-
Name __________________________________________________ Honors
Chemistry Summer Assignment
HONORS CHEMISTRY – SUMMER ASSIGNMENT In chemistry, you will be
learning the scientific names of elements and compounds, as well as
completing many mathematical calculations of chemical quantities.
Your summer assignment begins with learning some of these facts.
You will be quizzed on the names and symbols of the elements and
polyatomic ions in this packet throughout the year (given the
chemical symbol provide the properly spelled name or given the name
provide the proper
chemical symbol). You must know the spelling and symbol. All
elements are to be written as shown on this list with a capital
letter as the first letter and lowercase letter as the second
letter. Do not write in all caps, or in cursive. You will also be
quizzed on the metric prefixes, their meanings and the ability to
convert between them. Assignment:
1. Make flash cards (on index cards not cut strips of paper) of
the metric system prefixes at the bottom of this page, as well as
the elements and the polyatomic ions listed on the next page. a.
For the metric system, put the prefix on one side and the numerical
meaning on the other. b. For elements and ions, put the symbol on
one side and the name on the other. Please put only one element or
ion per card. c. You will be given a homework grade for the flash
cards. 2. Complete the attached worksheet packet, to be handed in
the first day of school as a second homework grade. You have been
given a periodic table in case there are elements in the worksheet
that are not on your list to memorize. Assume you have Chemistry
the first day of school, do not wait to find out your schedule!
Bring your cards and packet to school!
Metric Prefixes
Prefix Numerical Meaning
Kilo- (K___) 1000 (which is 103) BASE UNIT The main metric unit
(meter (m), liter (l) , gram (g), etc.) deci- (d___) 0.1 (which is
10-1 or a tenth) centi- (c___) 0.01 (which is 10-2 or a hundredth)
milli- (m___) 0.001 (which is 10-3 or a thousandth) micro (µ___)
0.000001 (which is 10-6 or a millionth)
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Name __________________________________________________ Honors
Chemistry Summer Assignment
ELEMENTS
Aluminum Al Argon Ar Barium Ba
Beryllium Be Bismuth Bi Boron B
Bromine Br Calcium Ca Carbon C Cesium Cs Chlorine Cl
Chromium Cr Cobalt Co Copper Cu Fluorine F
Gold Au Helium He
Gallium Ga Germanium Ge Hydrogen H
Iodine I Iron Fe Lead Pb
Lithium Li Magnesium Mg Manganese Mn
Mercury Hg Neon Ne Nickel Ni
Nitrogen N Oxygen O
Phosphorus P Platinum Pt Potassium K
Radon Rn Rubidium Rb Scandium Sc
Silicon Si Silver Ag
Sodium Na Strontium Sr
Sulfur S Titanium Ti
Tin Sn Uranium U Xenon Xe Zinc Zn
POLYATOMIC IONS
Polyatomic ions are groups of multiple atoms that have a charge
(positive or negative). The symbols shown below tell you what
elements are in the ion, how many atoms of each, and the charge.
For example: NH4+1 contains a nitrogen atom, four hydrogen atoms
and the entire group has a charge of +1.
Memory Hint: If you have two ions with similar names and the
only difference is the number of oxygen atoms in your ion: -ite
means smaller number of O -ate means larger number of O Hypo-
(smallest) and Per- (largest) are used if there are four ions with
similar names and different numbers of oxygen. ION NAME NH41+
ammonium ClO1- hypochlorite ClO21- chlorite ClO31- chlorate ClO41-
perchlorate CN1- cyanide OH1- hydroxide IO31- iodate NO31-
nitrate
ION NAME NO21- nitrite MnO41- permanganate CO32- carbonate O22-
peroxide SO42- sulfate SO32- sulfite PO43- phosphate CH3COO1-
acetate
-
Name __________________________________________________ Honors
Chemistry Summer Assignment
Memorization Tips: Elements/Symbols Over the years, my students
and I have developed several unique ways to help us remember the
symbols for the elements. Be warned - some are a little out
there!
Silver Ag If a person who is expecting a present of a gold
necklace receives a silver one. He might say, “Ag, I didn’t want
silver!”
Gold Au “Hey you, I want that gold necklace!” Said with “Hey
you” sounding like Au.
Bromine Br That brother of mine - Bro of mine! Calcium Ca “Caws
give milk!” Pronounced with an accent to make cows sound like
it’s spelled with an A. Chlorine Cl “You Clean with chlorine!”
Iron Fe “Fe, Fi, Fo, Fum, I’m an iron man!” Helium He If you
breathe in helium, you will laugh! He, He, He! Mercury Hg Greek
mythology - Hg stands for Helmet guy! Potassium K You will get
Kicked out of school for the double nasty! You can’t do the
first three letters and cannot say the next three! Sodium Na
“Naw, I don’t want any sodium!” Nickel Ni “Nick owes me a nickel!”
Oxygen O “Open your mouth wide to take in oxygen!” Lead Pb Pencil
broke! Silicon Si Silly con! Tin Sn A tin roof gets hot in the Sun.
Manganese Mn Take first three letters - Man Magnesium Mg Take first
three letters – Mag
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Name __________________________________________________ Honors
Chemistry Summer Assignment
Fill in the missing symbol/name of the element. The date of
discovery and the origin of the name are included for your
information only. You will only be responsible for the names and
symbols.
Symbol Name Date Origin of Name aluminum 1825 Latin, alumen =
astringent taste
Ar 1894 Greek, argos = neutral or inactive barium 1808 Greek,
baryos = heavy
Bi ~1450 German, wismut = white mass boron 1808 Arabic,
bawraq
Br 1826 Greek, bromos = stench C B.C. Latin, carbo = coal Cs
1860 Latin, caesius = blue
chlorine 1808 Greek, chloros = green gas Cr 1797 Greek, chroma =
color cobalt 1735 Greek, cobolos = goblin
Cu B.C. Latin, cuprum fluorine 1886 Latin, fluere = to flow
Ga 1875 Latin name, Gaul, of France germanium 1886 country,
Germany
Au B.C. Latin, aurum He 1895 Greek, helios = the sun H 1766
Greek, hydro genes = water former I 1811 Greek, iodos = violet
color
Fe B.C. Latin, ferrum lead B.C. Latin, plumbum magnesium 1803
Latin, magnesia = a place in Asia Minor
Mn 1774 Latin, magnes = magnet Hg B.C. Latin, hydragyrum = god
and planet
neon 1898 Greek, neo = new nickel 1750 German, goblin nitrogen
1772 Latin, nitro = native soda and gen = born
O 1771 Greek, oxys = sharp and gen = born P 1669 Greek,
phosphoros = light bringer
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Name __________________________________________________ Honors
Chemistry Summer Assignment
platinum 1735 Spanish, plata = silver K 1807 Latin, kalium radon
1900 originates from radium
Rb 1860 Latin, rubidius = red scandium 1879 Scandanavian
peninsula by its discoverer silicon 1823 Latin, silex = flint
Ag B.C. Latin, argentum sodium 1807 Latin, natrium
Sr 1808 town of Strontian, Scotland sulfur B.C. Latin, sulphur
tin B.C. Latin, stannum
Ti 1791 Greek mythology, first sons of earth U 1789 planet
Uranus Xe 1808 Greek, xenos = strange
zinc B.C. German, zink = like tin Write your answers in the
blanks below
1. Mg is _____________
2. Magnesium is _____________
3. Aluminum is _____________
4. Silicon is _____________
5. Fe is _____________
6. H is _____________
7. Cu is _____________
8. N is _____________
9. C is _____________
10. Helium is _____________
11. Oxygen is _____________
12. Copper is _____________
13. Calcium is _____________
14. Iron is _____________
15. Potassium is _____________
16. Hydrogen is _____________
17. Carbon is _____________
18. Nitrogen is _____________
19. O is _____________
20. F is _____________
21. Fluorine is _____________
22. Na is _____________
23. Sodium is _____________
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Name __________________________________________________ Honors
Summer Assignment
Spell the name of the following ions correctly: 1. NO21-
__________________________________ 2. CO32-
__________________________________ 3. ClO31-
__________________________________ 4. OH1-
__________________________________ 5. PO43-
__________________________________ 6. NH41+
__________________________________ 7. SO42-
__________________________________ 8. CN1-
__________________________________ 9. CH3COO1-
____________________________ 10. O22-
__________________________________ 11. NO31-
__________________________________ 12. IO31-
__________________________________ 13. MnO41-
____________________________ 14. ClO21-
__________________________________ 15. O22-
__________________________________
Write the symbol and charge of the following ions. 1. phosphate
____________________________ 2. sulfate
____________________________ 3. cyanide
____________________________ 4. hydroxide
____________________________ 5. carbonate
____________________________ 6. nitrate
____________________________ 7. acetate
____________________________ 8. chlorate
____________________________ 9. perchlorate
____________________________ 10. hypochlorite
____________________________ 11. iodate
____________________________ 12. nitrite
____________________________ 13. sulfite
____________________________ 14. peroxide
____________________________ 15. permanganate
_______________________
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Name __________________________________________________ Honors
Summer Assignment
Polyatomic Ion Puzzle
Across
2. chlorite
3. peroxide
5. chlorate
7. sulfite
8. permanganate
10. carbonate
11. nitrite
14. nitrate
15. sulfate
Down
1. hypochlorite
2. acetate
4. phosphate
6. iodate
9. hydroxide
10. perchlorate
12. cyanide
13. ammonium
DIRECTIONS: When writing symbols for this puzzle, write the
number before the + or – for the charge. For example, write N O 3 1
- NOT N O 3 - 1
-
Name __________________________________________________ Honors
Summer Assignment
Polyatomic Ion Puzzle
Across
2. ClO1-
3. CN1-
4. ClO31-
6. PO43-
11. IO31-
13. CO32-
14. MnO41-
15. NO21-
Down
1. SO32-
5. CH3COO1-
6. ClO41-
7. SO42-
8. OH1-
9. ClO21-
10. NH41+
12. O22-
15. NO31-
-
Name __________________________________________________ Honors
Summer Assignment
1 2 3 4
5
6
7
8 9
10 11
12 13 14
15 16
17
18
19 20 21
22 23
24
25
26
Across 5 He (6) 6 Ca (7) 8 O (6) 9 Fe (4) 10 Cu (6) 12 K (9) 15
Na (6) 16 Au (4) 18 H (8) 19 Be (9) 24 Cl (8) 25 S (7) 26 Ne
(4)
Down 1 N (8) 2 Ni (6) 4 As (7) 7 Zn (4) 11 P (10) 12 Pu (9) 13
Ar (5)
14 Si (7) 17 F (8) 19 B (5)
20 Pb (4) 21 Hg (7) 22 Ag (6) 23 C (6)
25
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Nam
e __________________________________________________ Honors
Sum
mer A
ssignment
Periodic Table of the Elements Q
uiz
Fill in
the b
lanks w
ith th
e ato
mic sym
bols o
f the first 2
0
elemen
ts. (You
do n
ot n
eed to
ad
d th
e atomic n
um
ber,
weig
ht o
r nam
e.) Write th
e elem
ent n
ames in
the b
lanks
belo
w.
Key:
elem
ent name
atomic num
ber
symbol
atomic w
eight
scandium
21
Sc 44.95591
titanium
22
Ti 47.867
vanadium
23
V
50.9415
chromium
24
Cr
51.9961
manganese
25
Mn
54.93805
iron 26
Fe 55.845
cobalt 27
Co
58.9332
nickel 28
Ni
58.6934
copper 29
Cu
63.546
zinc 30
Zn 65.409
gallium
31
Ga
69.723
germanium
32
Ge
72.64
arsenic 33
As
74.9216
selenium
34
Se 78.96
bromine
35
Br 79.904
krypton 36
Kr
83.798 rubidium
37
Rb
85.4678
strontium
38
Sr 87.62
yttrium
39
Y
88.90585
zirconium
40
Zr 91.225
niobium
41
Nb
92.90638 molybdenum
42
Mo
95.94
technetium
43
Tc [98]
ruthenium
44
Ru
101.07
rhodium
45
Rh
102.9055
palladium
46
Pd 106.42
silver 47
Ag
107.8682
cadmium
48
Cd
112.411
indium
49
In 114.818
tin 50
Sn 118.710
antimony
51
Sb 121.760
tellurium
52
Te 127.60
iodine 53
I 126.9045
xenon 54
Xe
131.293 (1)
(2) (3) (4) (5)
__________________
__________________
__________________
__________________
__________________
(6) (7) (8) (9) (10)
__________________
__________________
__________________
__________________
__________________
(11)
(12)
(13)
(14)
(15)
__________________
__________________
__________________
__________________
__________________
(16)
(17)
(18)
(19)
(20)
__________________
__________________
__________________
__________________
__________________
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2
Welcome to Chemistry. You will study the composition of matter
(anything that has mass and takes up space) and the changes it
undergoes. In order to prepare to study this material effectively,
you should have some background knowledge. This packet contains
that information. This packet should be completed over the summer
and returned at the specified time. It is your first assignment and
it will be discussed and graded. You may work with classmates, but
you will be held responsible for this material, so be sure you
understand it. I. Introduction There are 5 divisions of chemistry,
and they overlap somewhat:
1. Organic chemistry: the study of substances containing carbon
– often living, or once living things
2. Analytical chemistry: the study of the composition of
substances – how much of a chemical or substance is present 3.
Physical chemistry: the study involving prediction of the behavior
of chemicals – why does a substance do what it does under specific
conditions or just in general 4. Biochemistry: the study of the
chemistry of living organisms – how and why the chemistry of living
organisms works and affects lives 5. Inorganic chemistry: the study
of substances not containing carbon – nonliving things or things
that never lived, and how they interact or react. Questions: Which
division or divisions of chemistry might be used to examine: a) The
mechanism by which blood clots b) The amount of a toxic substance
found in a water supply c) The reason a metal melts at a specific
temperature d) The formation of chemical compounds made up of
metals
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3
II.The Scientific Method In order to pursue the study of matter
and the changes it undergoes, or to solve other
problems in science, it is important to use a logical and
predictable technique. The Scientific Method is one such technique.
It is a common sense approach applicable to not only scientific
questions, but many other, everyday quandaries as well.
There are variations on the Scientific Method, but most of them
contain the following steps: 1. Observations: Something needs
explanation – a problem, a question, an issue. 2. Hypothesis: A
suggestion which solves the problem, answers the question, explains
the issue is offered. It may be termed an “educated guess” because
it is based on some logical explanation or suggestion of cause.
Often stated as “If ___________, then ___________. 3.
Experimentation: The testing of the hypothesis, usually repeated
trials are necessary to assure that the results of the tests can be
accepted as genuine. 4. Theory: A theory, or explanation of the
results of the experiment is offered. The theory is a possible
answer, it cannot be proven to be true; it is possible to disprove
a theory. Some sources will list a fifth step to the Scientific
Method: 5. Scientific Law: A statement is offered which summarizes
the results of experiments and observations. A scientific law does
not try to explain the results, it merely states what the results
are. Questions: 1. Given the following five statements, identify
which of the five parts of the Scientific Method they describe. A)
Ash trees infested with ash borer insects die. B) Leaves are
falling off my ash trees and there are areas of bark which are
destroyed. C) Ash trees are infested by ash borer insects in
Oakland County. It is possible that my trees are also. D) I will
remove the bark in areas of my trees and look for characteristic
signs of ash borer disease. I will look at each tree in several
spots. E) If my trees have ash borer disease, they would lose their
leaves and their bark would be damaged. 2. Write your own example
of an observation or problem, and use the five steps of the
Scientific Method to solve it.
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4
III. Dimensional Analysis Many calculations are required in this
course. The technique used to solve most problems is called
dimensional analysis. It uses conversion factors in order to derive
answers. For example, suppose the question was “How many inches are
there in 3.5 feet?” There is a defined relationship between feet
and inches: 1 foot = 12 inches. If this relationship was written as
a fraction it could be said that 1 foot = 1 or, 12 inches = 1. This
is true because 1 foot = 12 inches. 12 inches 1 foot Now, returning
to my question, and recognizing that multiplying anything by 1 does
not change that which is multiplied, I could answer my question
using dimensional analysis. 3.5 feet · 12 inches = 42 inches Units
of feet/foot cancel, leaving inches. 1 foot 3.5 feet was multiplied
by a conversion factor that allowed feet to be converted to inches.
This method should be used whenever possible. Always begin these
problems with the amount given, above the amount was 3.5 feet. Then
use the appropriate conversion factor to convert from the units you
have to the units you want. In this example, the desired unit was
inches. Remember: the conversion factor is a fraction in which the
two components are EQUAL. i.e.: 1 foot = 12 inches. The conversion
factor can also be the number required for each, such as: each
student requires 2 pencils to take a test. How many pencils are
needed if 15 students are taking a test? 15 students · 2 pencils =
30 pencils Units of “student” cancel, leaving 1 student pencils
Questions: Use dimensional analysis to solve: 1. How many eggs are
there in 6.2 dozen? 2. Each automobile made by Ford requires 4
tires. How many tires are needed to manufacture 259 cars? 3. Each
student in chemistry requires 3 beakers in her lab drawer. How many
beakers are required for a class of 7 students?
-
5
IV. SI Units of Measurement The SI system is an internationally
accepted system of measurement. SI stands for the French “System
International”. It encompasses the metric system as well as
specific base units. It is necessary for a chemistry student to
know the base units of the SI system. QUANTITY MEASURED SI UNIT SI
SYMBOL Length Meter m Mass Kilogram kg Amount Mole mol Temperature
Kelvin K (not oK) Time Second s Volume Cubic meter (informally,
liter) m3 Pressure Pascal Pa Density Grams per cm3 or grams per mL
g/cm3 or g/mL Whenever these SI units are used, metric prefixes are
also used. These should be memorized as well. These prefixes allow
for using much larger and much smaller multiples of the base units.
For example, a megameter would be 1 x 106 meters. A kilogram = 1 x
103 grams. A centimeter = 1 x 10-2 meters. A micromole = 1 x 10-6
moles. PREFIX SYMBOL MEANING Mega M X 106 (a million times larger)
Kilo k X 103 (a thousand times larger) Deci da X 10-1 (ten times
smaller) Centi c X 10-2 (a hundred times smaller) Milli m X 10-3 (a
thousand times smaller Micro µ X 10-6 (a million times smaller)
Nano n X 10-9 (a billion times smaller) Pico p X 10-12(a trillion
times smaller) It is possible to convert between units by using
dimensional analysis. Example: How many kilograms are present in
5000 grams? 5000 grams · 1 kilogram (kg) = 5 kg 1x 103g (1000 g)
Example: How many nanograms are present in 20 g? 20 g · 1 x 109 ng
= 20,000,000,000 ng 1 g Questions: How many grams are present in 29
kg? How many millimeters are present in 2 meters?
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6
V. Scientific Notation In chemistry we use very large and very
small numbers to represent amounts and masses. In order to
effectively communicate these numbers without using lots of zeros,
we employ scientific notation. Each value is written as the product
of two numbers, one of which is a power of 10. The first number is
always a single digit number, which may or may not have any number
of decimal places. Numbers which are greater than 1 have a positive
exponent. Numbers which are less than 1 have a negative exponent.
For example: Converting to scientific notation: 36,000 would be
written as 3.6 x 104 23.67 would be written as 2.367 x 101 35069032
would be written as 3.5068032 x 107
0.003 would be written as 3 x 10-3 0.23999 would be written as
2.3999 x 10-1
0.000000000000000000000007 would be written as 7 x 10-24
Converting from scientific notation to equivalent decimal numbers:
3.55 x 103 would be written as 3550 6.887 x 10-6 would be written
as 0.000006887 Questions: Convert the following in scientific
notation: 45 ____________________ 699 _____________________ 2884
______________________ 28395____________________
0.344_________________ 45.83 _____________________
302.555___________________ 0.00043502 ______________________
Convert the following from scientific notation to their equivalent
decimal numbers: 3.692 x 106 ________________ 4.9 x
10-2___________________ 1.9734 x 105 __________________ 5.55050 x
10-9 __________________ 7 x 10-3__________________ 9.97 x
10-1_________________
VI. Accuracy and Precision
-
7
It is impossible for measurements to be perfect. Many sources of
error can exist when humans take measurements. All measurements
have some amount of uncertainty. Good precision concerning
measurements entails making a series of measurements. Precision is
increased with additional measurements. It is also increased with
an increased number of decimal places in the measurement. If I were
to use two scales to weigh an object, and one scale weighed to the
0.00 place, and the second to the 0.0000 place, I would get more
precise measurements with the scale that measured to the 0.0000
place (providing the level of accuracy in both was equal- see
below). If we were to take many measurements of the same object,
and the values obtained were very similar, We could state that our
measurements were PRECISE. That is, they agreed closely with one
another. If I were weighing a marshmallow on a scale that measured
to the hundredth’s place and obtained three values of 1.00 g, 1.01
g, and 0.99 g, I could state that I had good precision. My three
measurements were close to one another. These measurements were
PRECISE. Precision can only be determined if a SERIES of
measurements are taken. Good precision would require that the
measurements be close in value to one another. Does that mean that
they were also ACCURATE? Not necessarily. ACCURACY is defined as
how close the measurements are to the true value. If the true mass
of our marshmallow was indeed 1.00 g, then our measurements were
accurate and precise. If, however, the true mass of our marshmallow
was 1.50 g, then our measurements were precise, that is, they
agreed with one another, but not accurate, as they were far from
the true value, which was 1.50 g. In order to determine accuracy,
the true measurement must be known. Questions: 1. Three
measurements are made of the mass of a new automobile. The first is
3500 lbs., the second is 3502 lbs., and the third is 3501 lbs. The
actual mass is known to be 3500 lbs. Are these measurements of mass
accurate? ___________ Are they precise?_____________ 2. I buy a 2 x
4 board from Home Depot. I measure it with an extremely fine
measuring tool and find the width of the board to be 4 inches. Can
I say that the board was accurately measured by Home Depot?
_______________ Can I say that my measurement = 4 inches was
precise?______________ 3. In question # 2, how would it be possible
to increase the precision of my measurement?
__________________________________________________________________________________
4. I have a choice of two scales to use to weigh myself. The first
will weigh in pounds, the second scale in pounds with marks between
each to indicate the ½ pound weight. Which will be more precise if
both are accurate?
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8
VII. Significant Figures In chemistry the concept of significant
figures is always used when discussing or doing calculations with
measured amounts. Significant figures apply ONLY to measured
quantities. The principle behind significant figures is this: if I
am measuring any quantity, my measurement will consist of all the
digits I am sure of, and the last digit, which is a reasonable
estimate. Example: I use the grocery store scale to weigh my fruit.
The scale has divisions of pounds, with 10 divisions between each
pound to indicate tenths of pounds. So if I put my apples on the
scale and it reads 3.5 pounds and the arrow is not exactly on the
5, but it appears to be exactly between 3.5 and 3.6 pounds, I can
estimate that I have 3.55 pounds of apples. I am sure I have 3.5
pounds, and I think I have 3.55 pounds, but the last 5 is really an
estimate. I am sure of the 3.5, and reasonably sure of the 3.55,
but that last 5 is not certain. SO: significant figures include all
digits we are sure of PLUS the last one, which is a very good
estimate. Questions: Which digit of the following measurements is
NOT significant, but a very good estimate? 2.379 inches ___________
19.334 m ____________ 890.5 miles_______________ In addition, there
are rules that determine whether digits are significant or not.
Below are the Rules: 1. All non-zero digits and zeros BETWEEN
non-zero digits, are significant. Example: 245 has 3 significant
figures (sig figs) 2405 has 4 sig figs. 2. Zeros at the END of a
number and to the right of a decimal point ARE always significant.
Example: 23.0 has 3 sig figs. 23.00 has 4 sig figs. 23.000 has 5
sig figs. 3. Zeros appearing to the left of all non-zero digits are
NOT significant if they just act as place holders. Example: 0.0071
has 2 sig figs (the 7 and 1) The three zeros to the left of the 71
are Not significant because they are merely place holders. 4. Zeros
appearing to the right of non-zero digits where NO DECIMAL POINT is
present are NOT significant. These are merely placeholders. Example
7100 has 2 sig figs, the 7 and the 1. The two zeros are merely
placeholders. There is no decimal point. [If the number had been
written as 7100., there would be 4 sig figs because the decimal
point appearing after the 7100. makes the two zeros significant.]
5. In two cases, there are an unlimited number of sig figs., i.e,
they cannot be quantified. These two cases are a) when objects are
counted (you cannot have 5.3 toes on your left foot) and b) when
numbers are used as conversion factors, i.e.; 12 inches/ 1 ft. Here
sig figs do not enter the equation.
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9
VIII. Using Significant Figures in Calculations When making
measurements, and using those measurements in a calculation, the
result of the calculation can NEVER have more accuracy than the
measurement with the least amount of accuracy. So, in order to obey
this rule of Sig Figs, it is necessary to define how we will do
calculations. ADDITION OR SUBTRACTION If the calculation involves
addition or subtraction, the answer can have no more decimal places
than the measurement with the LEAST number of decimal places.
Example: 12.1 inches + 2.1 inches + 3.03 inches = According to our
calculator, the answer to this question is 17.23 inches. BUT: using
sig fig rules we would give the answer as ~ 17.2 inches. Yes, the
calculator says 17.23 inches, but sig fig rules say we can only
have an answer with ONE DECIMAL PLACE, because our least precise
measurements, the 2.1 and 12.1 inch measurements only have one
decimal place. OUR ANSWER CAN ONLY HAVE ONE DECIMAL PLACE if we are
ADDING OR SUBTRACTING. Do all the calculations on the calculator,
then ROUND for significant figures. If adding or subtracting, put
all the digits in the calculator, obtain an answer, then round off
to the correct number of sig figs. Questions: Do the following
calculations and round for significant figures: a) 8.7 g + 15.43 g
+ 19 g = b) 853.2 L – 642.333 L = MULTIPLICATION OR DIVISION When
multiplying or dividing, and using sig figs, the final answer can
have only as many sig figs as the measurement with the LEAST number
of sig figs. Example: 6.3 cm x 2.2 cm = 13.86 cm2 according to my
calculator. Using sig figs, my answer would be ~ 14. I must round
my answer to 2 sig figs because each measurement in the problem,
the 6.3 and 2.2 cm, each only have 2 sig figs. Example: 6.66 cm /
3.2 cm = 2.08125 cm2 according to my calculator. Using sig figs, my
answer is rounded to 2.1. The measurement in the problem with the
least number of sig figs is the 3.2 cm. Thus, my answer cannot have
more than 2 sig figs, because 3.2 has 2 sig figs. Questions: Do the
following calculations and round for significant figures: a) 4.32
cm x 1.7 cm = c) 5.40 m x 3.21 m x 1.871 m = b) 38.742 kg ÷ 0.421
kg = d) 38.0 in ÷ 2.121 in =
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10
IX. Qualitative and Quantitative Measurements Qualitative
measurements are descriptive – “It is really cold today.” This
sentence describes the temperature without stating an exact
temperature. Quantitative measurements are specific – “It is 10
degrees below zero today.” This sentence tells us exactly how cold
it really is. It quantifies, or gives an exact quantity, to the
temperature. Both qualitative and quantitative measurements have
value. Questions: Give an example of a qualitative measurement you
might make in everyday life. Give an example of a quantitative
measurement that you might make in everyday life. X. Temperature
Conversions Temperature measurements are important in chemistry. We
will employ two temperatures scales, the Celsius and Kelvin scales.
Conversion between the two is much simpler than any Fahrenheit
conversions you have learned in other science classes. Forget
Fahrenheit. It is not used in chemistry. The Celsius scale is named
after the Swedish astronomer Anders Celsius. Water freezes at 0oC,
and water boils at 100oC. Easy to remember and IMPORTANT. Room
temperature is between 20oC and 25oC, and body temperature is 37oC.
The Kelvin scale is named after Lord Kelvin, and is designated as
just Kelvins. Not degrees Kelvin. Kelvin degrees are the same size
as Celsius degrees, so that an increase of 1oC is the same
magnitude of temperature increase as an increase of 1 K. 0oC = 273
K SO water freezes at 273 K 100oC = 373 K SO water boils at 373 K
To convert from Celsius to Kelvin, simply add 273 to the Celsius
temperature. To convert from Kelvin to Celsius, simply subtract 273
from the Kelvin temperature. Example: 10oC = 283 K 298 K = 25oC.
Questions: Do the following conversions: 4 K = ______oC 170oC =
_______K 87 K = ______oC 222oC = ___________ K