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Chapter 1Introduction:

Matter and Measurement

John 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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Matter:

Anything that has mass and takes up space.

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Properties and Changes of

Matter

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Properties of Matter

• Physical Properties:□ Can be observed without changing a

substance into another substance.• Boiling point, density, mass, volume, etc.

• Chemical Properties:□ Can only be observed when a substance is

changed into another substance.• Flammability, corrosiveness, reactivity with

acid, etc.

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Properties of Matter

• Intensive Properties:□ Independent of the amount of the

substance that is present.• Density, boiling point, color, etc.

• Extensive Properties:□ Dependent upon the amount of the

substance 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:□ Changes that result in new substances.

• Combustion, oxidation, decomposition, etc.

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Chemical Reactions

In the course of a chemical reaction, the reacting substances are converted to new substances.

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Units of Measurement

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SI Units• There are two types of units:

– fundamental (or base) units;– derived units.

• There are 7 base units in the SI system.

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SI Units

• Système International d’Unités• Uses a different base unit for each quantity

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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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Volume

• The units for volume are given by (units of length)3.– SI unit for volume is 1

m3.

• We usually use 1 mL = 1 cm3.

• Other volume units:– 1 L = 1 dm3 = 1000 cm3

= 1000 mL.

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Uncertainty in Measurements

Different measuring devices have different uses and different degrees of accuracy.

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Density:

Physical property of a substance

d=mV

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Do now: A graduated cylinder is filled with 15.0 mL of water. An object with a mass of 29.66 g causes the total volume to increase to 23.4 mL. What is the density of the sample?

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Vocabulary review

• Mass : amount of matter in an object. • It is measured with a balance.• Unit in the SI system: grams g

• Weight: a measure of the pull that the gravity exerts over an object.

• If we stay in the same planet (same gravity) is used like the mass.

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Volume

• The space that an object occupies.• Unit of volume in SI Liter = L

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What is density?

• Density is a comparison of how much matter there is in a certain amount of space.

• IT IS AN INTENSIVE PHYSICAL PROPERTY. It is used to identify a substance.

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DENSITY OF WATER

• 1g/ml• That means that a gram of water has a

volume of 1 milliliter or 1 cubic centimeter.

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Which one is more dense?

• Now which one is more dense?

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What is density?

• Density = mass

volume

or• mass ÷ volume.

Units for density: g .

cm3

• Why are these the units for density?

ALWAYS REMEMBER

UNITS!

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Let’s try a density problem together

• Find the density of a yellow rock has a mass of 8 g and a volume of 4 cm3.

• Use your table S to determine what element it could be!

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IMMISCIBLE LIQUIDS

• If you pour together liquids that don’t mix and have different densities, they will form liquid layers.

• Liquids that don’t mix are said to be IMMISCIBLE

• The liquid with the highest density will be on the bottom.

• The liquid with the lowest density will be on the top.

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Liquid Layers• . Which layer has the highest

density?• Which layer has the lowest

density?• Imagine that the liquids have the

following densities: – 10g/cm3. 3g/cm3.– 6g/cm3. 5g/cm3.

• Which number would go with which layer?

• Is any of the liquids water?

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To measure the volume of an object

• If is a regular object measure the dimensions needed and use the formula

• cube= LxWxH• Cylinder = p h r2

• For an irregular object use the water displacement method.

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Liquid Layers – Try with your neighbor

• Which liquid has the highest density?

• Which liquid has the lowest density?

• Which liquid has the middle density?

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Liquid Layers – Try on your own!

• Imagine that the liquids on the right have the following densities:– 15g/cm3 10g/cm3

– 3g/cm3 9g/cm3

– 7g/cm3 12g/cm3

• Match the colors to the correct densities.

3g/cm3

7g/cm3

9g/cm3

10g/cm3

12g/cm3

15g/cm3

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Review

• What is the formula for density?• What happens if you pour together

liquids that have different densities?• Will the liquid on the top have the

highest or lowest density?• Will the liquid on the bottom have the

highest or lowest density?

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Density worksheet answers

• 1=1• 2=2• 3=2 D of Na 0.971 g/mL• 4=4 D of Mg 1.738 g/cm3• 5=2• 6 20g/24.4L = 0.82 g/L

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• HOMOGENEOUS : SAME PROPERTIES THROUGHOUT THE SAMPLE

• HETEROGENEOUS :• DIFFERENT PROPERTIES IN

DIFFERENT PARTS OF THE SAMPLE

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PURE SUBSTANCEShave a constant composition

• ELEMENTS – Made up of same kind of atoms. Could not be decomposed.

• COMPOUNDS – Made up of different kind of atoms CHEMICALLY COMBINED. Can be decomposed.

• Recognizable by formulas!

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Mixtures and Compounds

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Compounds

Compounds can be broken down into more elemental particles.

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Matter

• Atoms are the building blocks of matter.

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Matter

• Atoms are the building blocks of matter.• Each element is made of the same kind of atom.

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Matter

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

elements.

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SEPTEMBER 21

• ELEMENT , COMPOUNDS AND MIXTURES

• REVIEW FOR TEST• PHYSICAL AND CHEMICAL

PROPERTIES• PHYSICAL AND CHEMICAL

CHANGES• DENSITY

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MIXTURES

• Combination of two or more pure substances. Can be separated by physical means. They do not have a fixed composition.

• Can be homogeneous or heterogeneous.

• SOLUTIONS ARE HOMOGENEOUS MIXTURES.

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AQUEOUS SOLUTIONS

The solution is prepared using water as the solvent

(aq) means DISSOLVED IN WATER!!!

Na Cl (s) is a compound

Na Cl (aq) is a mixture!!!!

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Pure Substances

and Mixtures

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Pure Substances and Mixtures• If matter is not uniform throughout, then it is a

heterogeneous mixture.• If matter is uniform throughout, it is

homogeneous.• If homogeneous matter can be separated by

physical means, then the matter is a mixture.• If homogeneous matter cannot be separated

by physical means, then the matter is a pure substance.

• If a pure substance can be decomposed into something else, then the substance is a compound.

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Elements

• If a pure substance cannot be decomposed into something else, then the substance is an element.

• There are 114 elements known.• Each element is given a unique chemical symbol

(one or two letters).• Elements are building blocks of matter.• The earth’s crust consists of 5 main elements. (O,

Si, Al, Fe, Ca)• The human body consists mostly of 3 main

elements. (O, C, H)

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Elements

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Metals, Nonmetals, and Metalloids

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Symbols

• First letter of element in CAPITAL letter• Second or third letter in lower case.• Some elements have symbols different

from the english name • SODIUM Na• POTASSIUM K• CUPPER Cu• LEAD Pb

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• IRON Fe• MERCURY Hg• GOLD Au• SILVER Ag• TIN Sn

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Symbols from Latin Names

Element Symbol Latin nameCopper Cu cuprumGold Au aurumLead Pb plumbumMercury Hg hydrargyrumPotassium K kaliumSilver Ag argentumSodium Na natriumTin Sn stannum

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DIATOMIC ELEMENTS

•H2O2F2Br2I2N2Cl2

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Classification of Matter

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Classification of Matter

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Classification of Matter

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Classification of Matter

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Classification of Matter

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Classification of Matter

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Classification of Matter

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Classification of Matter

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Classification of Matter

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Classification of Matter

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MC ANSWERS

• 1 . A• 2 . A• 3. A• 4. D• 5. A• 6. B• 7. C• 8. D• 9. D

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Chemical Reactions

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Electrolysis of Water

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Separation of Mixtures

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Distillation:

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

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Distillation

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Filtration:

Separates solid substances from liquids and solutions.

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Chromatography:

Separates substances on the basis of differences in solubility in a solvent.

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Uncertainty in Measurement

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Significant Figures

• The term significant figures refers to digits that were measured.

• When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.

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Uncertainty in Measurement• All scientific measures are subject to error.• These errors are reflected in the number of

figures reported for the measurement.

Precision and Accuracy• Measurements that are close to the “correct”

value are accurate.• Measurements that are close to each other

are precise.

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• ACCURACY DEALS WITH THE EXACTESNESS OF THE MEASUREMENT, HOW CLOSE IT IS TO THE , TRUE, ACCEPTED OR STANDARD VALUE

• PRECISSION DEALS WITH REPRODUCIBILITY OF A MEASUREMENT.

• IF SEVERAL MEASUREMENTS GIVE A SIMILAR RESULT IT IS SAID THAT THE MEASUREMENT IS PRECISE

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Accuracy versus Precision

• Accuracy refers to the proximity of a measurement to the true value of a quantity.

• Precision refers to the proximity of several measurements to each other.

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October 5

• UNCERTAINTY IN MEASUREMENT• SIGNIFICANT FIGURES – Rules and

examples• DO NOW • Calculate the density of an object that

has a mass of 10.0 g and a volume of 3.0 mL.

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Measuring Volume by water displacement

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Significant Figures• The number of digits reported in a

measurement reflect the accuracy of the measurement and the precision of the measuring device.

• All the figures known with certainty plus one extra figure (estimated digit) are called significant figures.

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Sig fig in calculations

• In any calculation, the results are reported to the fewest significant figures (for multiplication and division) or fewest decimal places (addition and subtraction).

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Significant Figures

1. All nonzero digits are significant.

2. Zeroes between two significant figures are themselves significant.

3. Zeroes at the beginning of a number are never significant.

4. Zeroes at the end of a number are significant if a decimal point is written in the number or if they are to the right of a decimal point.

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0.0030 has 2 sf

400. Has 3 sf

700.00 has 5 sf

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EXAMPLE FOR ADDITION

• Copy and perform the following operation indicating the right number of sig fig

12

1.2

0.2

---------------------

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MULTIPLICATION AND DIVISION

5 x 100 =

0.00022 x 745 =

3469/ 5799=

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Examples: How many significant figures are in each of the following?

1. 52300 m

2. 0.000487 kg

3. 29.0400 s

4. 507 people

5. 230,050 cm

6. 45.600 A

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Tell the number of significant digits in each of the following measurements. 1. 48 cm __________ 2. 306.2 g __________ 3. 0.329 m __________ 4. 83.9520 °C __________ 5. 3700 mm __________ 6. 400. cm3__________ 7. 71.60 g __________ 8. 0.00432 mm __________ 9. 10.0 kg __________10. 3.60 × 1015 sec __________11. 6.24 × 10-4 m __________12. 82.000 g __________

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Tell the number of significant digits in each of the following measurements. 1. 48 cm __________ 2. 306.2 g __________ 3. 0.329 m __________ 4. 83.9520 °C __________ 5. 3700 mm __________ 6. 400. cm3__________ 7. 71.60 g __________ 8. 0.00432 mm __________ 9. 10.0 kg __________10. 3.60 × 1015 sec __________11. 6.24 × 10-4 m __________12. 82.000 g __________

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• How do scientist express the accuracy of a measurement?

• DO NOW : Observe the two instruments in my desk to measure volume, determine which would determine the volume of an Al cylinder with greater accuracy and explain in your notebook why.

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Percent Error

• To determine the accuracy of a measurement. It tells us how far our measured stands from an accepted or known value.

% error = I measured value – accepted value I X 100

___________________________________________

accepted value

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Example

• Calculate the percent error of the measurement for a student that determined that the density for Al is 2.5 g/mL. Hint use table T to determine the accepted value!

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1) 3.482 cm + 8.51 cm + 16.324 cm ____________________ 2) 48.0032 g + 9.17 g + 65.4321 g ____________________

3) 80.4 cm - 16.532 cm ____________________

4) 106.5mL - 30. mL ____________________ 5) 48.2 cm × 1.6 cm × 2.12 cm ____________________ 6) 8.3 m × 4.0 m × 0.9823 m

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7) 64.34 cm3 ÷ 8.149 cm ____________________ 8) 4.93 mm2 ÷ 18.71 mm ____________________

9) 0 57 mL x 760 mm/740 mm x 273K/250 K

10) 5 13 g x 44 962amu/15 874 a m u

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Answers

1. 28.32 cm

2. 122.61 g

3. 63.9 cm

4. 76 mL

5. 160 cm3

6. 33 m3

7. 7.895 cm2

8. 0.263 mm

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MULTIPLE CHOICE

1. 3

2. 2

3. 3

4. 2

5. 1

6. 1

7. 4

8. 1

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• 1. According to an accepted chemistry reference. the heat of vaporization of water is 540. calories per gram. A student determined in the laboratory that the heat of vaporization of water was 620. calories per gram. The student's results had a percent error of (1) 12.9, (2) 80.0, (3) 14.8, (4) 87.1

• 2. Which measurement contains a total of three significant figures?

• (1) 0.01 g (2) 0.0100 g (3) 0.010 g (4) 0.01000 g

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3. In an experiment the gram atomic mass of magnesium was determined to be 24.7. Compared to the accepted value 24.3, the percent error for this determination was (1) 0. 400, (2) 24.7, (3) 1.65, (4) 98.44. A student determined the melting point of a substance to be 55.2°C. If the accepted value is 50. 1°C the percent error in her determination is (1) 5.10, (2) 10.2, (3) 9.24, (4) 12.05. Using the rules for significant figures, the sum of 0. 027 gram and 0. 0023 gram should be expressed as (1) 0. 029 gram, (2) 0.03 gram, (3) 0.0293 gram, (4) 0.030 gram

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6. Which milligram quantity contains a total of four significantfigures? (1) 0.3010 mg (2) 3100 mg (2) 3010 mg (4) 30001 mg

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Scientific Notation• Numbers written in scientific notation include a

numeral with one digit before the decimal point, multiplied by some power of ten (6.022 x 1023)

• In scientific notation, all digits are significant.

• You should be able to convert from non-scientific notation to scientific and vice-versa.

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• Temperature • Definition• Instrument• Scales

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TEMPERATURE

• Is associated with heat but it is NOT HEAT. IT IS NOT A FORM OF ENERGY!!!! ( Heat is)

• Review: What is KINETIC ENERGY?

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KINETIC ENERGY (KE)

• Is associated with movement. • If an object is moving fast has high KE• If an object is moving slowly it has low KE

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Temperature• In scientific

measurements, the Celsius and Kelvin scales are most often used.

• The Celsius scale is based on the properties of water.□ 0C is the freezing point

of water.□ 100C is the boiling

point of water.

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Temperature:

A measure of the average kinetic energy of the particles in a sample.

If an object is at HIGH temperature its particles are moving FAST

At LOW temperature particles move SLOWLY

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Instrument to measure temperature

THERMOMETER

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FIXED POINTS OF A THERMOMETER

• BOILING POINT OF WATER

• FREEZING POINT OF WATER

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Temperature

• The Kelvin is the SI unit of temperature.

• It is based on the properties of gases.

• There are no negative Kelvin temperatures.

• K = C + 273.15

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Temperature

• The Fahrenheit scale is not used in scientific measurements.

• F = 9/5(C) + 32• C = 5/9(F − 32)

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Examples:

1. What is 35ºC in Kelvin? In ºF?

2. What is 183 K in ºC? In ºF?