SCIENTIFIC METHOD - Ms. Tabors Classroom · Web viewis a measure of how closely numerical values of a set of measurements of the same quantity made in the same way agree. (How close

Name: _______________________Chemistry 1 Notes,

Topic 1 – MeasurementI. Scientific Method / Measurement

Scientific Method – a logical process used to solve a problem

The Scientific Method has five steps:1) Define the problem (I wonder what causes that…?)2) Form Hypothesis (I think it might be…)3) Test Hypothesis (Experiment) (Let’s find out)4) Collect and Analyze Data (Here’s what I found out)5) Draw Conclusions (Here’s what it means)

A hypothesis = a testable statement (this is a better definition than an “educated guess”)

In an experiment, you get data, or information. This info can be one of two types: quantitative or qualitative.

numerical (quantitative) vs. non-numerical (qualitative) info examples:

quantitative qualitative

II. SI Measurement Fundamental units (standards) for SI:

length: meter (m) mass: kilogram (kg)

volume: liter (l) time: second (s)

temperature: Kelvin (K)

Conversions you should know: 2.54 cm = 1 inch 1.6 km = 1 mile 1 kg = 2.2 pounds 3.785 liters = 1 gallon 1 cm3 = 1 mL 1 dm3 = 1 L 0C = 273K

(but 1C increments = 1K increments) 0C = 32F

F = 9/5(C) +32; C = 5/9(F – 32)

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SI prefixes:

Prefix Symbol Meaningmega M 1 000 000kilo k 1000

hecto h 100deca da 10---- ---- ----deci d 1/10centi c 1/100milli m 1/1 000micro m 1/1 000 000

III. Accuracy and Precision Accuracy is a measure of how close a measurement is to the true or accepted value.

(How close is your answer to what you know to be the correct answer?)

Precision is a measure of how closely numerical values of a set of measurements of the same quantity made in the same way agree. (How close are the answers to each other?)

Think of it like a dartboard…

A second example: Say you’re doing an experiment to test the freezing point of water. You take two sets of measurements & get readings of:

1) 36, 37, 36F

2) 31, 32, 32F

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K H D – D C M i e e e e i l c c c n l o t a i t l

o i i

K H D – D C M i e i r h i n n e i o l g r d n c k

y k o i l n a g t

e


Measurements can never be completely exact; there is always some degree of uncertainty in all measurements.

If you’ve ever weighed anything on a digital scale (like yourself), you know what I mean: you step on the scale and get a measurement, but it fluctuates a little bit.

Let’s say the scale reads 172.4 pounds The last digit on any digital scale (the one that fluctuates) is the “uncertain” digit. Your actual weight is within 0.1 pounds of 172.4 So you indicate the uncertainty by stating that your weight is 172.4 0.1 pounds.

(Remember, though, that the uncertainty is not always 1 of the last decimal places / significant figures of the unit you are measuring. Often, the uncertainty is recorded on glassware that you use, or you can look it up. For example, the uncertainty on a graduated cylinder may be 0.2 mL. Especially on glassware, uncertainty may also be called tolerance.)

Remember that uncertainty is an indicator of precision.

Percent error and percent difference, by contrast, are measures of accuracy.

% error =

Observed−ExpectedExpected x 100

% difference =

Observed−Expected Expected x 100

We will typically use % error in our labs, because I want to know if you are above or below the value you were expecting. So, for example, if you were measuring the freezing point of water, and you recorded it as being 34F, then your percent error would be:

3


IV. Significant Figures

SIGNIFICANT FIGURES IN A MEASUREMENT

The number of significant figures in a measurement are all the digits you can measure or read for certain, plus one more that you estimate.

Consider the liquid in this graduated cylinder and in the thermometer…

SIGNIFICANT FIGURES IN A NUMBER

♦ To ascertain how many sig figs are in a given number, there are basically three rules. First ask the question, “Is the # smaller or larger than 1?”

If smaller than 1: Find the first non-zero digit. That digit, and every digit after it are significant. Every zero before it is NOT significant (they are just place holders).

If larger than 1: If there IS a decimal, then every digit is significant. If there IS NOT a decimal, find the last non-zero digit. That digit and every digit before it is

significant. Every zero after it is NOT significant (they are place holders).

ROUNDING

We’ll just use the basic rounding rules you learned in elementary school. When rounding, look at the number just to the right of (after) the last significant digit. If it is less than 5, leave the last sig fig unchanged. (“round down” or “round off”) If it is 5 or higher, increase the last sig fig by one. (“round up”)

SIGNIFICANT FIGURES IN A CALCULATION

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Addition & Subtraction with Sig Figs

First line up the decimals of the numbers, then do the math (don’t round yet). Your final answer can only be as precise as the least precise number in the problem. Round off to

same digit / decimal place as the least precise number in the problem (i.e. – the number ending furthest to the left).

Multiplication & Division with Sig Figs

Again, your answer can only be as precise as the least precise number in the problem. Your answer should have the same # of sig figs as the number in the problem with the fewest sig figs.

12.4 cm x 3.2 cm =

V. Scientific Notation

one digit to the left of the decimal coefficient should indicate the correct # of sig figs

coefficient x 10exponent

exponent = # of places you moved the decimal

if you had to move the decimal to the right, the exponent is negative (negative exponent = # smaller than one)

if you had to move the decimal to the left, the exponent is positive (positive exponent = # larger than one)

problems:

1. 74 000 4. 5.3 x 104

2. 0.000 050 5. 7.00 x 10–3

3. 0.000 602 6. 4.077 x 105

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3 000 000 + 4 327 862 = ?

3 000 000 + 4 327 862 7 327 862

=

156.74 – 38 = ?

156.74 – 38

118.74

=

2 . 50 g0 .04 cm3


VI. Density

Density: D =

MassVolume or

MV (often expressed in units of g/ml or g/L)

A piece of aluminum has a mass of 10.8 g and measures 2.37 cm long, 9.1 mm wide, and 0.0156 m deep. What is its density, in g/cm3?

♦ You need to be able to use your EOC Data Booklets in conjunction with density problems:

An unknown metallic cube has sides which are each 2.0cm long, and a mass of approximately 36.42g. What is the identity of this metal?

What would be the approximate mass of a 17.5cm3 crystal of sodium chloride (NaCl)?

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VII. Temperature Temperature:

The Celsius scale is set up with 0C = freezing point of water & 100C = boiling point of water (in other words, it’s relative to the freezing and boiling points of water.)

The Kelvin scale is set up with 0 K = absolute zero (the temperature at which all molecular motion ceases).

0C = 273 K So, to convert from C K, add: C + 273 = K.

Temperature is the average kinetic energy of the particles in a substance.

Heat is the total kinetic energy of the particles in a substance.

To distinguish: Imagine a bathtub filled with room temperature water and a cup filled with boiling water. The

water in the cup has a higher temperature (on average, the water molecules in the cup are moving faster). But the bathtub has more heat (because there’s much more water).

♦ You also need to be able to use the EOC Data Booklet for problems involving temperature:

You are presented with an organic compound which melts at roughly 89 K. What is the most likely identity of this compound?

VIII. Physical vs. Chemical Changes

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Properties of Matter

Matter has 2 kinds of properties: physical and chemical.

A physical property can be observed or measured without changing the chemical composition of the material (mass, length, volume, melting point, density, color, etc.). Similarly, a physical change doesn’t alter the chemical identity of the matter (if it

still the same formula after the change, it’s a physical change). Examples include: cutting a wire, grinding metal, allowing gas to expand, melting ice, etc.

Changes of state are all physical changes.

A chemical property refers to the ability of a substance to undergo a change that alters its chemical composition. During a chemical change (or reaction) the substance is converted into something new with a different chemical formula and properties. Examples include: flammability (wood tends to burn in air), oxidation (iron rusts),

decomposition (leaves change color in the fall).

The substances that undergo the chemical reactions are called reactants. The substances that are produced by this change are called products. sodium + chlorine sodium chloride

Na + Cl2 NaCl (reactants) (products)

You need to be able to identify and predict these four indicators of a chemical change:

1) formation of a precipitate (a solid) in solution

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Pb(NO3)2(aq) + KI(aq) PbI2(s) + KNO3(aq)

In this reaction, shown to the right, the yellow precipitate is PbI2(s).

(You will need to use the solubility rules on p.6 of your data booklet to be able to predict which product will be the precipitate.)

2) evolution of a gas (look for bubbling, as evidence) Zn(s) + HCl(aq) ZnCl2(aq) + H2(g)

3) color change This often goes hand in hand with formation of a

precipitate (two clear solutions are mixed together, producing a precipitate that is white or some other color).

Color change can also indicate an oxidation-reduction reaction (such as the one shown here: Cu(s) + AgNO3(aq) Cu(NO3)2(aq) + Ag(s)

However – don’t mistake color change due to dilution for evidence of a chemical reaction. It’s still the same color (because it’s still the same ion), just paler, because you’ve added more water.

4) absorption or release of heat Exothermic reactions release heat.

(Their temperature goes up and they feel hot to the touch.) Endothermic reactions absorb heat.

(Their temperature goes down and they feel cold to the touch.)

(You could add emission of light as a 5th indicator, but this is less common than the others.)

♦ Let’s discuss another set of tests you can perform to identify the products of certain reactions.

What a gas is produced in a reaction, you can sometimes perform a burning splint test to determine the identity of the gas. This test is especially useful if the gaseous product is oxygen,

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hydrogen or carbon dioxide. In the burning splint test, you use a match to light a longer wooden splint. You would have already conducted the reaction, and collected the gas produced (often in an upside-down test tube, or something similar). Then you insert the glowing splint into the test tube and observe the results. If the gas collected was: oxygen – the splint would glow brighter, because O2 supports combustion. hydrogen – the splint would ignite the hydrogen, because H2 is flammable. carbon dioxide – the splint would be extinguished, because CO2 does not support

combustion (it would displace oxygen from around the flame)

oxygen (O2) hydrogen (H2) carbon dioxide (CO2)

Another test to determine the presence of carbon dioxide is the lime water test. Lime water is a saturated solution of calcium hydroxide (Ca(OH)2). When carbon dioxide is bubbled through a solution of lime water, it reacts

with the calcium hydroxide to form calcium carbonate and water:

Ca(OH)2(aq) + CO2(g) CaCO3(s) + H2O(l)

Calcium carbonate (limestone or chalk) is insoluble in water, and precipitates out as a white solid. This indicates that the gas was CO2 and not some other gas.

Cobalt (II) chloride is a good indicator for the presence of water (or water vapor). Anhydrous CoCl2 is blue in color. When anhydrous CoCl2 is exposed to water, it absorbs six

water molecules, forming cobalt (II) chloride hexahydrate, [Co(H2O)6]Cl2 or CoCl2·6H2O, which is pink in color.

Often cobalt chloride paper is used in this test. Cobalt chloride paper is simply paper soaked in anhydrous CoCl2. The paper is blue, & turns pink when exposed to water.

CoCl2 + 6H2O [Co(H2O)6]Cl2 (or CoCl2·6H2O) blue pink

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Anhydrous Cobalt Chloride CoCl2 Hydrated Cobalt Chloride

[Co(H2O)6]Cl2 or CoCl2·6H2O

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SCIENTIFIC METHOD - Ms. Tabors Classroom · Web viewis a measure of how closely numerical values of a set of measurements of the same quantity made in the same way agree. (How close