chapter 1 matter and energy - bakersfield college1 1 chapter 1 matter and energy ... 1 - 5 pure...

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Chapter 1Matter and Energy

• Matter and its Classification

• Physical and Chemical Changes and Properties of Matter

• Energy and Energy Changes

• Scientific Inquiry

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Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Classifying Matter – An Exercise

• Look around your chemistry classroom. What objects do you see that are related to chemistry?

• If you were asked to classify these objects, what categories might you use to group similar objects together?

• Why group these similar items together? What makes them similar? What makes other objects different?

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Chemical Classifications

of Matter

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Matter – anything

that occupies space

and has mass

Pure Substances –

have uniform (the

same) chemical

composition

throughout and

from sample to

sample

Mixtures – are

composed of two

or more pure

substances and

may or may not

have uniform

composition

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Pure Substances

• Have uniform, or the same, chemical composition throughout and from sample to sample.

• Two kinds of pure substances– Elements

• An element is a substance that cannot be broken down into simpler substances even by a chemical reaction.

• Elements are separated further into metals and nonmetals.

– Compounds

• A compound is a substance composed of two or more elements combined in definite proportions.

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Pure Substances

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PureSubstances

Elements Compounds

Metals Nonmetals

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Classifying Metals and NonmetalsWhich of the following are metals? Which are nonmetals?1. Phosphorus – P

2. Gold – Au

3. Sulfur – S

4. Copper - Cu

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Figure 1.5

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Classifying Metals and NonmetalsWhich of the following are metals? Which are nonmetals?

1. Phosphorus - P 5. Bromine - Br 9. Lead - Pb

2. Gold - Au 6. Aluminum - Al 10. Tin - Sn

3. Carbon - C 7. Sulfur – S

4. Copper - Cu 8. Nickel - Ni

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Figure 1.5

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Mixtures

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Mixtures –consist of 2 or

more pure substances

HomogeneousMixtures

(solutions) -have uniform compositionthroughout

HeterogeneousMixtures – do

not have uniform

compositionthroughout

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Classifications of Matter

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Page 9

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Representations of Matter

• Matter is composed of atoms.

– An atom is the smallest unit of

an element that retains the chemical

properties of that element.

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Representations of Matter• Atoms can combine

together to form molecules.

– Molecules are composed of two or more atoms

bound together in a

discrete arrangement.

– The atoms bound together in a molecule can be from the same element or from different elements.

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Figure 1.11

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Practice - Representations of Matter

Identify the nonmetals in Figure 1.4. Explain the characteristics you considered in making your decision.

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Practice Solutions -Representations of Matter

Metals can be distinguished from nonmetals by the luster and ability to conduct electricity. Since we do not know how each of elements in Figure 1.4 conduct electricity, we need to use luster as our measure. Nonmetals are usually dull, with the exception of carbon (as diamond). Elements that are gases at room temperature are also nonmetals. Therefore, the nonmetals in Figure 1.4 are phosphorus, carbon, bromine, and sulfur.

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States of Matter

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• Discuss the following questions with someone near you.

– How does a solid differ from a liquid?

– How does a gas differ from a liquid?

– How does a solid differ from a gas?

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Table 1.2 States of Matter

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Particles are widely separated and move independently of one

another

Particles are randomly arranged and free to move

about until they bump into one another

Particles are fixed in place in a regular

array

Large volume change under pressure

Slight volume change under pressure

No volume change under pressure

Volume of containerIts own volumeIts own volume

Shape of container (fills it)

Shape of container (may or may not fill it)

Fixed shape

GasLiquidSolid

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States of Matter

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States of Matter

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Figure 1.15

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States of Matter

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Symbols Used in Chemistry• Elemental symbols

– a shorthand version of an element’s longer name

– can be 1-2 letters and can be derived from the Latin or Greek name [ex. Ag]

• Chemical formulas

– describe the composition of a compound

– use the symbols for the elements in that compound [ex. H2O and CO2]

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Symbols Used in Chemistry

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Ffluorine

CO2

carbon dioxide

H2Owater

CH4

methane (natural gas)

Agsilver

Hehelium

SymbolName

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Symbols Used in Chemistry

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Symbols Used in Chemistry• Symbols for physical

states

– are found in parenthesis by the

elemental symbol or chemical formula

– designate the physical state

[ex. solid, liquid, gas, aqueous]

– Also see Table 1.3

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(aq)aqueous

(dissolved in water)

(g)gas

(l)liquid

(s)solid

SymbolName

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Physical Properties of Matter

• Physical properties

– are properties that can be observed

without changing the composition of the substance

– Four common physical properties are:

• mass

• volume

• density

• temperature

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Mass• Mass:

– measures the quantity of matter

– is essentially the same physical quantity as weight, with the exception that weight is bound by gravity, mass is not

– common units are grams (g)

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Volume

• Volume:

– amount of space a substance occupies

– can be calculated by measuring the sides of a cube or rectangular side, then multiplying them

Volume = length * width * height

– common units are centimeters cubed (cm3) or milliliters (mL)

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Figure 1.18

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Density• Density:

– the ratio of the mass to its volume

density = mass

volume

– units are g/mL (solids and liquids) or g/L (gases)

– See Table 1.4 for a listing of densities for common substances

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Temperature• Temperature:

– a measure of how hot or cold something is relative to some standard

– is measured with a thermometer

– at which a phase change occurs is independent of sample size

– units are degrees Celsius (°C) and degrees Kelvin (K)

TK = T°C + 273.15

T°F = 1.8(T°C) + 321 -

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Physical Changes

• A physical change

– is a process that changes the physical properties of a substance without

changing its chemical composition

– evidence of a physical change includes:

• a change of state

– Example: water changes from a solid to a gas

• an expected change in color

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Chemical Changes• A chemical change

– is a process in which one or more substances are converted into one or more new substances

– also called a chemical reaction

– evidence of a chemical change includes:

• bubbling

• a permanent color change

• a sudden change in temperature

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Chemical Properties

• Defined by what it is composed of and what chemical changes it can undergo

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Practice - Physical vs. Chemical Changes

• Classify each of the following as a physical or chemical change:

– Evaporation of water

– Burning of natural gas

– Melting a metal

– Converting H2 and O2 to H2O

– Boiling an egg

– Crushing rocks

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Practice Solutions - Physical vs. Chemical Changes

• Classify each of the following as a physical or chemical change:– Evaporation of water �� physical change

– Burning of natural gas �� chemical change

– Melting a metal �� physical change

– Converting H2 and O2 to H2O �� chemical change

– Boiling an Egg �� chemical change

– Crushing Rock �� physical change

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Energy and Energy Changes• Energy

– is the capacity to do work or to transfer heat

• Two main forms of energy are:

– Kinetic energy: the energy of motion

– Potential energy: energy possessed by an object because of its position

• Other energies are forms of kinetic and potential energy (chemical, mechanical, electrical, heat, etc.)

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Energy and Energy Changes

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Energy Units• Units for energy are calories or joules

4.184 J = 1 cal

– A calorie is the amount of energy required to raise 1 g of water by 1°C.

– A kilojoule (kJ) or 1000 joules (J) is approximately the amount of energy that is emitted when a kitchen match burns completely.

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Energy Units

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The Scientific Method

• is an approach to asking

questions and seeking answers

that employs a variety of tools,

techniques, and strategies

• The method generally includes observations, hypotheses, laws, and theories.

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• Observations include:

– experimentation

– collection of data

• A hypothesis is a tentative explanation for the properties or behavior of matter that accounts

for a set of observations and can be tested.

• A scientific law describes the way nature

operates under a specified set of conditions.

• Theories explain why observations, hypotheses,

or laws apply under many different circumstances.

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Problem

Propose a Problem

Propose a MethodWhat Data would you need?

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Scientific NotationMath Toolbox 1.1

• A number written in scientific notation is expressed as:

C x 10n

where C is the coefficient (a number equal to or greater than 1 and less than 10) and n is the exponent (a positive or negative integer)– C is obtained by moving the decimal point to

immediate right of the leftmost nonzero digit in the number

– n is equal to the number of places moved to obtain C• If the decimal point is moved to the left, then n is

positive.• If the decimal point is moved to the right, then n is

negative.

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2.45 x 100Neither2.45

8.8 x 101Left88.

5.0607 x 10-3Right0.0050607

3.245 x 109Left3,245,000,000.

3.245 x 10-6Right0.000003245

3.245 x 103Left3245.

Scientific Notation

Direction Decimal Moved

Normal Notation

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Scientific Notation – Using Your Calculator

Math Toolbox 1.1The general approach to enter numbers expressed

in scientific notation into your calculator is:

1. Enter the coefficient, including the decimal point.

2. Press the EE or EXP (depending on your calculator model) to express the exponent. This button (EE or EXP) basically stands for “x 10—”.

3. Enter the exponent, using the change sign (+/- or (-)) to express negative exponents if needed.

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• You will use your calculator to perform mathematical operations in chemistry. Use these simple rules for mathematical operations involving exponents to confirm that your answer makes sense.

• Three operations are most common:

– Multiplication

(4.0 x 106)(1.5 x 10-3) = (4.0 x 1.5) x 106+(-3) = 6.0 x 103

When multiplying numbers in scientific notation, multiply the coefficients and add the exponents.

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– Division

(4.0 x 10-2)/(2.0 x 103) = (4.0/2.0) x 10(-2)-3 = 2.0 x 10-5

When dividing numbers in scientific notation, divide the coefficients and subtract the exponents.

– Raising the exponent to a power

(3.0 x 104)2 = (3.0)2 x 104x2 = 9.0 x 108

When raising numbers in scientific notation to a power, raise the coefficient to that power, then multiply the exponents.

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Practice - Math Toolbox 1.1

• Fill in the above table.• Multiplication and Division of Exponents

1. (6.78 x 103)(5.55 x 10-4) = ___________________2. (2.99 x 10-9) / (4.03 x 10-6) = _________________3. (7 x 103)4 = _______________________________

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14.82

0.00000834

299,800,000.

1.21

9,000,000,655.00

Scientific NotationNormal Notation

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Practice Solutions - Math Toolbox 1.1

1. (6.78 x 103)(5.55 x 10-4) = 3.76 x 100

2. (2.99 x 10-9) / (4.03 x 10-6) = 1.20 x 10-14

3. (7 x 103)4 = 2 x 1015

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6.3 x 10163

1.482 x 10114.82

8.34 x 10-60.00000834

2.99800000 x 108299,800,000.

1.21 x 1001.21

9.00000065500 x 1099,000,000,655.00

Scientific NotationNormal Notation

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Precision vs. AccuracyMath Toolbox 1.2

• The precision of a measured number is

– the extent of the agreement between repeated measurements of its value.

– If repeated measurements are close in value, then the number is precise, but not necessarily accurate.

• Accuracy is

– the difference between the value of a measured number and its expected or correct value.

– The number is accurate if it is close to its true value (much like hitting a bulls-eye on a dart board).

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Precision vs. AccuracyMath Toolbox 1.2

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Significant FiguresMath Toolbox 1.2

We need rules to determine the number of

significant figures in a given measurement.

1. All nonzero digits are significant.

Ex. 9.876 cm (4 significant figures)

2. Zeros to the right of the leftmost nonzero digit (often called leading zeros) are not significant.


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Significant FiguresMath Toolbox 1.2 - Continued

We need rules to determine the number of

significant figures in a given measurement.

3. Zeros to the left of the rightmost nonzero digit (typically called trailing zeros) are significant.


– Note: a decimal point is mandatory for #3 to be true. Otherwise the number of significant figures is ambiguous.

Ex. 98,760 cm (?? significant figures)

4. Zeros between two nonzero digits are significant.


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Significant FiguresMath Toolbox 1.2 - Continued

• When performing a calculation, your final answer must reflect the number of significant figures in the least accurate (most uncertain) measurement.

• The least accurate measurement is expressed differently depending on the mathematical operation you are performing.

– Multiplication and Division

234.506 cm

4455.9 cm

x 0.12 cm

1.3 x 105 cm

The least accurate measurement in a multiplication or division problem is the one with the smallest number of significant digits overall. Therefore, because the 3rd

measurement has 2 significant digits overall, the final answer must have 2 significant digits overall. 1 -

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Significant Figures Math Toolbox 1.2 - Continued

• When performing a calculation, your final answer must reflect the number of significant figures in the least accurate (most uncertain) measurement. – Addition and Subtraction

234.5 06 cm0.1 2 cm

+ 4455.9 cm4690.5 26 cm

The least accurate measurement in an addition or subtraction problem is the one with the smallest number of significant figures to the right of the decimal point. Therefore, because the 3rd

measurement has 1 significant digit to the right of the decimal point, the answer must have 1 significant digit to the right of the decimal point.

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• Fill in the above table.

• Calculate the following: 1. 14.6608 + 12.2 + (1.500000 x 102) = ____________________

2. (5.5 x 10-8)(4 x 1010) = _______________________________

6.65 x 1045

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1.2407661 x 10-2

9 x 1019

0.003416

19628.00

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# of significant digits Given number

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Practice Solutions –Math Toolbox 1.2

1. 14.6608 + 12.2 + (1.500000 x 102) = 176.9

2. (5.5 x 10-8)(4 x 1010) = 3 x 10-43

6.65 x 1045

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81.2407661 x 10-2

19 x 1019

40.003416

719628.00

226

# of significant digits Given number

19

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Units and ConversionsMath Toolbox 1.3

• Measurement– is the determination

of the size of a particular quantity

– Measurements are defined by both a quantity (number) and unit.

– Most scientists use SI (from the French for SystèmeInternationale) units (see top chart or chart on pg. 39).

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amount of substance

molmole

temperatureKkelvin

electric current

Aampere

timessecond

masskgkilogram

lengthmmeter

QuantitySymbolUnit

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Units and Conversions - Continued Math Toolbox 1.3

• Scientists also use the metric system to define base units of measure, with the understanding that a special prefix denotes fractions or multiples of that base (see chart on left).

1 -n10-9nano

µ10-6micro

m10-3milli

c10-2centi

d10-1deci

k103kilo

M106mega

G109giga

SymbolFactorPrefix

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Unit AnalysisMath Toolbox 1.3

A possible approach to problem solving involves 4 steps:

1. Decide what the problem is asking for.

2. Decide what relationships exist between the information given in the problem and the desired quantity.

3. Set up the problem logically, using the relationships decided upon in step 2.

4. Check the answer to make sure it makes sense, both in magnitude and units.

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Unit AnalysisMath Toolbox 1.3

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1. How many inches are in 2 kilometers?

[1 in = 2.54 cm; 100 cm = 1 m; 1000 m = 1 km]

2. What is the volume of a 14 lb block of gold?

[1 lb = 453.6 g; dAu = 19.3 g/cm3]

3. Dan regularly runs a 5-minute mile. How fast is Dan running in feet per second?

[1 min = 60 s; 1 mile = 1760 yds; 1 yd = 3 ft]

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1. How many inches are in 2 kilometers?

[1 in = 2.54 cm; 100 cm = 1 m; 1000 m = 1 km]

First, decide what the problem is asking for: inches.

Next, decide what relationships exist between the information given and the desired quantity: starting with 2 km, we can use the conversion factors (in the brackets) to convert from 2 km to inches.

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in 10 x 8 cm 2.54

in 1 x

m 1

cm 100 x

km 1

m 1000 x km 2 4

=

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2. What is the volume of a 14 lb block of gold?

[1 lb = 453.6 g; dAu = 19.3 g/cm3]

First, decide what the problem is asking for: volume.

Next, decide what relationships exist between the information given and the desired quantity: starting with 14 lb, we can use the conversion factors (in the brackets) to convert from 14 lb to volume (cm3).

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cm 10 x 3.3 g 19.3

cm 1 x

lb 1

g 453.6 x lbs 14 =

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3. Dan regularly runs a 5-minute mile. How fast is Dan running in feet per second?

[1 min = 60 s; 1 mile = 1760 yds; 1 yd = 3 ft]

First, decide what the problem is asking for: feet per second.

Next, decide what relationships exist between the information given and the desired quantity: starting with 5 minutes per mile, we can use the conversion factors (in the brackets) to convert from miles per minute to feet per second.

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ft/s 10 x 2 yd 1

ft 3 x

mile 1

yards 1760 x

s 60

minute 1 x

minutes 5

mile 1 1=

chapter 1 matter and energy - bakersfield college1 1 chapter 1 matter and energy ... 1 - 5 pure...

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