Next she drops the medallion into the cylinder and reads the new volume as 77.8 mL. Is the medallion platinum? (density 21.4 g/cm3 ) Or silver? (density 10.5 g/cm3 )

Example 2.6

Mercury has a density of 13.6/mL. What volume of mercury must be taken to obtain 225 g of the metal?

Example 2.7

Racing cars at the Indianapolis Motor Speedway now routinely travel around the track at an average speed of 225 miles per hour. What is the speed in kilometers per hour?

Uncertainty in Measurement

**A) Estimation (in ALL MEASUREMENTS) **

- some instruments have more digits involved(recorded)

- every measurement has 1 AND ONLY 1 uncertain digit, ALWAYS the digit on the right

- having more digits does not make it correct!!!

B) Precision and Accuracy

- terms used to describe the reliability of measurements.

- Precision - us the reproducibility or repeated measurements of the same thing.

-Accuracy - is the closeness of the measured value to the true or accepted value

Mistakes in precision or accuracy leads to ERRORS

2 Types of Errors

Indeterminate or Random Error- uncontrolled errors that result when your repeated measurement(s) effect precision NOT accuracy. Averaging out

Determinate Errors - systematic errors effecting measurement(s) the same way each time. This type of error effects accuracy NOT precision. Instruments

Significant Figures

Quantitative observations are measurements - Number/digit and Unit

Significant figures are used to interpret uncertainty in a system

--- exact numbers are counted or defined. NO UNCERTAINTY!! THEY ARE NOT A RESULT OF AN ESTIMATION.

Ex. A dozen = 12,

****4 Rules to Follow when using Significant Figures***

1) All NON-Zero digits are counted as SigFigs

2) All leading zeros are NOT counted

3) Captive zeros are counted.

4) Trailing zeros are only counted when a decimal point supersedes it

NO measured numbers are exact!!! There is always some uncertainty because we are estimating off of a scale. (Estimation)

--- with measured numbers - the last right most digit is uncertain; it’s the only one.

Ex. 25.34g, 8mL, 8.0 mL, 8.00mL

25.34 g

___ significant figures

___ certain digits ____ uncertain digits

______ +- _______

8mL

___ significant figures

___ certain digits ____ uncertain digits

______ +- _______

Example 2.8

Give the number of significant figures, certain/uncertain, and ranges for each of the following measurements.

a) A sample of orange juice contains 0.0108g of vitamin C.

b) A forensic chemist in a crime lab weighs a single hair and records its

mass of 0.0050060 g.

c) The distance between two points was found to be 5.030 x 103 ft.

d) In yesterday's bicycle race, 110 riders started but only 60 finished

Example 2.9(p.10)

Using the difference balances, three different students weigh the same object. They report masses below. Give the significant figures, certain/uncertain, and ranges for each measurement.

a) 1.611g

a) 1.60 g

b) 0.001611 kg

Example 2.10

Give the number of significant figures, certain/uncertain digits, and ranges for each of the following measurements.

A) 0.00100 m

B) 2.0800 x 102 L

C) 480 kg

Calculation and Significant Figures

-SF in the result of a calculation depend on the math operation preformed.

1. In X and / the number of significant figures is guided by the number of the lowest SFs. Count SFs in each number. The result of the calculation should only have the same number of SF.

EX.

2. In addition and division the SF is guided by the left most uncertain digit in the numbers used in the calculation. Result will be uncertain in that digit/place.

EX.

Rounding Off

- Concerned only with the digit to the right of the uncertain digit.

IFF that digit is 5 then keep it

EX.

Example 2.11

Without performing the calculations, tell how many significant figures each answer should contain.

A) 5.19 + 1.9 + 0.842

B) 1081-7.25

C) 2.3 x 3.14

D) the total cost of 3 boxes of candy at $2.50 a box

Example 2.12

Carry out the following mathematical operations and give each result to the correct number of significant figures.

A) 5.18 x 0.0208

B) (3.60 x 10-3 ) x 8.123 / 4.3

C) 21 + 13.8 + 130.36

D) 116.8 - 0.33

E) (1.33 x 2.8) + 8.41

Example 2.13

A) 4.56 x 1.4

B) 8.315 / 298

C) 12.11 + 18.0 + 1.013

D) 0.6875 - 0.1

E) (4.031 * 0.08206 * 373.1) / 0.995

F) (2.3232 + 0.2034 – 0.16 )* 4.0 x 103

Homework 1 due.

QUIZ I??

Problem Solving and Dimensional Analysis

Problem: Suppose you have to pick up doughnuts on your way to school. You need 2 dozen for the class, but you stop at a store that only sells by the doughnut.

How many do you get?

Problem: The cost of a jar of pickles is $2.06. You have $10.00. How many jars of pickles can you pick up?

Both problems involve conversions.

Conversion factor- a ratio of the two parts of statement that relates the two units (= 1)

Ex. 2.85 cm = ? in

Ex. 1.12 in = ? cm

Equivalence statement- exactly the same, just a different scale

ONES TO KNOW:

Converting from one step to another

Step 1 To convert from one unit to another, use the equivalence statement that relates the two units. The conversion factor needed is a ratio of the two parts of the equivalence statement.

Step 2 Choose the appropriate factor by looking at the direction of the required change (Make sure the unwanted units cancel).

Step 3 Multiple the quantity to be converted by the conversion factor give the quantity with the desired units.

Step 4 Check that you have the correct number of significant figures.

Step 5 Ask whether your answer makes sense

Example 2.14

An Italian bicycle has its frame size given as 62 cm. What is the frame size in inches?

Example 2.15

The length of the marathon race is approximately 26.2 mi. What is this distance in kilometers?

Example 2.16

Perform the following conversion.

A) 36 ft ( in ( cm

B) 36 in ( ft ( miles

C) 6.25 km( miles

D) 55 min ( hours ( seconds

*** WORKSHEET***

Temperature Conversions: An Approach to Problem Solving

Scales-

Fahrenheit Celsius Kelvin (or Absolute)

- used in US -metric - international

zero point 32o 0o -273.15

boiling 212o 100o 373.15

Converting between scales

*Temperature Conversion Formulas*

3. Celsius to Kelvin TK = ToC + 273

4. Kelvin to Celsius ToC = TK - 273

5. Celsius to Fahrenheit ToF = 1.80(ToC ) + 32

6. Fahrenheit to Celsius ToC = (ToF -32 )/1.80

Example 2.17

Temperature conversion: Celsius to Kelvin

The boiling point of water at the top of Mt. Everest is 70.oC. Convert this temperature to the Kelvin scale. (The decimal point after the temperature reading indicated that the trailing zero is significant.)

**Equation ToC + 273 = TK

Example 2.18

Liquid nitrogen boils at 77K. What is the boiling point of nitrogen on the Celsius scale.

*Equation ToC = TK - 273

Example 2.19

On a summer day the temperature in the laboratory, as measured on a lab thermometer, is 28oC. Express this temperature on the Fahrenheit scale.

**Equation ToF = 1.80(ToC ) + 32

Example 2.20

Express the temperature -40 C on the Fahrenheit scale.

Example 2.21

One of the body's responses to an infection or injury is to elevate its temperature. A certain flu victim has a body temperature of 101 0F. What is this temperature on the Celsius scale?

Example 2.22

Mercury thermometers are being phased out because of the toxicity of mercury vapor. A common replacement for mercury is the organic liquid isoamyl benzoate, which boils at 262 oC. What is its boiling point in oF, K?

1. Types of Matter

Rem: Chemistry is a natural science that deals with the changes that matter undergoes.

What is matter? -Anything that takes up space and has mass

Ex. Elements and Compounds

Element - a substance that cannot be separated into a simpler substances by chemical or physical changes. (Pure substance)

List some common elements

Carbon

Hydrogen

Oxygen

Nitrogen

Phosphorous

Sulfur

Elements are represented by symbols (see Periodic Table) = Chemical Symbols

Elements are the building blocks for compounds

Compound - a chemical combination of two or more elements in constant composition.

ALWAYS contains different elements, ALWAYS have the same composition (same combination)- also a pure substance

Mixture and Pure Substances

Virtually all matter consists of mixtures and substances.

Ex. Soil, Air,

Substance - a generic term referring to a sample of matter with uniform & definite composition.

Mixture - variable composition; 2 or more substances combined physically

Ex. Air, seawater,

2 Types of mixtures

Heterogeneous Homogeneous (solution)

- variable composition - uniform composition

Not uniform uniform properties throughout

Ex. Ex.

* most matter around us occurs naturally as a mixture.

Mixtures can be separated into their components by physical methods.

Ex. sifting sand from water, distillation, filtration, magnets,

Pure Substances- substances with definite and fixed composition.

Ex. Elements, compounds, Pure water

Example 2.23

Identify each of the following as a pure substance, a homogeneous mixture or a heterogeneous mixture.

A. gasoline D. brass

B. a stream with gravel on the bottom E. copper metal

C. air

Example 2.24

Classify each of the following as a pure substance, a homogeneous mixture or heterogeneous mixture.

A. Wine B. the oxygen and helium in a scuba tank

C. Oil and vinegar salad dressing D. Common salt (sodium chloride)

Separation of Mixtures

Mixtures can be separated by physical means.

Distillation-

Distillation Apparatus

Filtration-

Flow Chart

Matter

Pure Substance Homogeneous Mixture Heterogeneous Mixture

Element Compound Pure Substance

States of Matter

Solid -- has definite shape and definite volume

- incompressible

Ex:

Liquid -- has definite volume, but takes the shape of the container

- flows

Ex:

Gas -- no definite shape, no definite volume

Ex:

-Plasma -- compressed gas at high temperatures (sun and other stars)

* temperature and pressure determine the state of matter a substance is in, also forces acting on the particles

stronger forces = more rigid matter

Matter/Mass can NEVER be created nor destroyed!!!!!

Law of the Conservation of Mass

Proof: LaVoisier - experiments with combustion, No Mass lost.

Properties of Matter

Intensive -

Extensive-

Physical and Chemical Properties and Changes

If you see a friend you call them by name: unique characteristics (Physical)

hair, face, height, clothing, voice, glasses, ect.

Matter has properties.

Physical Properties Chemical Properties

- malleable - high or low melting point

- ductile - high or low freezing point

- hard/soft - reactivity

- high or low luster - stability

(odor, color, volume, state, (ability to form new substances)

density, melting point, boiling point

solubililty)

Example 2.25

Classify each of the following as a physical or chemical property

A. The boiling point of a certain alcohol is 78oC

B. Diamond is very hard

C. Sugar ferments to form alcohol

D. A metal wire conducts an electric current

Example 2.26

Which of the following are physical properties and which are chemical properties?

A. Gallium metal melts in your hand

B. Platinum does not react with oxygen at room temperature

C. This page is white

D. The copper sheets that form the "skin" of the Statue of Liberty have acquired a

greenish coating over the years.

Take a look at Water: H2O (the Mickey Mouse molecule)

Three dimensional diagram

What is really occurring when water undergoes the following changes?

Solid -----> Liquid -----> Gas (steam)

Melting Boiling

Matter undergoes changes.

Physical Changes Chemical Changes

Melting, boiling, cutting rusting, growing, decaying, decomposing

does NOT affect the composition - give the substance new properties

Physical Change-involves a change in one or more physical properties, but no change in the fundamental components that make up the substance. THe most common physical changes are changes of state: solid liquid gas.

Chemical Change- involves a change in the fundamental components of the substance; a given substance changes into a different substances. Chemical changes are called reactions: silver tarnishes by reacting with substances in the air; a plant forms a leaf by combining various substances from the air and soil; and so on.

Example 2.27

Classify each of the following as a physical or chemical change.

A. Iron metal is melted

B. Iron combines with oxygen to form rust

C. Wood burns in air

D. A rock is broken into small pieces

Example 2.28

Classify each of the following as a chemical or physical change or a combination of the two.

A. Milk turns sour

B. Wax is melted over a flame and then catches fire and burns.

Energy and Energy Changes

Energy(E)- is the capacity to do work

Ex. elements, gases, compounds, solids, liquids, cars, stoves, bunsen burners

2 Types of Energy(E)

Potential Energy(PE) Kinetic Energy(KE)

- energy due to position - energy due to motion of an object

Ex. Rock on the edge of a tall building Ex. A falling rock, Heat (often = KE)

How can we measure energy? Equations, mathematics

What are the units? Joules (SI system unit)

Joules (J) or calories (cal).

1 cal = amount of heat required to raise 1 g of pure water 1oC

1 cal = 4.184 J

Also Energy can NEVER be created nor destroyed (only converted from one form to another)!!!!! Law of the Conservation of Energy

FYI: 1 cal DOES NOT = calories you eat. Those are kilocalories or (C)

Example 2.29

Express 60.1 cal of energy in units of joules.

Example 2.30

How many calories of energy corresponds to 28.4 J.

Example 2.31

Determine the amount of energy (heat) in joules required to raise the temperature of 7.40g water from 29.0 oC to 46.0 oC. Hint: 4.184 J (specific heat for water)

g oC

Example 2.32

Calculate the joules of energy required to heat 454 g of water from 5.4 C to 98.6 C.

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