Chapter 1: Methods and Measurement
The Discovery Process
In the scientific world of controlled experiments, chance is seldom
acknowledged as a contributing factor in important discoveries. There are,
however, rare exceptions. In 1945 three men shared the Nobel Prize in
physiology or medicine for the discovery and isolation of penicillin, an
antibiotic medicine with great therapeutic potential.
Those three men were Alexander Fleming, Ernst Chain
and Howard Florey. Yet, despite the work of these three,
and related research by other scientists, most textbooks
credit a chance observation, made in 1928 by Fleming
alone, for the discovery of penicillin. How rare was this
serendipitous event and was the discovery of penicillin
really the result of an unexpected chance observation by a
single researcher?
Chemistry
Definitions of chemistry on the Web:
the science of matter; the branch of the natural sciences dealing with the
composition of substances and their properties and reactions
www.cogsci.princeton.edu/cgi-bin/webwn
the way two individuals relate to each other; "their chemistry was wrong from
the beginning -- they hated each other"; "a mysterious alchemy brought them
together"
www.cogsci.princeton.edu/cgi-bin/webwn
The study of the composition, structure, and other properties of substances and
the transformations that they experience.
www-lep.gsfc.nasa.gov/lepedu/glossary.html
scientific study of the composition, structure, and properties of substances and
the changes they undergo
pharmacy.ucsf.edu/glossary/
Major Areas of Chemistry
The Chemistry is typically viewed as having five major areas:
 Analytical
 Biological (Biochemistry)
 Inorganic
 Organic
 Physical.
A major in chemistry or biochemistry can lead to a variety of careers besides research.
These include the health professions, teaching, law, business, and management. A major
in one of these fields is also an excellent preparation for a number of interdisciplinary
areas including pharmacology, material science, nutrition and food chemistry,
neuroscience, toxicology, forensic science, and art conservation.
Scientific Method
The scientific method is typically noted for its orderliness and control; In fact, we are
taught that without these characteristics, experimental research may yield invalid
results. Therefore, chance should play little or no role in the process of the scientific
method. But what is chance? When is chance truly an accident and when is it
foreseeable? Historically, some chance discoveries have led to startling new ideas that
eventually directed important further scientific investigation of natural phenomena.
The scientific method has six steps
1. Observation.
2. Formulation of a question
3. Pattern recognition
4. Developing theories. (hypothesis and eventfully theory)
5. Experimentation
6. Summarizing information.( scientific laws)
Hypothesis - a guess to try to explain observed data.
The hypothesis is tested by designing and performing experiments.
Law - large amount of data summarized in a brief statement
Laws are universal and hold everywhere in the observable universe.
Boyle's Law is P1 V1 = P2 V2
Theory - best explanation for various phenomena as of today (tentative)
Theories are modified or discarded with new observations.
Theories are valuable for their predictive value.
Models - use of tangible items or pictures to represent invisible processes
Models in Chemistry
One of the most commonly used constructs of physical reality is a model. A model is a
simple way of describing and predicting scientific results, which is known to be an
incorrect or incomplete description. Models might be simple mathematical descriptions
or completely non-mathematical. Models are very useful because they allow us to
predict and understand phenomena without the work of performing the complex
mathematical manipulations dictated by a rigorous theory. Experienced researchers
continue to use models that were taught in high school and freshmen chemistry,
however they realize that there will always be exceptions to the rules of these models.
Science and Technology
Science
A method of learning about the physical universe by applying the principles of
the scientific method, which includes making empirical observations, proposing
hypotheses to explain those observations, and testing those hypotheses in valid
and reliable ways; also refers to the organized body of knowledge that results
from scientific study.
Technology
Technology is application of scientific knowledge to develop tools, materials,
techniques, and systems to help people meet and fulfill their needs.
Matter and Properties
Matter - occupies space (volume) and has mass.
Mass - measure of the quantity of matter in an object compared to a standard
mass.
Weight - measure of force of attraction between earth and object as using spring
balance.
Energy- Manifestations of matter
Matter and force (energy) are the two fundamental entities of which the universe is
composed. All that exists can be classified in these terms. All environmental
phenomena occur because of the interactions between matter and transformations of
matter in space and time. As the arrangements between forces and masses change, the
change is manifested in terms of energy.
List four different forms of energy: kinetic energy, potential energy
heat, light, electrical, mechanical, chemical (stored in bonds)
Physical properties
Physical characteristics/behavior of a substance are:
Temperature, mass, structure, color, taste, odor, boiling point, melting point, freezing
point, heat capacity, hardness, conductivity, solubility, density
Chemical changes and properties
Chemical Change
Chemical change is associated with change in the substances and a chemical reaction
which is normally nonreversible. E.g. Burning piece of magnesium in oxygen produce
magnesium oxide a new substance is formed.
Chemical properties
Chemical properties relates to the changes of substances making up the matter. For
example, corrosiveness refers to the readiness of certain metals to react with other
elements such as oxygen and chlorine. Flammability refers to readiness of organic
matter to react with oxygen to produce carbon dioxide and water. It also indicates the
reactivity of certain substances towards another. A listing of all the possible chemical
reactions of a substance is called its chemical properties.
CHEMICAL
PHYSICAL
CHANGE
Old substance destroyed
New substance formed
New form of old substance No
new substances formed
PROPERTIES
List of all possible
chemical changes
Description by sense- color,
shape, odor, etc. Measurable
properties - density, boiling point,
etc.
Chemical vs. Physical Changes.
 In a physical change, the substances are not altered chemically, but merely
changed to another phase (i.e. gas, liquid, solid) or separated or combined.
 In a chemical change, the substances are altered chemically and display different
physical and chemical properties after the change.
Practice on Identifying Chemical and Physical Properties
Water boils at 100 degrees Celcius. P
Diamonds are capable of cutting glass. P
Water can be separated by electrolysis into hydrogen and oxygen. C
Sugar is capable of dissolving in water. P
Vinegar will react with baking soda. C
Yeast acts on sugar to form carbon dioxide and ethanol. C
Aluminum has a low density. P
Ammonia is a gas at room temperature. P
Bromine has a red color. P
Dry ice, solid carbon dioxide, is sublimed at room temperature. P
Salt is dissolved in water. P
Iron rusts in a damp environment. C
Hydrogen peroxide decomposes to water and oxygen. C
Intensive and Extensive Properties
Intensive properties
Properties such as temperature and pressure do not depend upon the sample size.
One portion is then arbitrarily subdivided into three subportions. For intensive
properties, the value of the property for each subportion is the same.
Extensive properties
Properties such as mass and volume depend upon sample size. Two identically sized
portions of a solution are prepared.
For extensive properties, the sum of the values of the property for the subportions
equals the value for the whole portion.
Classification of matter
Anything that has mass and volume is matter. Matter is also defined as anything with
the property of inertia. All of the solids, liquids and gases that you may encounter in
your daily life would be classified as some type of matter. You are familiar with the
taxonomy of living things from Biology. Now you will learn how scientists classify
matter that makes up everything.
Atoms & Molecules
Atoms - smallest characteristic part of an element.
Molecule - group of atoms bound together as a neutral unit.
Elements
Elements are substances that cannot be separated into simpler substances. Salt is made
up of the elements sodium and chloride. Water is made up of the elements hydrogen and
oxygen.
Ions
Ion is a charged atom or molecule.
Na+ ( mono-atomic cation), Cl- ( mono-atomic anion)
Polyatomic ion NH4+ (polyatomic-cation), SO42- ( poly-atomic anion)
Homonuclear: made up of same kind of atoms (same elements).
Heteronuclear: made up of different kind of atoms (different elements).
Chemical Symbols
How are elements given symbols?
Chemical symbols can be one or two letters. The first letter is always a capital case and
the second letter is always a small case. Some symbols are taken from the Latin or
German names of elements.
Na = sodium, K = potassium, Fe = iron, Cu = copper, Ag = silver, Sn = tin, Sb =
antimony, W = tungsten,
Au = gold, Hg = mercury, Pb = lead
Compounds
A compound is a substance formed when two or more elements are chemically joined.
Water, salt, and sugar are examples of compounds. When the elements are joined, the
atoms lose their individual properties and have different properties from the elements
they are composed of. A chemical formula is used a quick way to show the composition
of compounds. Letters, numbers, and symbols are used to represent elements and the
number of elements in each compound. Chemical compound can be broken down into
its elements by chemical means.
Chemical compounds
Naming chemical compounds (systematic names)
Ionic compounds are formed from reaction of electropositive metal elements with
electronegative non metal elements
Eg. Name: aluminium bromide Formula: AlBr3
Name: sodium chloride
Formula: NaCl
Name: calcium oxide
Formula: CaO
Molecualr or covalent compounds are formed from reaction of non-metal elements with
non metal elements
Eg. Name: Carbon dioxide
Formula: CO2
Name: sulfur hexafluoride Formula: SF6
Name: calcium oxide
Formula: CaO
Classification of Matter
Matter
Anything with mass and volume.
Pure Substance
Matter with constant composition
Element
Substance
made up of
only one type
of atom
Examples gold, silver,
carbon,
oxygen and
hydrogen
Compound
Two or more
elements that are
chemically
combined
Mixture
Matter with variable composition
Heterogeneous Mixture
Mixtures that are made
up of more than one
phase
Examples - water, Examples - sand, soil,
carbon dioxide,
chicken soup, pizza,
sodium
chocolate chip cookies.
bicarbonate,
carbon monoxide
Homogeneous Mixtures
Also called solutions.
Mixtures that are made
up of only one phase
Examples - salt water,
pure air, metal alloys,
seltzer water.
Give examples of each of these categories of matter given below:
a) mixtures
d) substances
b) heterogeneous mixtures
e) elements
c) homogenous mixtures
f) compounds
a) Mixtures: Mixtures are aggregates of matter two or more substances.
b) Heterogeneous mixtures: Aggregates of matter containing regions with different
composition and properties. Eg. garbage and soil.
c) Homogeneous mixtures: Aggregates of matter where composition and properties of
the mixture is same throughout. E.g. milk and solution of alcohol. Solutions are always
homogenous mixtures.
d) Pure substances: Substance is what matter is made up of. A pure substance is made
up of either elements or compounds. Pure substances are also matter with constant
composition and properties. For example water is has a constant ratio of hydrogen and
oxygen , and it has a b.p. 100oC and m.p. 0oC.
e) Elements: Element is a pure substance that cannot be broken down to other
elements or pure substances by chemical means. E.g. copper and helium.
Measurement in Chemistry
Data, Results, and Units
Data
Facts represented in a readable language (such as numbers, characters, images,
or other methods of recording) on a durable medium. Data on its own carries no
meaning. Empirical data are facts originating in or based on observations or
experiences. A database is a store of data concerning a particular domain. Data
in a database may be less structured or have weaker semantics (built-in
meaning) than knowledge in a knowledge base.
Results
Results are the outcome of a designed experiment, often determined from
individual bits of data.
Units
In England units of measurement were not properly standardized until the 13th
century, though variations (and abuses) continued until long after that. For example,
there were three different gallons (ale, wine and corn) up until 1824 when the gallon
was standardized.
In France the metric system officially started in June 1799 with the declared intent of
being 'For all people, for all time'. The unit of length was the meter which was defined
as being one ten-millionth part of a quarter of the earth's circumference.
Le Systeme international d'Unites (SI units) officially came into being in October 1960
and has been officially recognized and adopted by nearly all countries, though the
amount of actual usage varies considerably. It is based upon 7 principal units, 1 in each
of 7 different categories Category
Name
Abbrev.
Length
metre
m
Mass
kilogram
kg
Time
second
s
Electric current
ampere
A
Temperature
kelvin
K
Amount of substance
mole
mol
Luminous intensity
candela
cd
Measurement in Chemistry
Metric Units One base unit for each type of measurement. Use a prefix to change the
size of unit.
Learn the Exponential Expressions, Decimal Equivalents, Prefixes, and Symbols for the
prefixes below:
Convert measurements to a unit that replaces the power of ten by a prefix:
6.80 x 10-9 m
7.14 x 10-6 s
2.54 x 10-2 m
4.56 x 103 g
2.88 x 10-3 g
Answers:
Volume
Area unit: m2 is the base unit of area.
1 cm = 10-2 m
square each number and unit in the conversion factor
(1)2 cm2 = (10-2)2 m2
1 cm2 = 10-4 m2
Volume unit: m3 is the base unit of volume.
1 cm = 10-2 m
cube each number and unit in the conversion factor
(1)3 cm3 = (10-2)3 m3
1 cm3 = 10-6 m3
This is the cubic centimeter.
English and Metric Units
Important Conversion Factors
12 in = 1 ft,
3 ft = 1 yd,
5280 ft = 1 mi,
5.5 yds (exactly) = 1 rod,
year
40 rods = 1 furlong,
16 oz = 1 pt,
2 pt = 1 qt,
4 qt = 1 gal, 20
8 qt = 1 peck,
16 oz = 1 lb,
2000 lbs = 1 ton,
scruples = 1 grain,
3 scruples = 1 dram,
4 pecks = 1 bushel, 96 drams = 1 lb
60 sec = 1 min
60 min = 1 hr
24 hr = 1 day
365.25 days = 1
1 mL = 1 cm3
Unit Conversion: English and Metric Systems
Important Metric to English and English to Metric conversions:
Mass: 2.205 lbs/kg,
Length: 1.609 km/mi,
Volume: 3.785 L/gal.
453.59 g/lb
2.54 cm/in, 3.281 ft/m
Temperature and Temperature Conversions
Astronomers and other scientists like to use a temperature scale called "Kelvin." On a
Kelvin thermometer water freezes at 273 degrees and boils at 373 degrees. Zero degrees
Kelvin is called "absolute zero." It is the lowest possible temperature of matter.
To convert degrees Celsius to Kelvin,
o
C -->K ; K = C + 273.15
simply add 273 degrees. (Example, 273 K = 0 deg. C)
To convert Kelvin to degrees Celsius,
o
C -->K ; K = C - 273.15
simply subtract 273 degrees. (Example, 273 K = 0 deg. C)
To convert Celsius to Fahrenheit
o
C-->oF ; C = 9/5C + 32
begin by multiplying the Celsius temperature by 9, then divide the answer by 5. Finally
add 32.
To convert Fahrenheit to Celsius
o
F-->oC ; C = 5/9 (F - 32)
begin by subtracting 32 from Fahrenheit temperature, then multiply by 5 divide and
the answer by 9.
Human body temperature is 98.6 oF convert this temperature to oC and K
scale.
o
C = 5/9 (98.6 - 32) = 5/9 (66.6) = 37.0
o
C--> K = 37.0 oC +273.15 = 310.2 K
Error, Accuracy, Precision, and Uncertainty
Precision versus Accuracy:
Precision = how measurements agree.
Accuracy = how close measurement is to true value.
Uncertainty is expressed in terms of the rounding off to significant figures
Significant Figures
Significant Figures = Digits which express the precision of a measurement.
In recording a measurement, write all the numbers with certainty
plus the FIRST NUMBER WITH UNCERTAINTY.
All non-zero NUMBERS are significant.
NOTE: ZEROS used to "place" the decimal are NOT significant figures.
0.015 grams = 2 SF (LEADING zeros BEFORE all the digits are NOT significant)
1.500 grams = 4 SF (TRAILING zeros after all the digits are SIGNIFICANT)
10.5 grams = 3 SF (SANDWICH zeros WITHIN a number are SIGNIFICANT)
0.0105 grams = 3 SF (examples of LEADING zeros BEFORE as well as SANDWICH
zeros WITHIN)
Exact numbers = definitions (not measurements).
Exact numbers have "infinite" numbers of significant figures.
2.54 cm = 1 inch
Significant Figures in Calculation of Results
How many significant figures are retained in addition/subtraction:
When adding or subtracting numbers, all numbers must have the same units.
The answer can have no more decimals than the measurement
with the fewest DECIMALS.
254
ml
25.4 ml
2.54 ml
======
281.94 ml
???
254 ml
- 54.1 ml
=====
199.9 ml
???
how many significant figures are retained in multiplication/division
The answer can have no more significant figures than the measurement
with the fewest SIGNIFICANT FIGURES.
(231.54 * 43)/433.4 = 22.972358
???
Practice exercises
How many significant figures are in the following numbers:
a) 0.0945
b) 83.22
c) 106
d) 0.000130
Deduce the number of significant figures contained in the following:
a) 16.0 cm
b) 0.0063 m
c) 100 km
d) 2.9374 g
e) 1.07 lb/in2
How many significant figures are in the following measurements?
a) 25.9000g
f) 21.2 m
b) 102 cm
g) 0.023 kg
c) 0.002 m
h) 46.94 cm
d) 2001 kg
i) 453.59 g
e) 0.0605 s
j) 1.6030 km
Significant Figures and Scientific Notation
Used to express very large or very small numbers
as decimal number between 1 and 10 times a power of 10.
133 000 grams = 1.33 x 105 grams
Positive exponent tells how many spaces to move the decimal to the right
to convert to a decimal number.
0.000 133 grams = 1.33 x 10-4 grams
Negative exponent tells how many spaces to move the decimal to the left
to convert to a decimal number.
Practice exercises
Express the following in scientific notation:
a) 0.00839 b) 83.264 c) 372 d) 0.0000208
e) 936,800
f) 1638
g) 0.0000568
h) 0.009117
Convert the following numbers into non-xponential form:
a) 2.3 x 10-3
b) 2.9 x 102
c) 3.92 x 10-4
d) 1.73 x 104
Adition/subtraction using scientific notation:
To add/substract numbers written in scientific notation, all number
must have THE SAME EXPONENT.
2.54 x 105 grams + 2.54 x 103 grams = 256540 grams
First, convert to same exponent.
NOTE: if decimal gets larger/exponent gets smaller.
If decimal gets smaller/exponent gets larger.
Second, decide how many significant figures to report.
2.54 x 105 grams + 0.0254 x 105 grams = 2.5654 x 105 grams = 2.57 x 105 grams
Mltiplication/division using scientific notation:
Multiplication: multiply the decimals, add exponents, use correct SF
1.35 x 10-5 times 2.35 x 103 = 0.031725
???
Division: divide the decimals, subtract exponents, use correct SF
1.35 x 10-5 divide by 2.35 x 103 = 5.745 x 10 -9
???
Rounding Off Numbers
In measuring or calculating, your answers will be expressed in significant figures.
Measurements and calculations are either demonstrating a physical reality or are
meaningful to you as a scientist You might be given specific directions about how
many significant figures you should use.
For each example, you are directed to round off to 3 significant figures.
1. If the first "extra" digit is LESS than 5-drop it. Now the last digit of the
number remains the same.

Ex. 4.321 becomes 4.32
2. If the first "extra" digit is 5 or MORE than 5, drop the number and increase
the last significant digit by 1.
Ex. 4.336 becomes 4.34
What is this "even/odd rule" I keep hearing about?
When digit is exactly 5:occasions when the number you are rounding off is exactly half
way between the two choices (the left most digit to be dropped is a 5 and there are no
nonzero digits after it) then it makes sense to round up half the time, and round down
the other half of the time. You choose when to round up and when to round down by
choosing whichever option will give an even number as the answer (last digit is 0, 2, 4,
6, or 8). Refer to the examples above.
Significant figure retained when combining numbers:
Calculate 3.21 cm x 15.091 cm
Calculate 3.82 x 1.1 x 2.003
Divide 13.87 by 1.23
Divide 0.095 by 1.427
Add 12.786 to 1.23
Add 3.961 to 24.6543
Subtract 2.763 from 3.91
Subtract 54.832 from 98.2
Carry out the following arithmetic operations:
(a) 65.336 + 47.893
(b) 32.4 – 0.128
(c) 65.3 * 0.065
(d) 6930 / 0.6975
(e) 45.65 + 32.63/8.000
Density and Specific Gravity
Definition
Density and specific gravity have very similar, but not quite identical definitions.
Density is the amount of something per unit volume. Most typically, one expresses the
mass per unit volume for a solid or liquid. For example, 5.2 g/cm3. We might express
this as g/m3.
For example, if we have 25 g of a material with a density of 0.798 g/cm3, what size
container will we require?
Specific gravity is a ratio of the mass of a material to the mass of an equal volume of
water at 4 oC (39 oF). Because specific gravity is a ratio, it is a unitless quantity. For
example, the specific gravity of water at 4 oC is 1.0 while its density is 1.0 g/cm3.
At 4 oC, the density of water is 1.0 g/cm3. Therefore, density and specific gravity have
the same numeric value at this temperature.
Food Calories (kilocalorie)
In scientific terms:
1000 calories = 1 kilocalorie (food calories) = 1 kcal = the energy it takes to raise the
temperature of 1kg of water by 1°C.

1 gram of protein = 4 Kilocalories (food calories)

1 gram of carbohydrate = 4 Kilocalories

1 gram of fat = 9 Kilocalories
One ounce = 28.3 grams.
Let's multiply and get an estimate that is easy to work with.
Factor label method for calculations
The key to easy conversions is to do them all the same way--use a routine. The steps to
one routine follow:

Write down the quantity you want to convert, including the units. (On every
conversion you want to change a quantity with its units to a quantity in some other
units.)

Multiply this by a "conversion factor" that has the units you want in the numerator
and the units you are starting with in the denominator. The numbers which go with
the units have to be given to you.
For example, you have to be told that 1 liter = 1.05 quarts if you want to convert
quarts to liters or liters to quarts. Usually one of the numbers in the two conversion
numbers will be 1.

If the 1 goes in the numerator of your conversion factor, the mathematical operation
you will have to do will be division. If the 1 goes in the denominator, the
mathematical operation you will have to do will be multiplication. You do not
always multiply the quantity you start with by the conversion factor.
For example, some people would multiply by 2.54 cm/in whether they were going
from inches to centimeters or from centimeters to inches. (They would get the
wrong answer when going from centimeters to inches.)
Conversion Examples
1. Convert 427 in3 to cm3 if 1 in3 = 16.4 cm3 (= 2.543 cm3).
Solution:
2. If 1 bushel = 35.4 liters, convert 16 liters to bushels.
Solution:
Make the following conversions. Use the following to assist you:
0.224 pounds = 1 Newton 2.54 cm = 1 in
3.281 feet = 1 meter
1.05 quarts = 1 liter
1.
5 qts. to liters
2.
10 Newtons to pounds
3.
20 centimeters to inches
4.
15 meters to feet
5.
10 foot pounds to Newton meters
Additional Practice
6.
Convert 25 U.S. gallons to Imperial gallons if 1 U.S. gallon equals 0.83 Imperial
gallon.
7.
Convert 2.50 American dollars to Mexican pesos if 1 American dollar equals 8
Mexican pesos.
8.
Convert 3 pints into cups if 1 cup equals 0.5 pints.
9.
Convert 5 tablespoons to ounces if 1 ounce equals 2 tablespoons.
10.
Convert 30 acres to square miles if 1 square mile equals 640 acres.
11.
Convert 50 feet to fathoms if 1 fathom equals 6 feet.
12.
Convert 5 square inches to square centimeters if 1 centimeter equals 0.3937 inch.
13.
Convert 60000 Newtons per square meter (Pascal) to psi if 1 lb equals 4.46 Newton
and 1 meter equals 39.37 inches
14.
Convert 30 psi to kiloPascal.
15.
Convert 25 m/sec to mph if 1 mile = 5280 ft.
16.
Convert 100 rpm to radians/second if 1 revolution = 6.283 radians
17.
Convert 100 N/m3 to lbf /ft3.
18.
Convert 20 gpm to liters/sec
Answers
1.
4.76 liters
10.
0.046 square miles
2.
2.24 pounds
11.
8.33 fathoms
3.
7.84 inches
12.
32.26 square centimeters
4.
49.5 feet
13.
8.68 psi
5.
13.51 Newton meters
14.
207.39 kPa
6.
20.75 Imperial gallons
15.
56.25 mph
7.
20 Mexican pesos
16.
10.47 radians/second
8.
6 cups
17.
0.623 lbf /ft3
9.
2.5 ounces
18.
1.27 liter/sec
You should be able to answer following questions:
1. Why do we have to worry about significant figures anyway?
2. How can you tell how many significant figures a particular number has?
3. Aren’t all whole numbers exact?
4. When you are rounding off numbers, how do you know whether to round up or
round down? And what exactly does "round up" or "round down" mean
anyway?
5. What is this "even/odd rule" I keep hearing about?
6. What happens if you have to round off a really big number?
7. How do you decide how many significant figures your answer should have when
you are adding or subtracting numbers?
8. How do you decide how many significant figures your answer should have when
you are multiplying or dividing numbers?
9.
What happens if you do a calculation where you have adding or subtracting
AND multiplying or dividing?
The Composition and Structure of the Atom
Matter and Structure
The structure of matter has been of interest since the times of the ancient Greeks, when
Demokritos proposed that the world was built up of elementary atoms (from the Greek
for indivisible). Only in the last century was the existence of atoms seriously
considered, and indeed only in this century was the physical reality of atoms finally
accepted.
Atomic Structure
We know now that the atom consists of a small, but massive, positively charged nucleus
which binds a number of negatively charged electrons in its electric field. This positive
nuclear charge determines the chemical properties of the elements. Thus an atom has Z
electrons each having a charge -e, with e = 1.602 x 10-19 coulombs. The nucleus has a
charge +Ze. Z is the atomic number which characterises an element.
Electron
The electron is, as far as we know, an elementary particle; i.e. it has no substructure.
But the nucleus is not elementary; it is built up of constituents, and the interactions of
these constituents and the structure they lead to is the subject of this course.
Nucleus
The nucleus consists (largely) of protons and neutrons (collectively known as
nucleons). A new force operates at the nuclear level to hold them together in the nucleus
-- this is the (strong) nuclear force, a short range interaction which does not extend
much beyond the nuclear boundaries.
Development of the Atomic Theory
Early scientist observed changes of matter (the first step in the scientific method). They
called these changes chemical reactions when there are changes in substances or the
physical properties of the matter. They also observed a pattern or a repeatable
observation in chemical reactions. They observed repeatedly that during chemical
reactions mass was neither destroyed nor created (conservation of mass), and elements
combine in constant proportions (definite proportions) and multiple proportion. In order
to summarize these patterns , they came up with three chemical Laws.
i) Law of Conservation of Mass:
ii) Law of Definite (Constant) Proportions:
iii) Law of Multiple Proportions:
To explain chemical Laws which includes law of constant proportions John Dalton in
1803-1807 came up with a Atomic Theory. The theory he postulated consists of four
postulates.
Dalton’s Theory
Postulates of Dalton's Atomic Theory:
i)
All matter is composed of indivisible atoms. Elements consists of tiny
particles called atoms. This postulate is the most important concept in his
theory to show clearly for the first time the existence of atoms.
ii)
An atom cannot be divided destroyed, converted or created.
iii)
An element is composed of identical atoms. They have similar mass and
volumes. This was necessary to explain the fixed properties of an element.
iv)
Atoms of different elements have different properties.
v)
Compounds are formed when atoms combine. This postulate was
necessary to explain the existence of compounds and the breaking of
compounds into elements.
vi)
Chemical Reactions involve rearrangement of atoms to form new
compounds. This postulate was necessary to define and describe chemical
reactions or changes.
Starting and final materials in a chemical reaction I called, reactants and products,
respectively.
For example, the combustion or burning of magnesium metal with oxygen can be
written with the formulas of all reactants and products: Reactants: (magnesium = Mg,
oxygen = O2), Products: Ash = magnesium oxide (MgO). The reaction is written as
Subatomic Particles:
Electrons
Electron was first discovered by J.J. Thompson using cathode-ray tubes or
Crook's tubes and calculated mass/charge ratio. According to modern atomic
theory an electron travels around a nucleus made up of protons and neutrons.
American physicist Robert Andrews Millikan (1868-1953) designed an
experiment to measure the electronic charge. Drops of oil were carried past a
uniform electric field between charged plates. lectrons are sub-atomic particles
with a mass of 9.11 x 10-28g ( 1/1833 times mass of a proton) and a negative
charge of 1.60 x 10-19 c (c=coulombs) or -1.60 x 10-19 c.
Nuclear atom
Rutherford’s experiment
A New Zealander, Rutherford fired Alpha particles at an extremely thin gold
foil. He expected them to go straight through with perhaps a minor deflection.
Most did go straight through, but to his surprise some particles bounced directly
off the gold sheet! This means there something in the atom that deflected the
alpha particles Rutherford hypothesized that the positive alpha particles had hit a
concentrated mass of positive particles, which he termed the nucleus.
Nucleus
The mass and the positive (+) charge of an atom are concentrated in the center.
This center is called nucleus, and it has radius of about 10-13 cm. Nucleus
contains protons and neutrons which are equal in mass.
Proton
Proton is a sub-atomic and sub-nuclear particle. A proton has a mass of 1.67 x 10-24 g
(which is 1833 times heavier than an electrons) and carries a positive charge of 1.60 x
10-19 c which is equal to the negative charge found on an electron . Protons gives a
positive charge to the nucleus of an atom. In a neutral atom, number of electrons and
protons are equal.
Atomic Number
Number of protons in a nucleus is called the atomic number (Z)
Neutrons
Chadwick discovered an uncharged particles in the nucleus to account for the missing
mass of atoms after considering number of electrons and protons. Chadwick observed a
particle in the nucleus of the same mass as a proton( 9Be + 4He --> 12C + 1n), but without
a charge in a nuclear reaction. His experiment led to the discovery of neutron. Neutrons
are sub-atomic and sub-nuclear particles. A neutron has a mass of 1.67 x 10-24 g which
is equal to that of a proton, but it is neutral. Neutrons along with protons contribute to
the mass or bulk of the matter or atoms.
Isotopes
Isotope: Atoms of the same element with different masses or different number
of neutrons in the nucleus. E.g. 1H (hydrogen with one proton and one
electron)and 2H (deuterium with one proton, one neutron and one electron).
Isotopic symbol: Element symbol (X) indicating number of protons or atomic
number (Z) written as left subscript, and mass number (A), total of number of
protons and neutrons written as left superscript.
A
12
X
E.g carbon-12: C
Z
6
or simply written as 12C (carbon-12) because once atomic symbol is known
atomic number is known from the periodic table.
Average Atomic Masses

We average the masses of isotopes using their masses and relative abundances to
give the average atomic mass of an element.

Naturally occurring Cl consists of 75.53% 35Cl (34.969 amu) and 24.47% 37Cl
(36.966 amu).

The average mass of C1 is
(0.7553)(34.969 amu) + (0.2447)(36.966 amu) = 35.46 amu
(Please see the periodic table!)

Atomic weight (AW) is also known as average atomic mass.

Atomic weights are listed on the periodic table.
Light and Atomic Structure
The process of exciting an atom, involves adding energy to the atom. This can be done
in a variety of ways. Simply heating a sample of an element up in an open flame will
excite electrons. Passing electricity through a sample of an element will excite
electrons. The colored lights observed when sky rockets explode are a result of burning
gunpowder exciting electrons within atoms of elements packed with the gun powder.
Flame colors
Fire Works
The Bohr Atom
Following diagrams to illustrate the Bohr model used to explain emission spectrum of
an atom.
Energy is absorbed when an electron goes from a lower (n) to a higher (n) level [e.g.
n=1 to n=3]
Energy is emitted when an electron goes from a higher (n) to a lower (n) level [e.g. n=5
to n=1]
• The lowest energy state is known as the ground state
• Higher states are known as excited states
the number n (n = 1, 2, 3, 4) called is called quantum number.
Hydrogen Spectrum
Modern Atomic Theory
Bohr's model of the atom is important because it introduced the concept of the
quantum in explaining atomic properties. However, Bohr's model ultimately needed
revision because it failed to explain the nature of atoms more complicated than
hydrogen. It took roughly another decade before a new more complete atomic theory
was developed - the modern atomic theory.
Louis de Brogllie introduced the wave/particle duality of matter (1921)
Heisenberg Uncertainty Principle: It is impossible to know simultaneously how fast
an electron is moving and its position with certainty.
Traditional (classical) physics had assumed that particles were particles and waves were
waves. However, de Broglie suggested that particles could sometimes behave as waves
and waves could sometimes behave as particles - the wave/particle duality of nature.
These electron clouds are denser in some regions than others. The electron density is
proportional to the probability of finding the electron at any point in time. According to
modern atomic theory electrons are shown as probabilities the space that an electron can
be found is called an atomic orbital. We will talk about different wave shapes of
atomic orbitals s, p d, and f later.
Electromagnetic Radiation and Its Effects on Our Everyday
Lives
The Electromagnetic Radiation Spectrum, or EMR Spectrum, is a full
"rainbow" of all colors of light, both visible and non-visible. To a large degree it
is a theoretical spectrum, in that no real system is capable of producing all colors
of EMR. The EMR Spectrum is generally broken down into three regions-visible, ultraviolet and infrared. By tradition, the spectrum is shown with the
colors in order of increasing wavelengths, decreasing frequencies and decreasing
energy values.
wave equation:
c = 
c = velocity of wave
 = frequency of wave
 = wavelength
Radio Waves: I like these waves; they are used for radio broadcasts, amateur
radio, television, and mobile phones. Different parts of the radio spectrum have
been allocated to the various services. Radio waves have a much longer
wavelength that light waves and less energy. The longest waves are several
kilometers in length. The shortest ones are only millimeters long.
Microwaves: Microwaves have such a short wavelength that they are very
easily absorbed by water. This is why they are used in microwave ovens. What
happens is that when the water in your dinner absorbs the microwaves, the
energy of the microwaves is converted into heat: it makes the water molecules
vibrate and rotate faster. Some people are frightened that the radio waves
coming out of their mobile phones are short enough to cook their brains. A fear
that is not confirmed by experiments
Infra Red: These radio/light waves have a very short wavelength; their
wavelength is longer than visible light. Infra-red can be detected by special
infra-red film. If the police shine an infra-red light on you they will be able to
take a picture of you at night using infra-red film: you will not know that they
have taken your photo. You have been warned!!! Animals like the pit-viper
have infra-red detectors so that they can find their prey in the dark.
The Visible Spectrum: Orange, Yellow, Green, Blue, Indigo, and Violet. are
the colors of the visible spectrum. We cannot see Infra-red, but we can feel it
warm our skin when we sit in the sun. Infra-red has a longer wavelength (less
energy) than Red light. We cannot see Ultra-violet light, but we feel our skin has
been burnt by the sun if we were in the sun too long yesterday. It is the Ultraviolet which is thought to cause skin cancer. UV light has a shorter wavelength
(more energy) than visible light. Amount of UV light we get is affected by
ozone layer.
Ultra Violet: These waves have very high energy and very short wave lengths;
shorter than visible light. Some animals like honey bees can see ultra-violet
light. Some plants have white flowers, at least you think that they are all white,
but they may appear to be different colors to a honey bee because of the
amounts of ultra-violet light which they reflect.
X-Rays: X-Rays have so much energy and such a short wavelength and higher
energy that they can go right through you. However, they cannot get through
bone as easily as they can get through muscle. This is because your bones
contain so much Calcium. If you have never had an X-Ray, try jumping off a
high wall and breaking a bone. The doctors will soon have you fixed: they take
an X-Ray to see which bones have been broken. If you break the bone in enough
places they will have to use steel bolts to fix you up. Long exposure to X-ray
known to cause cancer. X-Ray can also be used to find other problems in your
body. If the doctors want to look for ulcers in your guts, they can give you a
short lived isotope of barium meal. Like calcium in the bone, the barium absorbs
X-Rays so the doctors can look at parts of your guts and find your ulcers.
Gamma Rays: These are nasty ones produced in nuclear reactions. They have
very high energy and will even go through metals. So they can be used for
finding tiny cracks in metals. You cannot see the hairline cracks in an airplane
wing with the naked eye. If the plane has been thoroughly checked you should
be safe. Some radioactive materials produce gamma rays. Gamma rays and XRays can cause cancer, but gamma rays can also be used to destroy cancer cells:
this is radio-therapy.
CHEM 120 Homework Chapter 1
1) Which the following statement about scientific method is false
a) Scientific method is a procedure followed by scientist around the world.
b) It usually starts with an experiment, looking for patterns in the data, hypotheses, theory
then further experiments.
c) Scientific Theory will never become obsolete.
d) Patterns in observations are summarized as a scientific law.
e) Further experiments confirm and sometimes modify or replace a theory.
2. An example of a chemical compound is,
a)
air
b) brass (an alloy)
c) granite (a rock)
d) table salt
3. An example of a physical change
a) Baking cake b) Burning fire wood c) Electrolysis of water d) Sublimation of dry ice.
4. Samples of seawater obtained from various locations vary in the weight percentage of
various elements present in the sample. This suggests that seawater is
a. mixture.
b. a compound. c. an element. d. a pure substance.
5. Which name/symbol combination is wrong?
a. Tungstun/W
b. Fluorine/F
c.Silicon/Si
d. Selenium/Sn e. Silver/Ag
6. Compounds are different from mixtures because compounds
a. are composed of two or more substances.
b. have compositions that may vary from substance to substance.
c. always have the same set of physical properties.
d. can be separated by physical means into simpler substances.
7. Which of the following is a chemical property of iron?
a. occurs in an abundance of 4.7% in the earth's crust
b. melts at 1535 oC
c. attracted to a magnet
d. rusts when exposed to moist air
8. Convert -78 oC to oF.
a. -108 oF
b. -61 oF
c. -195 oF
9. The correct answer to report for
2.481 x 12.74 + 0.27 is
2.69
a. 12.0202
b. 12.020
c. 12.02
d. -351 oF
d. 12.0
e. 12
10. How many cubic meters equal one cubic millimeter?
a. 109
b. 106
c. 10-2
d. 10-9
CHEM 120 Homework Chapter 2
1) Which of the following name/discovery/experiment combination is not correct
a) Rutherford/ nucleus of an atom/-particle scattering experiment from a gold foil.
b) Thomson/discovery of electron/ Cathode Ray Tubes(CRT).
c) Chadwick/discovery of neutrons/ bombardment of 9Be with -particles
(9Be + 4He --> 12C + 1n).
d) Robert Andrew Millikan/protons, positively charged particles/Oil drop experiment.
2) Most of the volume of an atom is occupied by the space required for the,
a) moving electrons
b) protons
c) nucleus
d) neutrons
3) All isotopes of a given element,
a) possess different masses.
c) have the same atomic number)
b) possess the same chemical properties.
d) all the above.
4) The 207Pb2+ ion contains
a) 82 protons, 126 neutrons, and 80 electrons.
c) 82 protons, 126 neutrons, and 126 electrons.
b) 82 protons, 125 neutrons, and 80 electrons
d) 126 protons, 82 neutrons, and 122 electrons.
5) An atom with 35 protons and 45 neutrons would be an isotope of the element,
a). rhodium, Rh
b) neon, Ne
c) mercury, Hg
d) bromine, Br
6) How many microliters are in 89.63 L?
a) 8.963 x 10-4 L b) 8.963 x 10-7 L
c) 8.963 x 104 L d) 8.963 x 107 L
e) 8.963 x 1010 L
7) A box has dimensions of 1.51 in. by 2.74 in. by 4.72 in. What is its volume in cubic centimeters?
a) 1.19 cm3
c) 7.69 cm3
e) 49.6 cm3
b) 0.841 cm3
d) 320. cm3
8) If the isotopic ratio of the two boron isotopes 10B (10.013 amu) and 11B (11.009 amu) has been
altered from the ratio found in nature and now contains 48.73% 10B, determine the atomic weight
of the sample of boron.
a)10.811 amu b) 10.013 amu
c) 11.009 amu d) 10.524 amu
e) 10.498 amu
9) . In Bohr's theory of the atom, what is the number n (n = 1, 2, 3, 4)
called?
a) Quantum number b) Atomic number c) Mass number d) orbital
10) Explain the meaning/consequences of Heisenberg's
Uncertainty Principle.
a) Electron have wave properties
b) It is impossible to know the exact position and momentum of a particle,
such as an electron in an atom.
c) Electron absorb energy and get excited to a higher energy level.
d) Certain nuclei decay and change to other types of elements
Sample Test Questions
Chapter 1, Chemistry: Methods and Measurement
1. Name the branch of science that involves the study of matter and the changes it
undergoes.
Ans. chemistry
2. What do we call anything that has mass and occupies space?
Ans. matter
3. What is meant by the term "scientific law"?
Ans. a summary of a large amount of scientific information
4. What word means the application of scientific principles to meeting human needs?
Ans. technology
5. What is a hypothesis?
Ans. an attempt to explain an observation in a common-sense way
6. When does a hypothesis attain the status of a theory?
Ans. when sufficient experiments have been performed to confirm that the hypothesis is
correct
7. What does the prefix "centi-" mean?
Ans. 10-2 or 0.01
8. 1 microgram is equivalent to how many grams?
-6
g
9. How many centimeters correspond to 15.68 kilometers?
6
cm
10. How many pounds are represented by 764.6 mg? [Use: 1 pound = 454 g]
-3
Ans. 1.684 x
pounds
11. A typical aspirin tablet contains 5.00 grains of pure aspirin analgesic
compound. The rest of the tablet is starch. How many aspirin tablets
can be made from 100.0 g of pure aspirin? [Use: 1.00 g = 15.4 grains]
Ans. 308
12. A patient weighs 146 pounds and is to receive a drug at a dosage of
45.0 mg per kg of body weight. The drug is supplied as a solution that
contains 25.0 mg of drug per mL of solution. How many mL of the drug
should the patient receive? [Use: 1 pound = 454 g]
Ans. 119 mL
13. If a person smokes 10.0 packs of cigarettes a week and each cigarette
contains 5.00 mg of tar, how many weeks will she have to smoke to
inhale 0.250 pounds of tar? [Use: 20 cigarettes = 1 pack and 1 pound =
454 g]
Ans. 114 weeks
14. The cost of a drug is 125 francs per gram. What is the cost in dollars
per ounce? [Use: $1 = 6.25 francs and 1 ounce = 28.4 g]
Ans. $568/ounce
15. If one atom of carbon-14 weighs 14.0 atomic mass units and one atomic
-24
grams, what is the mass of 250.0
atoms of carbon-14 in grams?
-21
g
16. A patient needs 0.300 g of a solid drug preparation per day. How many
10.0 mg tablets must be given to the patient per day?
Ans. 30.0 tablets
17. When the value of an experimental quantity or measurement (e.g. the
mass of a tablet) is reported, it should consist of two parts. What are they?
Ans. number and units
18. What do we call a basic quantity of mass, volume, time, etc.?
Ans. unit
19. What is defined as the difference between the true value and our
measurement of that value?
Ans. error
20. In one sentence, define the word "accuracy"
Ans. the agreement between the true value and the measured value of a
quantity
21. What do we call the degree of agreement between replicate measurements
of the same quantity?
Ans. precision
22. What do we call the degree of doubt in a single measurement?
Ans. uncertainty
23. Express the number 0.000327730 to three significant figures using
scientific notation.
-4
24. How many signifi
Ans. 5
4
?
25. Write the number 3,000 using scientific notation and the proper number
of significant digits.
3
26. Write the number 0.00946 using scientific notation and the proper
number of significant digits.
-3
27. Provide the answer to the following problem using scientific notation
and the proper number of significant digits:
-2
-4
)=?
-5
28. Provide the answer to the following problem using scientific notation
and the proper number of significant digits:
4
3
)=?
1
29. Provide the answer to the following problem using the proper number of
significant digits: 0.004 + 26.59 + 3.2 = ?
Ans. 29.8
30. What is meant by the word "mass"?
Ans. the quantity of matter in a sample measured comparing to a standard wieght
31. What instrument is used to measure the mass of an object?
Ans. a balance
32. What commonly used mass unit is approximately the same as the mass of
one hydrogen atom?
Ans. atomic mass unit
33. What experimental quantity gives the force resulting from the pull of
gravity upon an object?
Ans. weight
34. What is the basic unit of volume in the metric system?
Ans. liter
35. List four different forms of energy, other than kinetic energy or
potential energy
Ans. any four of heat, light, electrical, mechanical, chemical
36. What Fahrenheit temperature corresponds to -168.7◦C?
Ans. -271.7◦F
37. What form of energy is usually absorbed or liberated during chemical
reactions?
Ans. heat
38. In the swinging of a pendulum, what two forms of energy are constantly
being interconverted?
Ans. kinetic energy and potential energy
39. What kind of energy is stored or the result of position or composition?
Ans. potential energy
40. What do we call the experimental quantity which gives the number of
particles of a substance (or their mass) contained per unit volume?
Ans. concentration
41. The density of an object is the ratio of its _______ to its _______.
Ans. mass; volume
42. If we assume the density of blood is 1.060 g/mL, what is the mass of
6.56 pints of blood in grams? [Use: 1 L = 2.113 pints]
3
g
43. A solid that had a mass of 189.6 g was found to have the following
measurements: length = 9.80 cm; width = 46.6 mm; height = 0.111 m. What
is its density in g/mL?
Ans. 0.374 g/mL
44. When a large object fell into a "full" swimming pool, it displaced 32.0
gallons of water. The density of the object is 1.65 g/mL. What is the
object's weight in pounds? Remember that any fully immersed object
displaces its own volume. [Use: 454 g = 1 pound; 1 L = 1.06 quarts]
Ans. 439 pounds
45. What is the branch of chemistry that is being applied in measuring the
concentration of an air pollutant?
A. analytical chemistry
B. biochemistry
C. inorganic chemistry
D. organic chemistry
E. physical chemistry
Ans. A
46. What do we call a statement of observed behavior for which no
exceptions have been found?
A. hypothesis
B. theory
C. law
D. model
E. result
Ans. C
47. The speed of light is 186,000 miles per second. What is its speed in
centimeters per second?
[Use: 5280 feet = 1 mile; 12 inches = 1 foot; 2.54 cm = 1 inch]
A. 3.01 x
11
cm/s
cm/s
12
cm/s
11
cm/s
10
cm/s
10
Ans. E
48. 1 kilometer equals how many millimeters?
A. 10-6 B. 10-3 C. 103 D. 104 E. 106
Ans. E
49. Round off 0.052018 to three significant figures.
A. 0.05
B. 0.052
C. 0.0520
D. 0.05201
E. 0.05202
Ans. C
50. Select the answer which best expresses the result of the following
calculation: 1.86 + 246.4 – 79.9208 = ?
A. 168
B. 168.3
C. 168.34
D. 168.339
E. 168.3392
Ans. B
51. The appropriate number of significant figures to be used in expressing
the result of 51.6 / 3.1416 is
A. 1 B. 2 C. 3 D. 4 E. 5
Ans. C
52. What Celsius temperature corresponds to -4.6◦F?
A. –20◦C
B. –20.3◦C
C. –23.0◦C
D. –10.9◦C
E. –68.4◦C
Ans. B
53. What Fahrenheit temperature corresponds to -40.0◦C?
A. –8◦F
B. 16.8◦F
C. –36.9◦F
D. –40.0◦F
E. –1.94◦F
Ans. D
54. What Kelvin temperature corresponds to 98.6◦F?
A. 310K
B. 310.0K
C. 31.00K
D. 132.0K
E. 199K
Ans. B
55. Which temperature scale does not use a degree sign?
A. Celsius
B. Kelvin
C. Centigrade
D. Fahrenheit
E. Absolute zero
Ans. B
56. If the density of carbon tetrachloride is 1.59 g/mL, what is the
volume, in L, of 4.21 kg of carbon tetrachloride?
A. 0.149 L
B. 0.378 L
C. 2.65 L
D. 6.69 L
E. 6690 L
Ans. C
57. What is the specific gravity of an object that weighs 13.35 g and has a volume of
25.00 mL? The density of water under the same conditions is
0.980 g/mL.
A. 1.335
B. 0.545 g/mL
C. 1.335 mL
D. 0.545
E. 0.980
Ans. D
58. T F Organic chemistry is the study of those chemical processes that
are found in living systems.
Ans. F
59. T F Hypotheses are not acceptable in the scientific method.
Ans. F
60. T F In the scientific method, a hypothesis carries more weight than a
theory.
Ans. F
61. T F Each piece of data is the individual result of a single
measurement.
Ans. T
62. T F The presence of some error is a natural consequence of any
measurement.
Ans. T
63. T F Errors which are scattered equally above and below the correct
value are called systematic errors.
Ans. F
64. T F It is possible for a set of measurements to be precise yet
inaccurate.
Ans. T
65. T F The number 0.068 has 3 significant figures.
Ans. F
66. T F The terms mass and weight are identical.
Ans. F
67. T F Mass is the force resulting from the pull of gravity upon an
object.
Ans. F
68. T F Equal masses of glass and steel at the same temperature will
generally have different heat energies.
Ans. T
69. T F Energy may be defined as the heat content of an object.
Ans. F
70. T F One Calorie is the amount of energy needed to raise the
temperature of one gram of water one degree Celsius.
Ans. T
71. T F Density and specific gravity can be expressed in the same units.
Ans. F
Chapter 2, The Composition and Structure of the Atom
1. In which state does matter have a definite shape and volume?
Ans. solid
2. In which state of matter are forces between particles least dominant?
Ans. gas
3. What kind of change does not alter the composition or identity of the
substance undergoing the change?
Ans. physical
4. Conversion of ice to liquid water or liquid water to steam is an
example of what kind of change?
Ans. physical
5. In one sentence, explain clearly what is meant by a "chemical change".
Ans. A chemical change results in different substance(s) from those
initially present.
6. What type of change is represented by the decay of a fallen tree?
Ans. chemical
7. Give an example of a chemical property of iron metal.
Ans. its tendency to form rust when exposed to the atmosphere
8. What do we call the starting and final materials in a chemical
reaction?
Ans. reactants and products
9. What type of property of matter is independent of the quantity of the
substance?
Ans. intensive
10. What word is used to describe properties of a substance that depend on
the quantity of substance? Give two examples of such properties.
Ans. extensive; mass, volume and others
11. What is meant in chemistry by the term "pure substance"?
Ans. The substance consists of only one component; it has a fixed
composition throughout.
12. What are two kinds of pure substance?
Ans. elements and compounds
13. Explain what is meant in chemistry by the word "compound".
Ans. a pure substance consisting of two or more elements chemically combined
in a definite ratio
14. What are the two classes of mixtures?
Ans. homogeneous and heterogeneous
15. Give an example of a heterogeneous mixture.
Ans. concrete, salt and pepper, smoky air and others
16. What kind of mixture is a solution of alcohol and water?
Ans. homogeneous
17. Which of the following terms are appropriate in describing an apple?
pure substance; element; compound; mixture; homogeneous; heterogeneous
Ans. mixture; heterogeneous
18. List the three primary particles found in an atom.
Ans. proton, electron, and neutron
19. What quantity does the mass number minus the atomic number represent?
Ans. number of neutrons
20. In a neutral atom, what number of particles is equal to the number of
protons?
Ans. number of electrons
21. Given that He helium has an isotope predict the number of electrons 42,
in a helium atom.
Ans. 2
22. How many neutrons are present in an atom of the isotope 3 7
Li?
Ans. 4
23. What is the value of the mass number in the isotope 53
131I?
Ans. 131
24. What is the term for atoms of the same element having different masses
due to a different number of neutrons?
Ans. isotopes
25. What is the process by which certain isotopes emit particles and
release large amounts of energy?
Ans. radioactive decay
26. What do we call electrically charged particles that result from the
gain of one or more electrons by a parent atom?
Ans. anions
27. Dalton proposed that all atoms of an element have identical properties.
Briefly, explain why this proposal is invalid.
Ans. Isotopes of an element have different properties, particularly their
mass.
28. J. J. Thomson in 1897 announced that cathode rays consisted of a stream
of _______.
Ans. electrons
29. In one sentence, state Rutherford's important contribution to our
knowledge of atomic structure.
Ans. He concluded that atoms have a small, heavy, positively charged nucleus
surrounded largely by empty space, occupied by electrons.
General, Organic, and Biochemistry, 3/e Page 9
30. A more general term for "light" is _______ _______ (two words).
Ans. electromagnetic radiation
31. What word means the study of wavelengths of light emitted and absorbed
by atoms?
Ans. spectroscopy
32. In Bohr's model of the hydrogen atom, what is an "orbit"?
Ans. An orbit is a circular path followed by the electron.
33. In Bohr's theory of the atom, what is the number n (n = 1, 2, 3, 4)
called?
Ans. quantum number
34. In one sentence, explain the meaning/consequences of Heisenberg's
Uncertainty Principle.
Ans. It is impossible to know the exact position and momentum of a particle,
such as an electron in an atom.
35. In modern atomic theory, Bohr's orbits are replaced by atomic orbitals.
What is an atomic orbital?
Ans. It is a region of space where an electron is likely to be found; or, an
electron "cloud".
36. Which state of matter has no definite shape or volume?
A. liquid B. solid C. vapor D. steam E. gas
Ans. E
37. Which of the following is NOT a physical property of matter?
A. odor
B. compressibility
C. flash point
D. melting point
E. color
Ans. C
38. What kind of change does NOT alter the composition or identity of the
substance undergoing the change?
A. molecular
B. endothermic
C. exothermic
D. physical
E. chemical
Ans. D
39. What kind of change always results in the formation of new materials?
A. molecular
B. exothermic
C. endothermic
D. physical
E. chemical
Ans. E
40. Which of the following is a chemical property?
A. flammability
B. color
C. hardness
D. odor
E. taste
Ans. A
41. Which one of the following is an example of an extensive property?
A. density
B. specific gravity
C. mass
D. hardness
E. boiling temperature
Ans. C
42. Which one of the following is an example of a pure substance?
A. ethyl alcohol
B. sugar water
C. salt and pepper
D. milk
E. sand
Ans. A
43. Air is a/an
A. element
B. compound
C. mixture
D. molecule
E. pure substance
Ans. C
44. What type of mixture is represented by a collection of salt and pepper?
A. atoms
B. molecules
C. solution
D. heterogeneous
E. homogeneous
Ans. D
45. What kind(s) of particles can be found in the nucleus of an atom?
A. protons
B. neutrons
C. electrons
D. protons and electrons
E. protons and neutrons
Ans. E
46. The total mass of the protons in any neutral atom is about _______
times the total mass of electrons in the atom.
A. 0.0005 B. 0.3 C. 1 D. 2 E. 2000
Ans. E
47. What is the quantity represented by the mass number (Z) minus the atomic
number?
A. number of atoms
B. number of neutrons
C. number of electrons
D. number of protons
E. number of particles in the nucleus
Ans. B
48. Which isotope of hydrogen has two neutrons?
A. hydrogen-1
B. hydrogen-2
C. hydrogen-3
D. deuterium
E. H2
Ans. C
49. Which of the following accounts for the fact that chlorine has an
atomic mass of 35.45 amu rather than a whole number?
A. isotopes
B. electrons
C. protons
D. radioactivity
E. isomers
Ans. A
50. Who discovered that cathode rays consist of a stream of negative
particles, electrons?
A. Crookes
B. Thomson
C. Geiger
D. Rutherford
E. Bohr
51. Who discovered the existence of the atomic nucleus?
A. Crookes
B. Thomson
C. Geiger
D. Rutherford
E. Bohr
Ans. D
52. In Rutherford &Geiger's experiment which led to the discovery of the atomic
nucleus, what type of particle or ray was fired at the gold foil
target?
A. alpha
B. beta
C. gamma
D. neutrons
E. cathode rays
Ans. A
53. Which of the following statements relating to Bohr's model of the
hydrogen atom, is incorrect?
A. The lowest energy orbit has quantum number n = 1
B. The highest energy orbits are furthest from the nucleus
C. In a transition from the n = 3 to the n = 1 level, light is emitted
D. Energy differences between energy levels can be calculated from the
wavelengths of the light absorbed or emitted
E. The greater the energy difference between two levels, the longer the
wavelength of the light absorbed or emitted
Ans. E
54. Who proposed that electrons could behave like waves, as well as like
particles?
A. Thomson
B. Rutherford
C. Bohr
D. de Broglie
E. Heisenberg
Ans. D
55. T F The term "solution" refers only to homogeneous mixtures of
liquids.
Ans. F
56. T F In the cal Calcium atom represented by the symbol , 60Ca, calcium-60, there are
40 protons, 20 neutrons and 20 electrons.
Ans. T
57. T F According to modern atomic theory, all atoms of a particular element have
identical physical properties.
Ans. F
58. T F According to modern atomic theory, an atom cannot be created, divided,
destroyed or converted to any other type of atom.
Ans. F
59. T F The atomic number of an ion tells us the number of protons that are present.
Ans. T
60. T F If an atom gains one electron, it becomes a cation.
Ans. F
61. T F The first experimentally based theory of atomic structure was
proposed by John Dalton.
Ans. T
62. T F J. J. Thomson was the first to state that an atom is mostly empty space.
Ans. F
63. T F Short wavelengths of electromagnetic radiation have more energy than long
wavelengths.
Ans. T
64. T F Rutherford was the first to use the term "orbit" to explain the fixed energy levels
of electrons.
Ans. F
65. T F Niels Bohr developed a theory which accounted for the lines in the visible
region of the hydrogen spectrum.
Ans. T