Transcript
99
Laboratory Equipment
A p p e n d i x A
Beaker
©H
ayde
n-M
cNei
l, L
LC
Florence Flask
Watch Glass
Litmus PaperStirring Rod
Filter Flask
Clamp Holder
Clamp
Ring Stand
Hot Plate
Weighing Dish
Spotting Plate
Spatulas and Scoops
Evaporating Dish
Mortar and Pestle
RingForceps
Safety Goggles
Tongs
Test Tube Brush
17 mL Crucible
44 mm Crucible LidErlenmeyerFlask
Evaporating Dish
44 mm Crucible Lid
Test Tube Holder
100
A p p e n d i x A • Laboratory Equipment
Appendix
A
PlasticWash Bottle
©H
ayde
n-M
cNei
l, L
LC
Funnel
Test Tubes
Ruler
Thermometer, –20°C – 100 °C
Buret Clamp
Buret
Pasteur Pipet
Petri Dish
Graduated Cylinders
Centrifuge
Dropper/BeralPipets
Test Tube Rack
BüchnerFunnel
VolumetricFlask
PlasticWash Bottle
101
Units and Constants
A p p e n d i x BTable B.1. International System of Units (SI Units).
Quantity Unit Abbreviation
mass kilogram kg
length meter m
time second s
temperature Kelvin K
amount of substance mole mol
electric current Ampere A
Table B.2. SI Derived Units.
Quantity SI Unit Alternate Name Alternate Units
volumem3
dm3 liter L
velocity sm
acceleration sm2
force skgm
2 Newton N
energys
kgm2
2
Joule J
densitymkg
3
cmg
3
frequency s1
Hertz Hz
pressuremskg
2 Pascal Pa
powers
kgm3
2
watt W
electric potentials Akgm
3
2
volt V
102
A p p e n d i x B • Units and Constants
Appendix
B
Table B.3. SI prefixes.
Prefix Symbol Meaning
tera T 1012
giga G 109
mega M 106
kilo k 103
hecto h 102
deca da 101
deci d 10–1
centi c 10–2
milli m 10–3
micro µ 10–6
nano n 10–9
pico p 10–12
When using the prefixes for conversions, there are two ways to set up the conversion factor, as shown below.
Example: How many nm in 1 m?
11
10 1 10mmnm nm
99# #=c m
or
110
1 1 10mnmm nm9
9# #=-c m
Either calculation is correct and both give the same answer. It just depends on whether you want to think of 1 m 1 109 nm or 1 nm 1 10–9 m.
103
A p p e n d i x B • Units and Constants
Appendix
B
Table B.4. Conversion factors.
Length1 inch 2.54 cm (exactly)
1 mile 5280 ft1 m 39.37 in
1 km 0.6215 mi1 light-year 9.46 1015 km
1 A 1 10–10 m
Volume1 cm3 1 mL
1 quart 0.9463 L1 fluid ounce 29.57 mL
1 gallon 3.785 L1 gallon 4 quarts1 quart 2 pints
Mass1 pound 16 oz
1 pound 453.6 g1 kg 2.205 pounds
TemperatureT(K) T(°C ) 273.15
T(°C) 95 [ T(°F) 32 ]
Pressure1 atm 1.013 105 Pa
1 atm 760 mm Hg1 atm 14.70 lb/in2
1 mm Hg 1 torr
Energy1 J 0.2390 calories
1 Calorie 1 kcal 1000 calories1 BTU 1055 J
1 kWh 3.6 106 J1 eV 1.60 10–19 J
Table B.5. Constants.
Constant Abbreviation Value
Planck’s constant h 6.6256 10–34 J∙s
Avogadro’s number 6.0221367 1023
Charge on an electron e 1.6022 10–19 C
Electron radius re 2.81792 10–15 m
Mass of an electron 9.109387 10–28 g
Mass of a proton 1.672623 10–24 g
Mass of a neutron 1.674928 10–24 g
Atomic mass unit amu 1.66057 10–27 kg
Molar volume 22.41383 L/mol
Gas constant R.
K molJ
8 314:
. gK molL atm0 0820::
Speed of light c 2.99792458 108 m/s
Acceleration due to gravity g .sm9 81 2
Rydberg constant for hydrogen RH 1.0967758 107 m–1
105
A p p e n d i x CIons
Common Monatomic IonsMonatomic ions not following general rules for charge. There are other possible ions for many of these metals but these are the most common and are the ones you are responsible for knowing.
Name Formula
Chromium Cr2+, Cr3+
Manganese Mn2+
Iron Fe2+ or Fe3+
Cobalt Co2+ or Co3+
Nickel Ni2+ or Ni3+
Copper Cu+ or Cu2+
Zinc Zn2+
Silver Ag+
Cadmium Cd2+
Tin Sn2+, Sn4+
Mercury Hg22+ or Hg2+
Lead Pb2+, Pb4+
106
A p p e n d i x C • Ions
Appendix
C
Polyatomic Ions
Name Formula
Ammonium NH4+
Carbonate Co32–
Hydrogen carbonate or bicarbonate HCO3–
Hypochlorite ClO–
Chlorite ClO2–
Chlorate Clo3–
Perchlorate ClO4–
Chromate Cro42–
Dichromate Cr2O72–
Cyanide CN–
Thiocyanate SCN–
Hydroxide OH–
Nitrate No3–
Nitrite NO2–
Phosphate Po43–
Hydrogen phosphate HPO42–
Dihydrogen phosphate H2PO4–
Permanganate MnO4–
Peroxide O22–
Sulfite SO32–
Sulfate So42–
Hydrogen sulfate or bisulfate HSO4–
oxyanions and oxyacids
Description Anion Acid
2 less oxygens than “ate” compound Hypo ite Hypo ous acid
1 less oxygen than “ate” compound ite ous acid
MEMORIZE ate ic acid
1 more oxygen than “ate” compound Per ate Per ic acid
Example: Sulfate is SO42–, remove two of the oxygens to get hyposulfite (SO2
2–) which becomes hyposulfurous acid (H2SO2) with the addition of two hydrogens.
107
A p p e n d i x DSolubility
Solubility Rules1. Compounds containing alkali metal ions (Li+, Na+, K+, Rb+, Cs+) and the ammonium
ion NH4+ are soluble.
2. Nitrates (NO3–), bicarbonates (HCO3
–), acetates (C2H302–), and chlorates (ClO3
–) are soluble.
3. Halides except Ag+, Hg22+, and Pb2+ are soluble.
4. Sulfates (SO42–) are soluble, except Ag+, Hg2
2+, Pb2+, Ca2+, Sr2+, and Ba2+.
Insolubility Rules1. Carbonates (CO3
2–), phosphates (PO43–), and chromates (CrO4
2–) are insoluble except for those containing alkali metal or ammonium ions.
2. Hydroxides (OH–) are insoluble except for those containing alkali metals. Hydroxides of Ca2+, Sr2+, and Ba2+ are slightly soluble.
3. Sulfi des are insoluble except when paired with alkali metal ions, ammonium, Ca2+, Sr2+, or Ba2+.
108
A p p e n d i x D • Solubility
Appendix
D
ksp Values for Some Common Salts
Compound Formula ksp
aluminum hydroxide Al(OH)3 4.6 10–33
barium carbonate BaCO3 5.1 10–9
barium chromate BaCrO4 2.2 10–10
barium hydroxide Ba(OH)2 5 10–3
barium sulfate BaSO4 1.1 10–10
calcium carbonate CaCO3 3.8 10–9
calcium fluoride CaF2 5.3 10–9
calcium hydroxide Ca(OH)2 5.5 10–6
calcium phosphate Ca3(PO4)2 1 10–26
copper(I) chloride CuCl 1.2 10–6
copper(I) sulfide Cu2S 2.5 10–48
copper(II) chromate CuCrO4 3.6 10–6
copper(II) hydroxide Cu(OH)2 2.2 10–20
iron(II) carbonate FeCO3 3.2 10–11
iron(II) hydroxide Fe(OH)2 8.0 10–16
iron(II) sulfide FeS 6 10–19
iron(III) hydroxide Fe(OH)3 4 10–38
lead(II) chloride PbCl2 1.6 10–5
lead(II) chromate PbCrO4 2.8 10–13
lead(II) hydroxide Pb(OH)2 1.2 10–5
lead(II) sulfate PbSO4 1.6 10–8
lead(II) sulfide PbS 3 10–29
lithium carbonate Li2CO3 2.5 10–2
lithium fluoride LiF 3.8 10–3
magnesium carbonate MgCO3 3.5 10–8
magnesium fluoride MgF2 3.7 10–8
magnesium hydroxide Mg(OH)2 1.8 10–11
magnesium phosphate Mg3(PO4)2 1 10–25
nickel(II) carbonate NiCO3 6.6 10–9
silver bromide AgBr 5.3 10–13
silver carbonate Ag2CO3 8.1 10–12
silver chloride AgCl 1.8 10–10
silver chromate Ag2CrO4 1.1 10–12
silver iodide AgI 8.3 10–17
silver nitrite AgNO2 6.0 10–4
silver sulfide Ag2S 6 10–51
silver sulfite AgSO3 1.5 10–14
zinc carbonate ZnCO3 1.4 10–11
zinc hydroxide Zn(OH)2 1.2 10–17
zinc sulfide ZnS 2 10–25
109
A p p e n d i x EStandard Deviation
In most real experiments, the “true” value of a quantity is not known. Therefore, we must fi nd a way to use our data to get the best possible estimate of the true value for the quantity being determined. One common estimate of the true value is the mean (X–
). The mean is simply the arithmetic average of all the data points:
...XnX
nx x x xni 1 2 3= =+ + + +/
where X–
= the mean value (or average), = “the sum of,”Xi = the individual data points (i = 1, 2, 3, …, n), andn = the total number of data points.
One way to express precision is by means of the standard deviation. To discuss this, we must fi rst discuss the normal distribution.
If a very large number of determinations of a quantity are done, all of the values will not be exactly the same, due to random errors.
On these graphs, X–
represents the mean, which is the best estimate of the true value. The width of the curve indicates the precision of the measurements. A tall, thin curve would indicate good precision, while a broad, fl at curve would show poor precision.
The standard deviation can be used to measure the width of a normal distribution. The standard deviation is defi ned as:
1s
nX Xi
2
=-
-
^^
hh/
where s = the standard deviation, n = the number of observations.
110
A p p e n d i x E • Standard Deviation
Appendix
E
The usefulness of the standard deviation is that it is expressed in units of the original measurement, and can be used to describe the position of any observation rela-tive to the mean. It can be shown mathematically that, for a distribution with an infi nite number of replicate measurements, 68.3% of the observed values will fall within ± 1s of the mean; 95.5% will fall within ± 2s of the mean; and 99.7% within ± 3s of the mean.
Measured characteristic
Freq
uenc
y
68%
95%
99%
-3s -2s -1s +1s +2s +3s
68%68%
95%95%
99%99%
Figure E.1.
ExampleSuppose that a density determination of a liquid is done in the laboratory, and the following data are obtained:
Experiment Number Density (g/mL)
1 2.60
2 2.90
3 2.70
4 2.90
5 2.50
From this data, calculate the average and standard de-viation for the results.
Step 1: Calculate the sum of the data: Xi = 13.60
Step 2: Calculate the average of the values:
13.60 2.720XnX
di== =
/
Step 3: Calculate the deviation of each result (d = | Xi – X
–| ), the sum of the deviations ( | d | ), the square of d
values ( | d |2 ), and the sum of the square of the d values ( | d |2 ). Tabulate these values in a new table:
Experiment
NumberXi | Xi – X
–| | Xi – X
–|2
1 2.60 | 2.60 – 2.720 | = 0.12 0.014
2 2.90 | 2.90 – 2.720 | = 0.18 0.032
3 2.70 | 2.70 – 2.720 | = 0.020 0.00040
4 2.90 | 2.90 – 2.720 | = 0.18 0.032
5 2.50 | 2.50 – 2.720 | = 0.22 0.048
n = 5 Xi = 13.60 (Xi – X–
)2 = 0.13
Step 4: The standard deviation can then be calculated from the formula:
. .s5 10 13 0 177=-
=^ h
Thus, we could state that the result of the density de-termination together with its standard deviation is 2.72 ± 0.18 g/mL. (Note that the average cannot have more signifi cant fi gures than the measurements that make up the average and that the standard deviation has the same number of decimal places as the average.)
The value of the standard deviation gives us some idea of the spread of our data points, or the precision of our determinations. A student with a standard deviation in this case of 0.100 will have a higher degree of precision in his or her experiment, but it does not necessarily mean that the experiment has a high degree of accuracy.
It is very important to realize at this stage that you can
have a very small deviation in your data (indicating high
precision) but your result may be signifi cantly off (if the
accuracy is low).
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