Lab Equipment: Glassware

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September 30, 2014 Lab Equipment: Glassware Beakers are used to hold liquids. Can be used to heat liquids at relatively low temperatures. Erlenmeyer flasks are used to hold liquids. Sloped sides = ideal for swirling mixtures

Transcript of Lab Equipment: Glassware

Page 1: Lab Equipment: Glassware

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Lab Equipment: Glassware

• Beakers are used to hold liquids.• Can be used to heat liquids at

relatively low temperatures.

• Erlenmeyer flasks are used to hold liquids.

• Sloped sides = ideal for swirling mixtures

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Lab Equipment

• Crucibles are used to heat solids at high temperatures.

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Lab Equipment: Pick up tools

• Beaker tongs pick up hot beakers.

• Crucible tongs pick up hot crucibles.

• Test tube clamps pick up hot test tubes.

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Lab Equipment: Measuring volume

• Graduated cylinders accurately measure liquid volumes.

• Always read from the bottom of the meniscus

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Lab Equipment: Measuring mass

• Electronic balance measures mass to .001 g.

• Record ALL of the digits and the units• Make sure to tare the balance before

use.

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Lab Equipment: Heating tools

• Bunsen burner is a high temperature heat source.

• Connect to gas and use a striker to light.

• Hot plate is used to heat at low temperatures.

• It still gets HOT--use caution.

• Temperature probe measures temperatures!

• Read temperature on digital screen.

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Lab Equipment: Small volumes

• Test tubes hold small volumes of liquids.

• Easy to observe contents.• Test tubes go in a test tube rack or

if hot, in a beaker.

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Lab Equipment: Evaporation

• Watch glass is used to slowly evaporate liquids.

• Evaporating dish are used to evaporate or heat liquids or solids at low temperatures.

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Lab Equipment: Miscellaneous

• Rubber stoppers close off test tubes and flasks.

• Some have holes for tubing.

• Scoopula is used to transfer solids

• Stirring rod is used for stirring.

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You will use other lab equipment and they will be introduced along with labs.

Now on to measurements...

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Measurements

• Measurements are quantitative observations.• What are some kinds of quantitative observations you

might make?

Temperature MassLengthVolume

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Student A and Student B measured the same paperclip.

• Student A recorded a measurement of 1.0 inches.• Student B recorded a measurement of 2.54 cm.

• What happened?

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Length

Mass

Time

Amount of substance

Temperature

Volume

meter (m)

kilogram (kg)

liter (L)

seconds (s)

mole (mol)

kelvin (K)

Units

• Units tell us what scale or standard of measurement is being used.

• A unit of measurement is a defined quantity.• In science, the SI system (international system) is

used.

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Prefixes

• Prefixes change the size of the units and are added before (pre) the base unit.

• The metric system prefixes are base 10> Each "step" is either:

– 10 times larger– 10 times smaller

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Example 1: Practicing Prefixes

• Millimeters are 10 times larger than __________

• 1 deciliter = liter

• 1 gram = mg

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Metric System Conversions

• First find the differences in powers of ten.• When going from a LARGER to a smaller prefix, move

decimals to the right.• When going from a smaller to a LARGER prefix, move

decimals to the left.• *Remember, moving decimals = multiplying or dividing

by ten

Example 2:• 5.5 L = ml• 190 cm = m• 3.1 km = cm• 980 ng = g

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1 Complete the following conversion:

100 millimeter = meter

A 100000 meter

B 1 meter

C 0.1 meter

D 10 meter

E 0.01 meter

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2 Complete the following conversion:

579 mm = m

A 579000 m

B 5790 m

C 0.0579 m

D 0.579 m

E 0.00579 m

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Scientific Notation

• In chemistry, we sometimes have to deal with measurements that are very BIG or very small.

• Scientific notation allows us to write very big or very small numbers in a shortened way.

• Numbers are expressed as a product of two factors:> A number between 1 and 10 > A factor of 10 (10n)> Example: 1500 = 1.5 x 103

Did you know?• The mass of a proton is

0.00000000000000000000000000167262kg.

• 1 gram of sugar (sucrose) is approximately 1760000000000000000000 molecules of sugar.

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How to put a number in scientific notation:

1) Move the decimal point to produce a factor between 1 and 10 (one digit to the left of the decimal)

2) Multiply by the appropriate exponent of 10 (10n)

I. Count how many places the decimal point was moved (n).

II. If the decimal moved to the right, the power of 10 is negative (10-n)

III. If the decimal moved to the left the power of 10 is positive (10n)

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Example 3:Put the following numbers in scientific notation.• 0.000567

• 18113100

• 0.0998

• 774000

How to put a number in scientific notation:

1) Move the decimal point to produce a factor between 1 and 10 (one digit to the left of the decimal)

2) Multiply by the appropriate exponent of 10 (10n)

I. Count how many places the decimal point was moved (n).

II. If the decimal moved to the right, the power of 10 is negative (10-n)

III. If the decimal moved to the left the power of 10 is positive (10n)

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3 What is 890010 in scientific notation?

A 8.9001 x 10^5

B 8.9001 x 10^-5

C 8.9001 x 10^-3

D 8.9001 x 10^3

E 8.9001 x 10

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4 What is 5.61 x 10^-2 in standard notation?

A 0.000561

B 56.1

C 561

D 0.0561

E 5.61

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Uncertainty in Measurements

• When making measurements, there is always some limitation to how "exact" they can be. This can stem from two main sources:> Uncertainty due to equipment limitation (how well

our tools allow us to measure)– Calibration– Graduation (the marks used to read

measurement values)> User variation

Think back to your lab--at which station did you explore uncertainty due to equipment limitation? At which station did you explore human variation?

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Uncertainty in Measurements

• We need to minimize uncertainty.• How can we do that?

> Pick the right tools> Calibrate when necessary> Multiple trials> Consistently follow procedures between trials.

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How close to the actual or accepted value a series of measurements are.

How closely grouped a series of measurements are to each other.

Accuracy and Precision

• We need to evaluate our data for uncertainty and error.• Go back to your lab-- what definitions of accuracy and

precision did you come up with?• Precision

• Accuracy

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Accuracy

• Accuracy can be measured by percent error:

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Example 4:

Calculate the percent error for the following:

A student determines the density of an unknown metal to be 1.54 g/cm3. The accepted value of the density of the metal is 1.61 g/cm3.

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Example 5:Which of the following students' data is most accurate? Most precise? The actual density is 1.59 g/cm3.

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So now that we know a little bit about measurements...What 3 things are important when making measurements? Think back to your lab experience and what we have talked about in class.

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Lets look at one source of uncertainty in measurements: Limitations in our equipment• Think back to station 8. What happened when you

tried to measure a volume using the graduated cylinder and the beaker?

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Significant Figures

• When we record measurements, we have to make sure that we only record as many digits as we can actually measure.> That is, we have to express our uncertainty in our

measurement.

• Significant figures or sig. figs. are the recorded numbers of a measurement. It includes> All of the certain digits> and one uncertain digit

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Significant Figures: What does that look like?• Lets look at an example...

1 98765432

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Example 6:Read the volume using the correct significant figures.

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Determining the number of sig figs• Ignore leading zeros• Ignore trailing zeros, unless they come WITH a

decimal point.• Everything else is significant.• Exact numbers (conversions, counting numbers)

have infinite # of sig figs.• Example:

> 0.0005811000> 23000

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Determining the number of sig figs

Stop at the first nonzero digit, then count all the digits.

Pacific-AtlanticPresent

Absent

• 0.0005811000• 23000

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Example 7Determine the number of sig figs in the following measurements:

• 9870 m

• 0.00113 cm

• 657.13 g

• 100.00 lb

• 71,005 km

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5 How many sig. figs are in this measurement: 0.0302220 ft

A 3

B 4

C 5

D 6

E 7

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6 How many sig figs are in this measurement: 67.43 mL

A 1

B 2

C 3

D 4

E 5

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Operations with Sig. Figs.

• When measurements are used in calculations, the result needs to maintain the degree of uncertainty from the measurements.

• There are two sets of rules:> multiplication and division> addition and subtraction

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Multiplication and Division

• When multiplying or dividing measurements, the final answer can only have as many significant figures as the measurement with the fewest.

• For example:

2.30 cm x 1.5 cm x 19.02 cm = 65.619 cm3

"calculator answer"

2.30 cm x 1.5 cm x 19.02 cm = 66 cm3

Correct answer with sig figs.

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Addition and Subtraction

• When adding or subtracting measurements, the answer can only have as many decimal places as the measurement with the smallest number of decimal places.

• For example:

23.5 cm + 2.544 cm + 31.03 cm = 57.074 cm

"calculator answer"

23.5 cm + 2.544 cm + 31.03 cm = 57.1 cm

Correct answer with sig figs.

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Rounding off

• Round so that you have the correct number of sig figs.> if the digit to be removed is less than 5, the digit

before it stays the same.> if the digit to be removed is greater than 5, the

digit before it increases by 1.

• Round at the very END of your calculations.

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Example 8Complete the following using sig figs:

• 4.55 + 3.0 =

• (5.1 + 0.66)/1.33 =

• (5.67 x 12)/.098 =

• 991.0 x 65 =

• 890.11 - 433.0198 =

• (354 + 312 + 481)/(5.4 x 102)=

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7 Solve the following and round using correct sig figs: 0.0013 + 0.010

A 0.011

B 0.0113

C 0.01

D 0.012

E

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8 Solve the following and round using correct sig figs: 89.1 x 36

A 3207.6

B 3.2 x10^3

C 3200

D 3210

E

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Conversions: Dimensional Analysis

• Allows you to change from one unit to another.• Remember, just because you change the unit, the

amount doesn't change. • How to do unit conversions via dimensional analysis:

1. Set up an equality from the unit you have to the unit you want.

2. Set up the equality as a fraction so that the unit you have is on the bottom, and the unit you want is on top.

3. Multiply the measurement with the fraction. Units should cancel.

• Example: How many centimeters is 3.5 inches?

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Example 9:• 4.56 L = deciliter

• 5.4 ft = meters 1 ft = 0.3048 m

• 60.1 miles = km 1 mile = 1.60934 km

• 0.89 lbs = g 1 lb = 0.453592 kg

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9 1 in = 2.54 cm. The length of a desk is 43.1 cm. How many inches is that?

A 109 in

B 17.0 in

C 0.98 in

D 32 in

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How to measure amount of matter:• A mole is a unit used to measure amount of atoms,

molecules, ions, etc. (particles)• 1 mole = 6.022 x 1023 particles (Avogadro's number)• 1 mole of atoms has a mass equal to the average mass

in grams of an element on the periodic table> Example: 1 mole of H has a mass of 1.01 g

http://sciencenotes.org/periodic-table-wallpaper-muted-colors

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10 What is the mass in grams of 1 mole of sulfur?

A 16 g

B 2.5 g

C 32.07 g

D 10.1 g

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Moles, mass, and # of particles

• We can use dimensional analysis to convert between moles, mass, and # of particles.

• Example: What is the mass (g) of 2.3 moles of helium?

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mass (g) moles # particles

molar mass 6.022 x 1023

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11 If you want to convert from moles to grams, what conversion factor do you use?

A molar mass

B avogadro's number

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Example 10:

• 4.3 g F = moles F

• 0.889 mol Na = atoms of Na

• 5.4 x 1023 atoms of Cl = moles of Cl

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12 What is the mass in grams of 2 moles of helium?

A 4.00 g

B 8.00 g

C 13.11 g

D 2.00 g

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13 How many atoms of zinc are in 20 g of zinc?

A 1.84 x 10^23

B 8.0 x 10^22

C 1.0 x 10^7

D 2.1 x 10^1

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