Chemistry

Chapter 1 and 2Scientific Method and Measurement

VocabularyVocabulary HeadingsHeadings

Important InfoImportant Info

Distinguish between a scientific law and a scientific theory.

• Scientific Law– Observation of a natural

event– Summary of what occurs– Does not try to explain

why something occurs only tells what occurs

• Scientific Theory– Explanation of events– Tries to explain why

something occurs– Supported by several

experiments– Must be able to predict

what happens by using the theory

Explain and apply the steps of the scientific method.

• Observation– Seeing something that makes you

ask a question• Formulate a question

– What exactly do you want to learn?

• Research the question– Find out what others have said

about your question• Develop a hypothesis

– Use the information you found in research to develop an educated guess in answer to your question

• Experiment/Collect Data– Test only one variable at a time– Use a control

• Control is a version of the experiment where nothing is changed

• Draw conclusions– Analyze the data– Did the data support your

hypothesis?

Describe the relationship between pure science and technology.

• Pure science studies things that may never be useful

• Pure science seeks only to know• Technology is useful• Technology is applied science• Studying far off galaxies is pure science• Creating a vaccine for a disease is technology

Distinguish between an independent and dependent variable.

• Independent variable• The variable that you

change in the experiment• If you place one plant in the

window and one in the closet the variable you are changing is the amount of light

• Amount of light would be the independent variable

• Dependent variable• The variable that changes

because of the change in the independent variable

• The plant in the closet is only 6 cm tall. The plant in the window is 12 cm tall.

• Height of the plant would be the dependent variable.

Identify the SI base units and compare the base units to derived units.

QuantityQuantity UnitUnit AbbreviationAbbreviation

LengthLength MeterMeter mm

MassMass KilogramKilogram kgkg

TimeTime SecondSecond ss

TemperatureTemperature KelvinKelvin KK

Electric currentElectric current AmpereAmpere AA

Amount of Amount of substancesubstance

MoleMole molmol

Luminous Luminous intensityintensity

CandelaCandela cdcd

Identify the SI base units and compare the base units to derived units

• Derived units are made up of more than one base unit

• Examples of derived units : g/cm3, g/mol, m3

• If it is not on the list of base units, it is a derived unit.

Define the main prefixes used in the metric system.

PrefixPrefix SymbolSymbol MeaningMeaning Multiple of Multiple of base unitbase unit

DecaDeca dada TenTen 1010

HectoHecto hh HundredHundred 100100

Kilo-Kilo- kk ThousandThousand 10001000

Mega-Mega- MM MillionMillion 1 000 0001 000 000

Giga-Giga- GG BillionBillion 1 000 000 0001 000 000 000

Define the main prefixes used in the metric system.

PrefixPrefix SymbolSymbol Meaning Meaning Multiple of Multiple of base unitbase unit

Deci-Deci- dd TenthTenth 1/101/10

Centi-Centi- cc HundredtHundredthh

1/1001/100

Milli-Milli- mm ThousandthThousandth 1/10001/1000

Micro-Micro- MillionthMillionth 1/1,000,0001/1,000,000

Nano-Nano- nn billionthbillionth 1/1,000,000,001/1,000,000,0000

Convert between different metric units.

• G _ _ M _ _ k h da (base) d c m _ _ _ _ n • Start to the right of the given• Move the decimal to the right of the unknown• Add zeroes to hold the decimal place• 2km = ____________ cm• Move from the right of the k to the right of the c• 5 spaces to the right, need 5 zeroes to hold the place• 2km = 200,000cm

Convert the following measurements

• 300 mm = _______m• 2 km = ________hm• 0.90 cm = ______mm• 5.67 dm = ______km• 3.6 Gm = ________m• 4.5 g= _________g• 34 kg = ________mg• 45 ms = ________s

• 6.7 nm = ________m• 37 kg = ________cg• 23 ml = ________ kl• 9.7 dam = ______m• 0.0054 cg = _____mg• 0.5 cm = _______mm• 0.68 Mg = _______cg• 1Gm= _________nm

Solve density problems.• Density is the Mass divided by the Volume

– D = M/V• The equation can be rearranged to solve for any of the three variables.

– M = D x V– V = M/D

• Example• A block of aluminum occupies a volume of 15.0 mL and weighs 40.5 g.

What is its density? • M = 40.5 g• V = 15.0 mL• D = M/V D = 40.5g/15.0 mL D = 2.7 g/mL

Density Problems cont.

• What is the weight of the ethyl alcohol that exactly fills a 200.0 mL container? The density of ethyl alcohol is 0.789 g/mL.

• D = 0.789 g/ml (3 sig figs)• V = 200.0 mL (4 sig figs)• M = D x V M = 0.789g/ml x 200.0 mL• M = 158 g (3 sig figs)

Density Problems cont.

• What volume of silver metal will weigh exactly 2500.0 g. The density of silver is 10.5 g/cm3.

• D = 10.5 g/cm3

• M = 2500.0 g• V = M / D V = 2500.0 g/ 10.5 g/cm3

• V = 238 cm3 (3 sig figs)

Density Problems cont.• A rectangular block of copper metal weighs 1896 g. The dimensions of the

block are 8.4 cm by 5.5 cm by 4.6 cm. From this data, what is the density of copper?

• First calculate the volume. (V = l x w x h)• V = 8.4 cm x 5.5 cm x 4.6 cm V = 212.52 cm3

• There are only 2 sig figs in the problem so you must round to 210 cm3

• Now calculate the density.• M = 1896 g• V = 210 cm3

• D = M/V D = 1896 g/ 210 cm3 D = 9.03 g/cm3

• Can only have 2 sig figs therefore the answer would be 9.0 g/cm3

Density Practice

1. Mercury metal is poured into a graduated cylinder that holds exactly 22.5 mL. The mercury used to fill the cylinder weighs 306.0 g. From this information, calculate the density of mercury.

2. A flask that weighs 345.8 g is filled with 225 mL of carbon tetrachloride. The weight of the flask and carbon tetrachloride is found to be 703.55 g. From this information, calculate the density of carbon tetrachloride.

3. Calculate the density of sulfuric acid if 35.4 mL of the acid weighs 65.14 g.

4. Find the mass of 250.0 mL of benzene. The density of benzene is 0.8765 g/mL.

5. A block of lead has dimensions of 4.50 cm by 5.20 cm by 6.00 cm. The block weighs 1587 g. From this information, calculate the density of lead.

6. What is the volume of a substance with a mass of 0.35 g and a density of 0.9 g/ml?

Use scientific notation to represent very large and small numbers

3.75 x 108

• Scientific Notation • only one number allowed in front of the decimal• 8 is called the exponent• If the exponent is positive, move the decimal to the right the

same number of times as the exponent• If the exponent is negative, move the decimal to the left the

same number of times as the exponent• 375,000,000 is the number above written in standard format

Use scientific notation to represent very large and small numbers

• 300 000 000• Only 1 number allowed in

front of the decimal• Count the number of times

the decimal must be moved this number becomes the exponent

• If the original number is larger than one the exponent is positive

• 3.0 x 108

• 0.000 000 03• If the number is smaller

than one the exponent is negative

• 3.0 x 10-8

Write the following numbers in scientific notation

• 800 000 000 m• 0.0015 kg• 60 200 L• 0.00095 m• 8 002 000 km• 0.000 000 000 06 kg• 602 000 000 000 000 000 000 000 atoms

Write the following numbers in standard form

• 4.5 x 105

• 7.009 x 109

• 4.6 x 104

• 3.2 x 1015

• 3.115 x 10-8 • 6.05 x 10-3

• 1.99 x 10-10

• 3.01 x 10-6

Distinguish between accuracy and precision.

• Precision is how close two measurements are to each other.

• 1.45 and 1.44 are precise

• Accuracy is how close a number is to an accepted value.

• If the accepted value is 9 then the above numbers are not accurate.

Accuracy and Precision

• Numbers may be:• Accurate only

– At least one of the numbers is close to the accepted values, but not close to the other measurements

• Precise only– The numbers are close to one

another but not to the accepted value

• Accurate and precise– The data is close to one

another and close to the accepted value

• Neither accurate nor precise– The data is not close to the

other measurements nor to the accepted value

Precision or Accuracy?

Define significant figures and know when to use them.

• When numbers are measured, measurements are always taken to the first number that is estimated (guessed).

• This is the last significant figure.• If the measurement is 100, the actual number could

be anywhere from 50-149.• If the measurement is 100.0 then it can only vary

from 99.5 and 100.4.• Significant figures indicate precision of

measurement.

Use significant figures in problem solving.

• If the decimal is present come in from the pacific to the first nonzero number (1-9) and count all remaining numbers left to right.

Use significant figures in problem solving.

• If the decimal is absent come in from the Atlantic to the first nonzero number (1-9) and count all remaining numbers from right to left.

Use significant figures in problem solving.• The answer to a problem cannot have more significant figures

than the number in the problem with the fewest significant figures.

• Once the correct number of significant figures has been reached, zeroes are used as place holders

• 76543210 written in 3 significant figures would be 76500000• Cannot drop the zeros. You wouldn’t want me to pay you 10

dollars if I owed you 1000. Dropping zeroes changes the value of the number. Round then hold the place with zeroes to keep the value the same.

• If the number following the last significant figure is 5 or greater the last significant figure goes up one.

• If the number following the last significant figure is 4 or less the last significant figure stays the same.

Determine the number of significant figures in each of the following:

• 1.560• 1560• 0.01560• 300000• 290100000• 0.000002390• 0.000000000002• 14.9800• 100

• 100.0• 20000.0• 3009000

Round the following to 3 significant figures:

• 4.900• 4.905• 20087• 653456• 928227• 5.596• 300.0 (try scientific notation)

Lab Equipment

Erlenmeyer Flask Graduated Cylinder Triple Beam Balance

Watch glassBeakers

Volumetric Flask

Lab Equipment

Test Tubes Test Tube Rack Test Tube Holder

CruciblesCrucible Tongs

Bunsen Burner