THHS.technologyintheclassroom.

Mr. K.
Vocabulary Skill
Word Scramble
Fill in the blanks in the paragraphs below with the correct words by unscrambling the letters to the left of the blanks.
A HORYET________________ is the most logical explanation of events that occur innature. This explanation often results from the use of the SCINIFICTE DETMOH
___________________________________ . First, a problem must be stated. Then, after
gathering information, a SOPTYHESHI_______________________ , or logical solution, isformed. The solution is then tested in an experiment. The factor that is tested in an
experiment is called the BAVALIER____________________ . This factor is not contained
in the RONTLOC PESTU_____________________________ . Sometimes principles of
MECHRYIST_______________________ , the study of how substances change and combine, are involved in experiments. At other times, principles of SHICYPS
_______________ are involved.
The CERTIM MYSTES_____________________________ is used for all measure‑
ments in science. Some basic units in this system are the TILRE_____________ for
volume, the LIKRAGMO__________________ for mass, the degree ELUCSIS
_______________ for temperature, and the TREEM____________ for length. Some‑
times the TESYNID_________________ of a substance, or its mass per unit volumemust be measured. If liquid volume is being measured, the reading is taken at the
SISMUNEC_____________________ If dimensional analysis is needed, a NOSVIERONC
TOFRAC______________________ must be used.


___________________________________________________________________
Using Science Skills: Making observations

Water Education Activity
Melting Ice and Freezing Water
1. Place several ice cubes and a thermometer in a glass with a little water and wait
several minutes. What is the temperature as the ice melts?______________ This temperature is the melting point of ice.
2. Place a glass of water, with a thermometer in it, in the freezer. Observe the temperature every 10 or 15 minutes. What is the temperature of the water when
it begins to freeze?______________ This temperature is the freezing point of water.
3. How does the melting point of ice compare with the freezing point of water?
Metric Conversions
Units Prefixes
length — meter (m) kilo (k) — 1000
volume — liter (L) centi (c) — 1/100
mass — gram (g) milli (m) — 1/1000
1. 1 km =_____________
2. 1 m =______________ cm
3. 1 cm =_____________ mm
4. 25 km =______________ m =_____________ cm =______________ mm
5. 5000 mL =_____________ cL =______________ L =______________ kL
6. 350 L =_____________ mL
7. 55 kL =_____________ cL =______________ mL
8. 31,768 mg =_____________ g =______________ kg
9. 4351 cg =_____________ g=___________ kg
10. 166 mg =_____________ cg =______________ g =______________ kg
Name ___________________________________ Class _______________________ Date __________________

___________________________________________________________________
Using Science Skills: Solving problems
Scientific Measurements

1. A rectangular block of wood is 11.3 cm long, 7.20 cm wide, and 10.4 cm deep. It has a mass of 825 g.
a. Find the area of the smallest face.
b. Find the volume of the block.
c. Find its density.
2. What is the mass of a 50-cm3 quantity of water?
3. What is the density of a 200-gram mass whose volume is 300 cm3?
4. If the density of sea water is 1.04 g/cm3, what is the mass of 1000 cm3 (1 liter) of sea water?
5. If a 540-gram mass has a density of 2.7 g/cm3, what is its volume?

___________________________________________________________________

Using Science Skills: Applying definitions

Little Millie Metric

When making measurements, it is important to use the correct metric unit. Read the story below and fill in the blanks with the right unit. Choose the unit from the following list. Each unit will be used only once.
milliliter liter gram kilogram millimeter
centimeter meter kilometer degrees Celsius
It was a beautiful day. The temperature was a mild 27_________________ . LittleMillie Metric was packing a lunch basket to take to Grandma's house. She carefully
poured 500__________________ of homemade lemonade into a bottle, which she put in
the basket. Then she placed 0.5_________________ of cheese, 200_________________ ofroast beef, and several large chocolate chip cookies in the basket. She strapped the basket on the back of her moped and took a quick check of the gas tank. There were
several________________ of gasoline left.
After traveling a little more than 5_________________ , Millie discovered she had lost her way. At the next corner she spotted a very hairy character leaning against a
lamppost. He looked to be only 11/2______________ tall. He had a wolfish grin and dark, piercing eyes. He introduced himself as Mr. W and offered to help Millie.
Mr. W's directions turned out to be the long way to Grandma's house. So Millie arrived an hour late. She let herself in and found Grandma in bed. Grandma did not look well. Had she always been so hairy? And those ears! "My goodness," Millie gasped, "what big ears you have, Grandma—at least 15__________________________________________________ ." To which Grandma replied, "The better to hear you, my dear." Then Millie spotted the teeth.
"My, Grandma, what big teeth you have—longer than 25__________________ ." This last observation made Grandma very angry. She jumped out of bed and snarled, "The better to eat you!" Now Millie realized that this was not Grandma, but the hairy Mr. W! Unfortunately for Mr. W, Millie held a black belt in karate. In just a few seconds, she took care of him!
Millie found Grandma in the closet. When they finally sat down to eat lunch,Millie couldn't help thinking that maybe next time she'd just have lunch delivered.
B. Now fill in the blank next to each unit with its correct abbreviation.
1. milliliter
2. centimeter
3. liter
4. meter
5. gram
6. kilometer
7. kilogram
8. millimeter
9. degrees Celsius ___


GAS LAWS


PROBLEMS
1. What Celsius temperature is needed to double the volume of a gas if the original temperature is 30° C? (Pressure remains constant.)
2. What pressure is required to convert 200.0 ml of a gas at STP (standard temper­ature and pressure) to 400.0 ml without any temperature change?
3. What volume will a gas occupy if 56 liters at 273° C are cooled to 0° C without a pressure change?
4. What pressure will be necessary to compress 2.0 liters of a gas at I atmosphere pressure to a volume of 0.50 liter if the temperature remains constant?
5. 3.40 liters of oxygen gas were collected at 37° C and .50 atm pressure. At 1.0 atm, the volume became 6.80 liters. What is the new Celsius temperature of the gas?
6. What volume will a gas occupy if 60.0 ml at –245.7° C and 1.75 atm pressure are adjusted to STP?


7. 48 liters of carbon dioxide gas are collected at STP. What will the volume of the gas be at 273° C and 1.75 atm?
8. 150 liters of a gas are collected at 27° C and 2.25 atm pressure. When the pres­sure is decreased to 1.5 atm, the volume becomes 200 liters. What is the new Celsius temperature?
9. 22.4 liters of a gas are collected at 1.25 atm pressure and 87° C. What pres­sure is required to reduce the volume to 5.6 liters when the temperature is reduced to –33° C?
10. 11.2 liters of a certain gas is collected at STP. What will the volume of the gas be at 273° C and 1.9 atm pressure?
11. Calculate the pressure required to reduce the volume of a gas from 350 milli­liters at 2.0 atm pressure to 175 milliliters with the temperature constant.
12. 300 liters of a gas are collected at 2.75 atm pressure and 7.0° C. If the pressure is increased to 3.25 atm and the temperature is decreased to –33° C, what is the new volume?
13. 24 liters of a gas are collected at 1.75 atm pressure and 27° C. Calculate the vol­ume of the gas if the pressure is increased to 3.25 atm and the temperature is reduced to –73° C.
14. 600 ml of a gas is collected at 2.80 atm pressure and at 37° C. When the pressure increases to 4.0 atm, the volume increases to 840 ml. Determine the new Cel­sius temperature.
15. Calculate the pressure required to reduce the volume of a gas from 400 milli­liters at 1.75 atm to 200 milliliters with the temperature constant.



WHO AM I?

The following chart contains clues about the identities of some of the elements in the periodic table.


1. I have very good ability to conduct electricity. I am never found
alone in nature. When I combine with other elements, I usually give
up my one valence electron. I am the only element in my group with
a one-letter symbol.


2. I do not conduct electricity and am usually found in the gaseous
state. I do not bond well with other elements. I can be found in
some bulbs used in signs.


3. I am a gas, but I combine very easily with many other elements. I
usually form ionic bonds. I frequently form a –1 ion in those ionic
bonds. I am the lightest element in my group.


4. I am a very tough, durable element. I can give up two electrons,
but I sometimes give up more than two when bonding. I am the
main element found in steel.


5. I am never found alone or unbonded in nature. I most commonly
form a +2 ion when bonding. I have the second highest number of
protons in my family.


6. In my family the elements are all metals except for me. I have
three valence electrons.


7. Although I am in a family of nonmetals, I am found as a solid. If I
combine with calcium, two atoms of me but only one atom of cal‑
cium are required.


8. I usually form covalent bonds. I have five valence electrons. I
have the highest atomic mass in my group.


9. In my family there are nonmetals, metalloids, and metals. I have
the same number of protons as the sum of the protons in the two
elements directly above me in the periodic table.


10. Elements in my family usually form covalent bonds. We have
two fewer valence electrons than noble gases. I am almost twice as
heavy as the lightest element in my group.


HEAT

Fill in the letter of the answer that best completes each statement.
1. Which scientist investigated the relationship between objects in motion and heat?
a. J. J. Thomson b. Lord Kelvin c. James Joule d. Gabriel Fahrenheit
2. Which of the following is not a method of heat transfer?
a. conduction b. ventilation c. convection d. radiation
3. When warmer molecules collide with cooler molecules,
a. energy is transferred from cooler molecules.
b. energy is transferred to cooler molecules.
c. a phase change takes place.
d. an insulator is formed.
4. As molecules become warmer, which of the following does not happen?
a. Molecules become less dense.
b. Convection currents form.
c. Molecules are less closely packed.
d. Cold is transferred away from the surroundings.
5. A temperature of —10°C corresponds to what temperature on the Kelvin scale? a. 263°K b. 283°K c. 0°K d. —10°K
6. Of the following solution temperatures, select the one that would have the highest average kinetic energy.
a. 10°C b. 20°C c. 300°K d. 290°K
7. The ability of a substance to absorb heat energy is called
a. calorie. b. specific heat. c. temperature. d. conduction.
8. An increase in temperature indicates
a. slower moving molecules. c. a rise in specific heat.
b. a phase change. d. the addition of heat.
9. Which of the following would not affect a temperature change in a substance?
a. mass of the substance c. addition of calories
b. heat of fusion d. specific heat
10. Which of the following would require the greatest number of calories?
a. heating 100 g of water from 20°C to 40°C
b. heating 1 g of water from 25°C to 100°C
c. heating 100 g of water from 42°C to 43°C
d. heating 10 g of water from 25°C to 100°C
11. The amount of heat needed to change a substance from the solid phase to the liquid phase is called
a. specific heat. c. heat of fusion.
b. heat of vaporization. d. melting point.
12. The amount of heat needed to change a substance from the liquid phase to the gas phase is called
a. specific heat. c. heat of vaporization.
b. heat of fusion. d. boiling point.
13. During a phase change,
a. extreme temperature changes occur. c. calories remain constant.
b. temperature changes occur slowly. d. temperature does not change.
14. The specific heat of ice is
a. 540 cal/g. b. 0.5 cal/g C°. c. 80 cal/g. d. 1500 cal.
744 Physical Science © 1988 Prentice-Hall, Inc.15. If 200 g of water in a calorimeter increases temperature by 5 C°, how many calories have been absorbed?
a. 200 b. 500 c. 1000 d. 40
16. How much heat is needed to melt 100 g of ice at 0°C?
a. 8000 cal b. 540 cal c. 7200 cal d. 100 cal
17. Which of the following is not part of the process of melting ice cubes?
a. Heat is absorbed.
b. Forces of attraction are overcome.
c. The temperature of the ice rises.
d. Average kinetic energy remains the same.
18. As heat energy is added to a solid, the
a. temperature always increases.
b. specific heat increases.
c. kinetic energy of the molecules increases.
d. molecules slow down.
19. The rise of fluid in a thermometer is a result of
a. a phase change. c. the change in volume of water.
b. a calorimeter. d. thermal expansion.
20. When water freezes, it
a. gives off 540 calories. c. expands.
b. contracts. d. becomes more dense.
21. A device that controls temperature is called a
a. calorimeter. b. thermostat. c. thermometer. d. Kelvin scale.
22. Scientists define work as
a. heat. c. force causing an object to move.
b. a phase change. d. energy.
23. Energy contained in a substance is called
a. internal energy. b. temperature. c. heat. d. work.
24. When a substance does work, it
a. lowers its temperature. c. gains internal energy.
b. raises its temperature. d. loses internal energy.
25. Thermal expansion can best be explained in terms of
a. an increase in density. c. bimetallic strips.
b. the kinetic energy of molecules. d. the loss of internal energy.


Atomic Structure


I. A neutral atom of argon contains (how many?)__________ electrons.
2. When ail of the electrons in an atom are in the lowest available energy levels, the atom is in the
3. A neutral atom containing ten negative charges is the element______
4. As an electron's distance from the nucleus increases, its energy content
5. The position of an atom in the Periodic Table and the number of its electrons are both given by its
6. The maximum number of valence electrons possible in any outermost shell is
7. The number of valence electrons increases with increase in________ from the beginning to the end of Period 2.
8. The spectral lines of hydrogen correspond to________ energy levels.
9. The number of valence electrons in helium is_________ but the numberof valence electrons in all other inert gases is ______

10. Atoms emit radiant energy in quanta when _____electrons fall to a lower energy level.
11. The M shell corresponds to principal quantum level
12. If an electron has absorbed energy and has shifted to a higher energy level, the electron is said to be________________
13. The M shell attains its maximum number of valence electrons with the element -----and its maximum number of electrons with the element-----
14. The average region through which an electron moves is a(an) ------
15. In any p sublevel, there are (how many?) ------- p orbitals.
16. The maximum number of electrons possible in any p sublevel is ------
17. The number of orbitals in the outermost shell of neon is -------
18. A sublevel designated 4p means a p sublevel in the ----principal energy level.
19. A 3d orbital has (more, less)___ energy than a 3p orbital.
20. The only two kinds of orbitals which may occur in the outermost shell are the____________ and___________
21. In lithium, the orbital of highest relative energy is the ------
22. If n=3, the total number of orbitals is -----; the number of d orbi‑
tals is __________
23. The element having atomic number 36 is -----. The number of
electrons in its 3d sublevel is__
24. No more than -----electrons can be accommodated in an orbital.
25 -----orbital in a sublevel must contain one electron before anyorbital in the sublevel can contain two electrons.
26. The electron configuration Is22S22p63s23p64S1 is that of the element----
27. The electron configuration 1s22s22p63s23p63d104s24p1 is that of the element____________ , which contains (how many?) ----electron(s) in its outermost orbital.
28. An atom is chemically -----when all of the orbitals in the outer­most shell are completely filled.
29. Elements may react to form ions developing electron configurations like those of the----
Atomic Structure
MULTIPLE-CHOICE QUESTIONS
1. Within the Periodic Table, the minimum number of electrons that occurs in the K shell is (1) 1, (2) 2, (3) 4, (4) 8.
2. The valence shell of a neutral atom in the ground state is (1) the principal energy level that is farthest from the nucleus, (2) the principal energy level that is nearest the nucleus, (3) the electron configuration of an inert gas, (4) the p sub­level that is farthest from the nucleus.
3. An excited electron of hydrogen emits the least energy when it drops from quantum levels (1) 4 to 2, (2) 5 to 1, (3) 2 to 1, (4) 5 to 2.
4. Bohr postulated that electrons could only exist in certain paths around the nucleus which he termed (1) orbits, (2) orbitals, (3) quanta, (4) spectral lines.
5. According to Bohr, the more energy an electron possesses, (1) the more stable it is, (2) the smaller the size of the quantum emitted, (3) the farther it is from the nucleus, (4) the closer it is to the nucleus.Spectral lines of the element are caused by (1) electrons falling to lower energy


ME TO WE

How much does it cost to make a differene in someone life? As 'little as ten cents "
That's the remarkable conclusion of a 1972 study which set out to see how a little bit of gratitude can af­fect the way we behave.
Researchers in San Francisco and Philadelphia placed a dime in the coin return slot of shopping mall payphones too see how people would react to finding it. Moments after an unsuspecting shopper dis­covered the dime, a man dropped a folder full of papers nearby.
Happy with their find, 84 per cent of those with the new dime in their pocket stopped to help. But of all the other people walking by, less than five per cent lent a hand.

This simple study points to a pow­erful equation that psychologists have been talking about for years: gratitude leads to happiness, which leads to compassion.
.
People who, regularly feel thankful even for 'things as small as a newfound dime—enjoy higher life satisfaction and feel more positive emotions such as happiness, hope, and vitality. In turn, they forge stronger relationships, are healthier and more likely to be concerned about the welfare of others.That's the power of gratitude
But let's put science aside for a moment. Think back to a time in your life when someone did some­thing nice for you. Maybe a neigh­bor shoveled your driveway after a blistering storm or your spouse made you dinner after a long day at work. How did you feel? It prob­ably brightened your day, made you thankful, and even encouraged you to return the favor.
That is the core of the life philos­ophy we call Me to We. By using gratitude to reach out to others, we can, live more fulfilling lives, feel a deeper sense a happiness build stronger communities, and work to­ward a greater shared humanity.
It sounds simple, but true grati­tude is a skill. It involves training our minds to notice, savor, and re­member the positive, to develop a sense of wonder and a greater ap­preciation for the world.
Living Me to We uses that grati­tude as the fuel to focus less on "me" and more on "we"—our com­munities, our nation, and our world. It creates a path to happiness that is less about collecting material goods and more about making positive contributions to the world.
Reaching out to others takes cour­age, but the payoffs are endless

In our hectic world, we rarely tab the time to pause and think abou how lucky we are. But consider that if you bought a coffee this morning you already spent the day's salary' of the 3 billion people in the work who earn less than $2 a day. Even by reading this you are luckier that the 800 million people who are illiterate and the 120 million children who are too poor to go to school.
Most of us in North America are pretty fortunate.
So take a moment to think about what you are truly thankful for. caring family? A night out with your friends? Freedom to speak your mind? Make a list, keep it updated and use your appreciation for all that you have to make a difference in someone else's life.
You might ask what two brothers from the suburbs of Toronto know about personal fulfillment. The truth is we are just the messengers. Our inspiration comes from the amazing people we have met during a decade of international charity work. Fron playing soccer with street kids in Brazil to sharing moments of reflection with Mother Teresa and the Dalai Lama, we have learned a valuable lesson:

Being thankful for what we have and sharing it with others is the true path to happiness. In other words, live from Me to We.

This philosophy is made up of a series of small decisions that we all make every day: What do I want out of life? What is really important to me? What makes me truly happy? Will I stop to help? Will I give back? These require careful thought and determined action.
For some, living Me to We sim­ply means getting to know co-work­ers and neighbors or spending more quality time with family. For others it means volunteering, mentoring, or campaigning for social justice. Whatever it is, do it with gratitude, empathy and compassion. Even small actions make a big differ­ence—for you and for the people around you.

Me to We is a growing movement of people who are interested in mak­ing the most out of their lives and the lives of others. It is made up of individuals who have learned to live each day to the fullest, take nothing for granted, and leave a legacy of goodwill.
What do you want your legacy to be?

For more information on how you can become part of this movement, see our new book Me to We: Find­ing Meaning in a Material World.

International child rights activists Craig and Marc Kielburger are the co-authors of "Me to We: Finding Meaning in a Material World." They are also founders of Free The Children and Leaders Today. www. metowe.org, www.freethechildren. com. www.leaderstoday. com.


NUCLEAR CHEMISTRY


1. When two light nuclei combine into a heavier nucleus, the process is called (1) fusion, (2) fission, (3) nuclear disintegration, (4) beta decay.
2. The atomic number of a transuranium element is (1) 92, (2) 93, (3) 84, (4) 83.
3. The property of radioactivity that led to its discovery was its (1) electro­magnetism, (2) ionizing power, (3) penetrating power, (4) mass.
4. If the mass number of an element were to increase while the atomic number remained the same, the resulting product would be a(an) (1) ion, (2) nuclide with a different positive charge, (3) different isotope of the same element, (4) dif­ferent element.
5. When a transmutation occurs, there must be (1) a decrease in atomic number, (2) an increase in the number of nucleons in the nuclide, (3) a change in the mass number, (4) a change in the net nuclear charge.
6. An alpha particle is a(an) (1) helium nucleus, (2) electron, (3) neutron, (4) proton.
7. The type of radiation with the greatest penetrating power is (1) infrared, (2) gamma, (3) beta, (4) alpha.
8. Gamma radiation has (1) no charge and small particles, (2) a small mass and a charge of I –, (3) no mass and a charge of 1-tT , (4) no mass and no charge.
9. Radioactive changes differ from ordinary chemical changes in that radioactive changes (I) release energy, (2) involve changes in the nucleus, (3) result in an absorption of energy, (4) produce explosions.

10. Gamma radiation travels at (I) 20,000 miles/sec, (2) 10 000 miles/sec, (3) 186,000 miles/sec, (4) 100,000 miles/sec.
11. When compared with beta particles, alpha particles have (1) more mass and more ionizing power, (2) less speed and less ionizing power, (3) the most pene­trating power and the most ionizing power, (4) the least penetrating power and the least ionizing power.
12. Which are the most similar? (I) gamma rays and X-rays, (2) beta rays and X-rays, (3) alpha rays and beta rays, (4) alpha rays and gamma rays.
13. All of the following are useful in detecting or measuring radioactivity except the (I) electroscope, (2) spectroscope, (3) Geiger counter, (4) photographic plate.
14. A nuclide is transmuted by natural radioactive decay (1) in a 1:1 ratio, (2) always toward a lower atomic number, (3) in random fashion, (4) in the direc­tion of greater stability.
15. A transmutation is a reaction in which (I) atoms are transferred between two molecules, (2) the atomic number changes, (3) electrons are exchanged from one atom to another, (4) the nucleus is unchanged.
16. In nuclear equations, the letter Z is used to designate the (I) mass number, (2) symbol of an element, (3) number of neutrons, (4) atomic number.
17. As a radioactive element decays, its half-life (1) decreases, (2) increases, (3) remains the same.
18. The rate of decay of a radioactive element (I) increases with decreased pres­sure, (2) is not subject to change by physical means, (3) can be decreased by lowering the temperature, (4) can be increased by raising the temperature.
19. When the nucleus of an atom emits a negatively charged particle, the atomic number of the element (1) decreases, (2) increases, (3) remains the same.
20. The most stable of the following atomic nuclei is the nucleus of the atom of
( I) 2 Np, (2) 29pu, (3) 2(8)Tb, (4) 2,343u.
21. When a neutron decays, it is transformed into a(an) (I) proton, electron, gamma radiation, (2) proton, electron, neutrino, (3) electron, positron, neutrino, (4) alpha particle, beta particle, gamma radiation.
22. A radioactive isotope has a 5 year half-life. After a given amount decays for 15 years, what fraction remains of the original isotope? (1) 1/15, (2) 1/2, (3) 1/8, (4) 1/5.
23. Strontium-90 has a half-life of 28 years. At the end of 56 years an original sample of 8 g will weigh (1)1g, (2) 2 g, (3) 8 g, (4) 4 g.
24. The mass of the nutieus of an atom, when compared to the total mass of the individual particles comprising it, is (1) greater, (2) less, (3) the same.
25. The energy equivalent of the mass defect is called (1) the binding energy, (2) the fission fragments, (3) the moderating effect, (4) the zone of stability.
26. As a neutron is added to a nucleus, the atomic number of the element (I) decreases, (2) increases, (3) remains the same.
Base your answers to questions 27-31 on the diagram at the top of page 118 which represents a possible nuclear decay scheme.
27. The scheme indicates that a stable nucleus may be formed that has a mass of (1) 206, (2) 210, (3) 214, (4) 222.
28. In going from point A to point B, the element emits which particle? (1) an alpha, (2) a beta, (3) a neutron, (4) a deuteron.



PERIODIC TABLE





1. In early efforts to classify the elements Mendeleev and others arranged the elements in order of increasing
a. atomic numbers. c. atomic radii.
b. atomic masses. d. ionization energies.
2. The elements with similar properties fall into the same columns in the periodic table. The groups of elements with similar properties are called
a. periods. c. octaves.
b. triads. d. families.
3. Mendeleev's periodic table had several empty places. As it turned out Mendeleev was able to use these gaps to
a. reveal the electronic structure of atoms.
b. discover the properties of atomic numbers.
c. find errors in the identification of elements.
d. predict properties of undiscovered elements.
4. Moseley's work led to the current periodic table which arranges the elements on the basis of increasing
a. electron energy. c. atomic volume.
b. atomic number. d. density.





Use the periodic table to solve the following problems. Each letter represents a single element.

5. Three elements are represented by the letters A, B, and C. The three elements have the same number of valence electrons. Element A has electrons in three energy levels and element C is the lightest of the three. Element B has a total of 19 electrons. Identify the elements.
6. Four elements are represented by the letters J, K, L, and M. The four elements are all active metals and have the same number of valence electrons and a filled valence subshell. Element J is the most metallic and has electrons in five energy levels. Element M has a larger atomic radius than either K or L but is smaller than J. K has a higher first ionization energy than does L Identify the elements.




7. As the atomic number increases within a period, the radius of the particle
a. increases. c. stays the same.
b. decreases. d. changes in an irregular way.

8. Which of the following quantities can be measured directly?
a. electronegativity c. both a. and b.
b. ionization energy d. neither a. nor b.

B. Problems
Use the periodic table given to solve the following problems. Each letter represents a single element.
9. The particles U, V, W, X, and Y all have three shells of electrons and the same number of electrons. They are listed in order of increasing size. The only neutral atom is W. U reacts with Y to form the compound UY and with X to form the compound UX2. Identify the elements.
10. Antimony atoms may share electrons with bromine atoms, with sulfur atoms, or with nitrogen atoms. Use the electronegativity values for Sb (1.8), S (2.5), Br (2.8) and N (3.0) to rank the Sb-Br, Sb-S, and Sb-N bonds in order of decreasing attraction of the electrons for the antimony atom.

MULTIPLE CHOICE
11. Mendeleev left spaces in his periodic table and predicted several elements and their
a. atomic numbers. c. properties.
b. colors. d. radioactivity.

12. Mendeleev attempted to organize the chemical elements based on their
a. symbols. c. atomic numbers.
b. properties. d. electron configurations.


13. Mendeleev is credited with developing the first successful
a. periodic table. c. test for radioactivity.
b. method for determining atomic number. d. use of X rays.
14. In developing his periodic table, Mendeleev listed on cards each element's name, atomic mass, and
a. atomic number. c. isotopes.
b. electron configuration. d. properties
15. Mendeleev predicted that the spaces in his periodic table represented
a. isotopes. c. permanent gaps.
b. radioactive elements. d. undiscovered elements.
16. Mendeleev's table was called periodic because the properties of the elements
a. showed no pattern.
b. occurred at repeated intervals called periods.
c. occurred at regular time intervals called periods.
d. were identical.
17 The person whose work led to a periodic table based on increasing atomic number was
a. Moseley. c. Rutherford.
b. Mendeleev. d. Dimocritus
18. Who used his experimental evidence to determine the order of the elements according to atomic number?
a. Meyer c. Stas
b. Ramsay d. Moseley
19. The most useful source of general information about the elements for anyone associated with chemistry is a
a. calculator. c. periodic table.
b. table of metric equivalents. d. table of isotopes.
20. Evidence gathered since Mendeleev's time indicates that a better arrangement than atomic mass for elements in the periodic table is an arrangement by
a. mass number. c. group number.
b. atomic number. d. series number.
21. What are the elements whose discovery added an entirely new row to Mendeleev's periodic table?
a. noble gases c. transition elements
b. radioactive elements d. metalloids
22. The discovery of the noble gases changed Mendeleev's periodic table by adding a new
a. period. c. group.
b. series. d. sublevel block.

23. In the modern periodic table, elements are ordered according to
a. decreasing atomic mass. c. increasing atomic number.
b. Mendeleev's original design. d. the date of their discovery.
24.The periodic law states that the physical and chemical properties of elements are periodic functions of their atomic
a. masses. c. radii.
b. numbers. d. structures.
25. The periodic law states that the properties of elements are periodic functions of their atomic numbers. This means that the_____________________________ determines the position of each element in the periodic table.
a. mass number c. number of protons
b. number of neutrons d. number of nucleons
26. The periodic law allows some properties of an element to be predicted based on its
a. position in the periodic table. c. symbol.
b. number of isotopes. d. color.
27. The periodic law states that
a. no two electrons with the same spin can be found in the same place in an atom.
b. the physical and chemical properties of the elements are functions of their atomic numbers.
c. electrons exhibit properties of both particles and waves.
d. the chemical properties of elements can be grouped according to periodicity but physical properties cannot.

28. A horizontal row of blocks in the periodic table is called a(n)
a. group. c. family.
b. period. d. octet.

29. To which group do lithium and potassium belong?
a. alkali metals c. halogens
b. transition metals d. noble gases
30. To which group do fluorine and chlorine belong?
a. alkaline-earth metals c. halogens
b. transition elements d. actinides





31. Atomic size is determined by measuring the
a. radius of an individual atom.
b. distance between nuclei of adjacent atoms.
c. diameter of an individual atom.
d. volume of the electron cloud of adjacent atoms.
32. A measure of the ability of an atom in a chemical compound to attract electrons is called
a. electron affinity. c. electronegativity.
b. electron configuration. d. ionization potential.
33. The element that has the greatest electronegativity is
a. oxygen. c. chlorine.
b. sodium. d. fluorine.
34. One-half the distance between the nuclei of identical atoms that are bonded together is called the
a. atomic radius. c. atomic volume.
b. atomic diameter. d. electron cloud.
35. Ionization energy is the energy required to remove____________ from an atom of an element.
a. the electron cloud c. an electron
b. the nucleus d. an ion

36. In a row in the periodic table, as the atomic number increases, the atomic radius generally
a. decreases. c. increases.
b. remains constant. d. becomes unmeasurable.

37. As the atomic number of the metals of Group 1 increases, the ionic radius
a. increases. c. remains the same.
b. decreases. d. cannot be determined.

38. The ionization energies required to remove successive electrons from one mole of calcium atoms are 590 kJ/mol, 1145 kJ/mol, 4912 kJ/mol, and 6474 kJ/mol. The most common ion of calcium is probably
a. Ca+ , . c. Ca3+.
b.Ca2+. d. Ca4+.
39. you move down the periodic table from carbon through lead, atomic radii
a. generally increase. c. do not change.
b. generally decrease. d. vary unpredictably.

40. When chemical compounds form, valence electrons are those that may be
a. lost only. c. shared only.
b. gained only. d. lost, gained, or shared.

41. The number of valence electrons in Group 1 elements is
a. 1. c. 8.
b. 2. d. equal to the period number.
42. The number of valence electrons in Group 17 elements is
a. 7. c. 17.
b. 8. d. equal to the period number.
43. The number of valence electrons in Group 2 elements is
a. 2. c. 18.
b. 8. d. equal to the period number.
Essay:
44. What can you predict about the properties of xenon and helium, both in Group 18 in the periodic table? Why

REASONING EXERCISES
45. Distinguish between the basis of Mendeleev's Periodic Table and the basis of the modern Periodic Table.
46. What is the Periodic Law? How was it determined?
47. Distinguish between the periodic behavior of the valence electrons along Period 2 and in Group 0 (zero). How does the number of valence electrons illustrate the Peri­odic Law?
48. The elements show an increase in various properties with an increase in atomic number throughout the Periodic Table. How is this increase manifested along Period 2? How is it manifested along Group IA?
49. What is meant by the representative elements? Distinguish between the transition metals and the metals of Groups IA and IIA.
50. Distinguish between the metalloids and the nonmetals.
51. Why are the alkali metals extremely reactive? Distinguish between the chemical behavior of the alkali metals and the halogens.
52. Letting M stand for metal, write a general formula for the metals of Group IIIA and bromine.
53. Where do you think the relative increase in atomic volume (size) is the greater—along a period, or along a group? Explain your reasoning.
54. How many principal energy levels are there in calcium? How many of these principal energy levels are in the kernel? How many valence electrons are there? Describe the change in electronic structure along Group IIA.

REASONING EXERCISES
55. What two groups have the strongest tendency to form ionic bonds with each other?
56. Describe the relation between metallic characteristics and the oxidation number as we move from left to right along a period.
57. Distinguish between the atomic radius and the ionic radius. Why is the ionic radius of sodium smaller than the atomic radius?
58. Describe the periodic change of atomic radii (a) along a period, (b) down a group. What reasons can you give for each trend?
59. How do ionic radii change systematically along a group? In the case of bromine, which is larger— the atomic radius or the ionic radius?
60. Show by reasoning that an alkali metal atom exerts a lesser relative force of attraction for a bonding electron than a halogen atom does.
61. Distinguish between each of the following from the standpoint of periodic progres­sion along an A group and a period: (a) atomic radius, (b) ionization energy, (c) elec­tronegativity, (d) metallic properties.
62. Review Section 8.4 of Chapter 8 on ionization energy. Then answer the following: (a) Why do ionization energies decrease down a group in the Periodic Table? (b) Why do ionization energies increase across a period?
MULTIPLE-CHOICE QUESTIONS
63. The modern Periodic Table is based primarily on the work of Mendeleev and
(1) Moseley, (2) Thomson, (3) Newlands, (4) Dobereiner.
64. The scientist who proposed that the periodic properties of the elements were a function of their increasing atomic weights was (1) Mendeleev, (2) Dobereiner, (3) Newlands, (4) Moseley.
65. If the elements of the Periodic Table were arranged by increasing atomic weight, all of the following would be in correct order except (1) Ar and K, (2) I and Xe, (3) Cu and Zn, (4) Se and Br.
66. All of the following are properties of most metals except (1) conduction of heat,
(2) reaction with water to free hydrogen, (3) malleability, (4) ductility.
69. The main reason why atomic volume decreases with increasing atomic number along a period is because of increasing (1) number of neutrons, (2) number of shells, (3) nuclear charge, (4) interelectronic repulsion.
72. Down any one group with increasing atomic number, ionic radii (1) decrease, (2) increase, (3) remain the same.

REASONING EXERCISES
73. Why does the reactivity of the alkali metals increase with increasing atomic number? Why does the reactivity of the halogens decrease with increasing atomic number?
74. Why do you think alkaline earth metals occur in nature only as compounds? How are they changed to their free state? Why is sodium more reactive than magnesium?
75. Nitrogen and phosphorus are both nonmetals of Group VA. Since the reactivity of nonmetals tends to decrease with increasing atomic number, how can we explain the fact that nitrogen and phosphorus provide an exception to this trend?
76. Why does oxygen exist in the free state in spite of its high reactivity?


CHEMICAL BONDING


1. A chemical bond results from the mutual attraction of the nuclei of atoms for (a) electrons;
(b) protons; (c) neutrons; (d) dipoles.
2. An ionic bond results from electrostatic attraction between (a) ions; (b) dipoles; (c) electrons;
(d) orbitals.
3. A bond with an ionic character of less than 5% is considered to be (a) polar covalent; (b) ionic;
(c) nonpolar covalent; (d) metallic.
4. A nonpolar covalent bond is unlikely when two different atoms join because the atoms are
likely to differ in (a) electronegativity; (b) density; (c) state of matter; (d) polarity.
5. Bond length is equivalent to the distance of (a) maximum kinetic energy; (b) minimum potential
energy; (c) maximum potential energy; (d) one-half the diameter of the electron cloud.
6. In many compounds, atoms of main-group elements form bonds so that the number of
electrons in the outermost energy levels of each atom is (a) 2; (b) 6; (c) 8; (d) 10.
7. In order to draw a Lewis structure, it is NOT necessary to know (a) bond energies;
(b) the types of atoms in the molecule; (c) the number of valence electrons for each atom;
(d) the number of atoms in the molecule.
8. Multiple covalent bonds may occur in atoms that contain carbon, nitrogen, or (a) chlorine;
(b) hydrogen; (c) oxygen; (d) helium.
9. Methane, CH4, contains how many covalent bonds? (a) 3 (b) 4 (c) 5 (d) 8
10. A formula unit of an ionic compound (a) is an independent unit that can be isolated and studied; (b) is the simplest ratio of ions that gives electrical neutrality; (c) describes the crystal
lattice; (d) all of the above.
11. The ions in an ionic compound are organized into a(n) (a) molecule; (b) Lewis structure;
(c) polyatomic ion; (d) crystal.
12. Lattice energy is an indication of the (a) strength of an ionic bond; (b) strength of a metallic
bond; (c) strength of a covalent bond; (d) number of ions in a crystal.
13. A compound that vaporizes readily at room temperature is most likely to be a(n)
(a) molecular compound; (b) ionic compound; (c) metal; (d) brittle compound.
14. The electron-sea model of a metallic bond consists of (a) stationary electrons; (b) electrons
bonded to particular positive ions; (c) valence and nonvalence electrons; (d) mobile electrons
shared by all the atoms.
15. Metals are malleable because the metallic bonding (a) holds the layers of ions in rigid
positions; (b) does not produce ions; (c) allows one plane of ions to slide past another;
(d) allows easy breaking of the bonds in a metal.
16. To appear lustrous, a material must be able to (a) form crystals; (b) absorb and re-emit light of
many wavelengths; (c) absorb light and change it all to heat; (d) change light to electricity.
17. VSEPR states that the electrostatic repulsion between electron pairs surrounding an atom
causes (a) an electron sea to form; (b) positive ions to form; (c) these pairs to be oriented as far
apart as possible; (d) light to be reflected.
18. According to VSEPR theory, the shape of an AB4 molecule would be (a) linear; (b) octahedral;
(c) bent; (d) tetrahedral.
19. Due to the presence of an unshared pair of electrons, the shape of an AB3E molecule, such as
ammonia, is (a) tetrahedral; (b) linear; (c) triangular pyramidal; (d) octahedral.
20. Compared to molecular bonds, the strength of intermolecular forces is generally (a) weaker;
(b) stronger; (c) about the same; (d) too variable to compare.

21. In a chemical bond, the link between atoms results from the attraction between electrons and
(a) Lewis structures; (b) nuclei; (c) van der Waals forces; (d) isotopes.
222. The attraction of an atom for the shared electrons that form a covalent bond between it and another atom is called its (a) electron affinity; (b) electronegativity; (c) resonance;
(d) hybridization.
23. Bonds with more than 50% ionic character are considered to be (a) polyatomic;
(b) polar covalent; (c) ionic; (d) nonpolar covalent.
24. Nonpolar covalent bonds are not common because (a) one atom usually attracts electrons more strongly than the other; (b) ions always form when atoms join; (c) the electrons usually remain equally distant from both atoms; (d) dipoles are rare in nature.
25. The energy released in the formation of a covalent bond is the difference between zero and the (a) maximum potential energy; (b) kinetic energy of the atom; (c) minimum potential energy; (d) bond length expressed in nanometers.
26. The principle that states that atoms tend to form compounds in such a way as to allow each atom to have 8 electrons in its outermost energy level is called the (a) rule of eights;
(b) Avogadro principle; (c) configuration rule; (d) octet rule.
27. In writing a Lewis structure, the central atom is the (a) atom with the greatest mass; (b) atom with the highest atomic number; (c) atom with the fewest electrons; (d) least electronegative atom.
28. In writing a Lewis structure for a polyatomic ion, one electron must be added for each unit of
(a) kcal/mol; (b) mass; (c) polarity; (d) negative charge.
29. How many potassium ions are there in a formula unit for the compound potassium chloride, KCl? (a) 0 (b) 1 (c) 2 (d) 3
30. In a crystal, the valence electrons of adjacent ions (a) repel each other; (b) attract each other;
(c) neutralize each other; (d) have no effects on each other.
31. Lattice energy is the energy released in the formation of a(n) (a) Lewis structure; (b) polar covalent compound; (c) nonpolar covalent compound; (d) ionic compound.
32. Ionic compounds are brittle because the strong attractive forces (a) allow the layers to shift easily; (b) cause the compound to vaporize easily; (c) keep the surface dull; (d) hold the layers in relatively fixed positions.
33. The electron-sea model of bonding is characteristic of (a) covalent bonding; (b) metallic bonding; (c) ionic bonding; (d) hydrogen bonding.
34. When a metal is drawn into a wire, the metallic bonds (a) break easily; (b) break with difficulty;
(c) do not break; (d) become ionic bonds.
35. To appear lustrous, a metal must be able to (a) form crystals; (b) absorb and re-emit light of many wavelengths; (c) absorb light and change it all to heat; (d) change light to electricity.
36. The idea that electrostatic repulsion between electron pairs surrounding an atom causes these pairs to be oriented as far apart as possible is called (a) VSEPR theory; (b) the hybridization model; (c) the electron-sea model; (d) Lewis theory.
37. In VSEPR theory, a molecule that is triangular planar in shape is classified as (a) AB2; (b) AB3;
(c) AB4; (d) AB6.
38. The water molecule, with two unshared electron pairs, has the VSEPR designation (a) AB2E2;
(b) AB2; (c) AB6; (d) AB2E.
39. In orbital notation, the hybridized orbitals responsible for the shape of CH4 are identified as (a) 1s21p3; (b) sp2; (c) 2s22p6; (d) spa.


A. Multiple Choice
Select the word, number, or phrase that best completes each statement and write its letter in the answer space at the left.
40. The bonds between the atoms in a polyatomic ion such as CO32 are
a. metallic. b. ionic. c. covalent. d. any of the above.
41. The hybridization of the atoms of an element in a chemical reaction results in the
a. demotion of electrons to lower energy levels.
b. formation of a number of orbitals with identical energy.
c. elimination of electrons from the valence shell.
d. formation of ionic bonds from covalent bonds.
42. Substances made of molecules that are capable of hydrogen bonding have unexpected properties including all of the following except
a. high heats of vaporization. c. high boiling temperatures.
b. low melting points. d. low vapor pressures.
43. Which of the following properties would not be associated with ionic solids?
a. no molecular units c. conductor of electricity
b. fixed positions of ions d. high melting point
44. The force of the attraction that holds two atoms together is
a. a magnetic attraction. c. a gravitational attraction.
b. an electrical attraction. d. a nuclear attraction.
45. In a dot diagram, the chemical symbol at the center of the diagram represents the
a. nucleus of the atom.
b. valence electrons.
c. kernel of the atom.
d. inner electrons of the atom.
46. The atoms with the lowest ionization energies are found in the group of elements called
a. noble gases. b. nonmetals. c. metals. d. semimetals.
47. A molecule in which four shared electron pairs are spaced symmetrically around a central atom has the shape of a
a. tetrahedron. b. planar. c. pyramid. d. trigonal bipyramid.
48. Which of the following statements about polar bonds and polar molecules is not correct?
a. polar molecules contain polar bonds
b. polar molecules can contain nonpolar bonds
c. nonpolar molecules can contain only nonpolar bonds
d. nonpolar molecules can contain polar bonds
49. The weak forces of attraction that result from the shifts of electron positions in mol‑ecules have the special name of
a. network forces. c. hydrogen forces.
b. metallic forces. d. van der Waals forces.
________

50. Hydrogen bonds between molecules are formed in compounds where hydrogen atoms are bonded to atoms of elements with
a. low electronegativity and large size.
b. high electronegativity and small size.
c. low electronegativity and small size.
d. high electronegativity and large size

51. A coordinate covalent bond differs from an ordinary covalent bond in only the
a. amount of ionic character of the bond.
b. number of electrons involved in the bond.
c. strength of the bond formed by the electrons.
d. source of the electrons forming the bond.
52. When metals and nonmetals combine to form compounds, each of the nonmetals formsthe electron configuration of
a. the noble gas nearest to it in the periodic table.
b. an element with two electrons in its valence shell.
c. an active metal or an active nonmetal.
d. an inactive metal or an inactive nonmetal.
53. When a bond is formed between two atoms the level of energy of the bonded atoms is
a. lower than before bonding. c. greater than before bonding.
b. the same as before bonding. d. unrelated to whether or not a bond forms.
54. The particles formed by atoms that lose one or more electrons are
a. anions. b. cations. c. subions. d. superions.
55. Two atoms of nitrogen form a triple bond. What is the ratio of electrons involved in thetriple bond to the total number of electrons in the valence shells of the two atoms?
a. 3 : 10 b. 6 : 18 c. 3 8 d. 6 : 10
56. The ratio of numbers of shared
pairs of electrons to unshared pairs of electrons to produce the molecular shape shown would be
a. 3 : 1. c. 4 : 0.
b. 1 : 3. d. 2 : 2.
57. Substances most likely to be gases at room temperature would be made up of
a. metallic crystals. c. nonpolar molecules.
b. ionic crystals. d. polar molecules.
58. The characteristic of metallic bonding that distinguishes it from other bonds is the
a. fixed position of the valence electrons.
b. polar property of substances with metallic bonds.
c. freedom of movement of the valence electrons.
d. directional nature of the bonds formed.






SIGNIFICANT FIGURES



Truth in Imperfect Measurement

Only those who have the patience to do simple things perfectly will acquire the skill to do difficult things easily. -- Johann von Schiller

Being honest and telling the truth while avoiding hype is critical for the information age. But this is nothing new for scientists. Isaac Newton, Robert Boyle, and Antoine Lavoisier in their quest to abandon alchemy and superstition established a custom of precisely communicating clear information.
The avoidance of cute introductory jokes and antidotes often makes science see dry and impersonal. But scientists were centuries ahead of most other professions in recognizing the needs of information exchange.
It was no accident that scientists were among the first to adopt computers, that a scientist at CERN (European Nuclear Research) invented WWW, or that the first U.S. Internet hub was at the Stanford Linear Accelerator Center. The skill of communicating measurements honestly is at the core of effective information exchange.
Uncertainty in Measurement
Measurements are not as perfect as other aspects of mathematics. Each of our senses has limited acuity for detecting slight differences. And each measuring tool also has a tolerance for possible errors. Significant Figures are ONE way to tell the truth about the precision of a measurement.
We should distinguish "accurate" from "precise." Precision and accuracy are terms used to describe the quality of a measurement. Some view these terms as synonyms, but in fact they are different.
Precision indicates the degree of reproducibility of a measurement. It depends on how well you make a measurement.
Accuracy describes how close a measured value is to the true value. It depends upon the quality and calibration of the measuring device.
Example
Imagine working in a carnival, guessing people's mass. Lets say the same person comes back 4 different times and each time you guess a different mass. If the person's true mass is 60 kg ( about 132 lbs) and

your guesses are:
· 56 kg
· 65 kg
· 70 kg
· 51 kg
· Average = 60.5 kg These represent good accuracy but poor precision.
If instead your guesses had been:
· 68 kg
· 69 kg
· 67 kg
· 68 kg
· Average = 68.0 kg They would have represented poor accuracy but good precision.
It is often difficult to determine accuracy. Scientists rely on repeated measurements using different tools and different methods to determine accuracy.
Significant Figures:
Why worry about significant figures? Since the quality and calibration of our equipment and the care of our measurement can influence the accuracy and precision of our measurement, we need a way to describe what confidence should be placed in our reported measurement. There are several ways to do this: significant figures is the easiest.
Two Examples to Ponder:
When you add a quart of oil to your car, do you need to measure out 1.000 quarts in volumetric flask, or is it sufficient to just add about 1 quart?
If promoter sells tickets to a concert, do they need to know if there are 19,852 seats or is it good enough to know that there are 20,000 give or take 5000?
The simplest way to describe the uncertainty of a measurement is to combine information about the quality of the measurement with the number of the measurement. We let the number of digits reported for a measurement tells us about the certainty of a measurement. The rule is
that we know all of the digits with certainty up to the last digit, which is estimated.
Example
Lets say you wanted to measure your weight using a bathroom scale with a digital readout. The readout might say you weigh 164 lbs. Providing the scale is well calibrated the measurement tells you that your exact weight is between 163.5 and 164.5 pounds. It would be OK to say you weigh about 160 lbs but it would be misleading to claim that you weigh 163.892 lbs. It is acceptable to discard information if it won't be needed, but it's a lie to claim more information than is known.

Rules for significant figures follow. They are an important part of experimental science, but they don't make for a very exciting lecture. They allow you to properly interpret measured results, and in certain cases they can save you a lot of time.
You will be expected to use the correct number of significant figures in lab reports as well as on quizzes and exams.
Such skill is merely useful in science but will be invaluable for the measurements you rely on during the rest of your life!
Several Points of Clarification:

Significant figures are the easiest way to deal with uncertainty in measurements, but it is not a perfect system. The Mesopotamians invented zero over 2000 years ago. We use it to distinguish when we know there is none. But we also use the zero as a place holder (zero wasn't invented 2 years ago).

Zeros used only as place holders are not significant figures while their other use is. Several rules to work around this ambiguity will follow.
Some definitions involve exact numbers. (100 cm = 1 m, 12 = 1 dozen) Numbers can be exact by count. (5 playing on a basketball team)
Exact numbers have an infinite number of significant figures, so when they are used in a calculation they do not change the number of significant figures.

Significant Figures Rules:

· The left most digit which is not a zero is the most significant digit.
· If the number does not have a decimal point, the right most digit which is not a zero is the least significant digit.
· If the number does have a decimal point, the right most significant digit (—>) is the least significant digit, even if it's a zero.
· Every digit between the least and most significant digits should be counted as a significant digit.
For example,
according to these rules, all of these numbers have three significant digits: 123
123,000
123

1.23 x 106 1.00
0.000123
How many significant figures should one retain_ in the final answer to a problem?
· For results obtained using addition or subtraction, the number of places after the decimal point in the result should be less than or equal to the number of decimal places in every term.
· For results obtained using multiplication or division, the number of significant figures in the result should be equal to the number of significant digits in the least precise number (the number with the fewest significant digits given).
It is sometimes good practice to give one more significant figure than is required by these rules; this helps prevent rounding errors if the number is used in later calculations. This extra digit should be in a smaller font to indicate its lesser significance.
Details and Examples
Introduction
Significant figures are a shorthand way to express how certain one is about one's data and calculations coming from that data. While significant figures are by no means as precise as detailed calculations of the uncertainty of a value, they are a very useful way to estimate uncertainty quickly.
Uncertainty and its Meaning
Any value that is the result of a scientific measurement has some uncertainty. The most precise way to state the uncertainty of a measurement is to write it as a number, plus or minus the expected error in that number. Scientists often use the standard deviation for their expected error. After a measurement has been repeated many times, the standard deviation is determined by averaging the absolute amounts by which each measurement differs from the mean.
Fine Print
For example, if you measured a wire's length 30 times, and got an average length of 28.3 cm, with an average error in that length of about 0.2 cm, you would write the length as (28.3 ± 0.2) cm if you wanted to be precise about your measurement results. This means that the majority of your measurements fall between 28.1 and 28.5 cm. (To be more specific, it means that 68% of the measurements fall between those two values.)
If, for example, we then wished to find the volume of this wire, and we had a measurement of 2.31 ± 0.07 mm for the wire's diameter, we would have to put these numbers into the formula for the volume of a cylinder. This would require doing separate calculations for the largest and smallest errors in addition to the calculation for the average values. To be precise in our treatment of uncertainty, we need some complicated mathematics to see how the errors on the individual quantities translate through the formula to become errors on the volume.

But for most measurements you will ever use, a simpler system of dealing with uncertainty will be adequate. This system is the system of significant figures.
Significant Figures: The Simplest Method for Expressing Uncertainties
If we just want an approximate idea of the extent to which a value is certain, and we either don't want to learn the mathematical techniques, or don't want to spend the time to apply them, we can keep track of the amount of certainty of a piece of data simply by paying attention to the number of digits we use to express it. That is to say, where we choose to round off our number tells where we think uncertainty creeps in. For example, if we have a length of 12.37 ± 0.10 cm, we just call the length 12.4 centimeters, to three significant figures. When we express a number with three significant figures, what we are saying is that the first two digits are essentially exactly correct, and the last one is uncertain by a small amount (generally it is only uncertain by about ± 1). In the example above, we rounded our answer to 12.4 cm because our answer is uncertain to ± 0.1 cm, viz., our answer is uncertain in the last digit by about 1.
How Many Digits to Use?
The question of the greatest practical importance is how many digits to include in your final answer. This is important because, as explained above, the number of digits you include in your answer shows the reader the precision of the data leading to the answer, and the accuracy of the answer. It might be useful to read this section again after reading through the following sections which explain how to determine the number of significant digits.
Addition and Subtraction
When adding and subtracting numbers, the rules of significant figures require that the number of
places after the decimal point in the answer is less than or equal to the number of decimal places in every term in the sum. (Treat subtraction as adding the same number with a negative sign in front of it.) If some of the numbers have no digits after the decimal point, use the same basic rule, but don't record any digits to the right of the last digit in the least significant number. Clarify these rules, are
some examples:
2355.2342 15600.00 15600 13.7 137000
+ 23.24 + 172.49 + 172.49 + 1.3 + 1330
2378.47 15772.49 15800 15.0 138000
Note it is not unusual for a sum to have more significant figures than the measurements added. This is why finding an average gives greater information than a single measurement.
Also note that a difference often has fewer significant figures. This apparent shortcoming is sometimes used by scientists in reverse! One of the most sensitive tests is to measure a null difference to verify that the much larger opposing forces are equal to an accuracy not directly measurable. Inverse squared force laws have been verified to a large number of significant figures this way.
Multiplication and Division

(Calculators are not programmed to do significant figures because they have no way to recognize which numbers are constants.)When multiplying and dividing numbers, the number of significant digits you use is simply the same number of significant figures as is the number with the fewest significant figures.
Some examples:
13.1 13.10 13.100 1500 15310 1.00
x 2.25 x 2.25 x 2.2500 x 2.315 x 2.3 x 10.04
29.5 29.5 29.475 3400 35000 10.0
Someday calculators may be able to do significant figures; but in the meantime the operators need that wisdom.
Why do Multiplication and Addition Have Different Rules?
When you add two numbers, you add their uncertainties, more or less. If one of the numbers is smaller than the uncertainty of the other, it doesn't make much of a difference to the value (and hence, uncertainty) of-the final result. Thus it is the location of the digits, not the amount of digits that is important.
When you multiply two numbers, you more or less multiply the uncertainties. Thus it is the percentage by which you are uncertain that is important -- the uncertainty in the number divided by the number itself. This is given roughly by the number of digits, regardless of their placement in terms of powers of ten. Hence the number of digits is what is important.
Which Digits are "Significant?"
In order to figure out how many significant figures to put into your final answer you must figure out how many significant figures are in each of the numbers you are working with. The rules are best explained separately for fundamental constants, physical constants, numbers not ending in 0 and numbers ending in 0.
Fundamental Constants
Fundamental constants are numbers without units of any kind that come strictly from mathematics; they are not "measured" like most quantities in science. Some examples of these are regular integers such as 2, 10, 14, or 27; fractions such as the 4/3 in the formula:
V = (4/3)piR3,
for the volume of a sphere; and constants such as and e, where e = 2.71828... is the base of the natural logarithms.
When you have one of these numbers, you should never let it determine how many significant figures you have. If the number is an integer or a rational fraction, just assume it has more significant figures that the least accurate of the measurements. When you have an irrational number like pi, look up as many digits as you need so it has more digits than required by the least accurate measurement.

Physical Constants
When you are dealing with a physical constant such as Planck's constant, the speed of light, or the charge of an electron, you should remember that these numbers are found by experiment and do not have any purely mathematical definition. So there may be some occasions where you have a piece of data that is more certain than the best value of your physical constant. In these cases, it is acceptable to let your physical constant define your uncertainty, and hence your number of significant figures. To avoid this problem it has been a top priority of some scientists to measure these physical constants very accurately.
Numbers Without Zeroes at the End
Numbers without zeroes at the end are the simplest case. When a given piece of data ends in a digit other than zero, all the digits in that measurement are significant digits.
Numbers with Zeroes at the End
Numbers ending in zero are more complicated because zero has two different meanings. In these cases, you must determine whether a zeroes is a significant digit representing the quantity of "none" or just a place holders.Zeroes right (---) of a decimal point are always significant digits because without them the actual value of the number is no different, so we assume they are placed there to show additional certainty in the value of the number. (Zeroes between other significant digits are also always significant.) Any other zero between the last non-zero digit and the decimal point is not a significant digit.
Examples:
· 130 has two significant digits.
· 130.0 has four significant digits.
· 1.000000 has seven significant digits.
· 100000 has one significant digit.
· 1.30 x 102 is the way you write 130 if you want to make it clear that there are 3, not 2 significant digits in the number.
Perhaps someday there will be universal adoption of another symbol, perhaps 0, to represent "none" and end the confusion. For example, 1300 would have 3 significant digits.
Exception to the Rules?
For those doing calculations on a calculator (or computer), it is wise to carry through all the digits in the calculator until the very end of the problem, and then truncate your final answer to the correct number of significant digits. In other words, during some intermediate step of the calculation, don't attempt to eliminate the insignificant digits in your calculator or write down an intermediate answer with fewer digits and then re-enter that new number into the calculator. Following the rules above, you might be tempted to round off midway through a problem, but doing so could introduce a small (but sometimes non-negligible) amount of "truncation error." This is especially prone to happen if you have a complicated calculation involving many steps where truncation errors could accumulate. In addition, it is possible to make an error re-entering a number. These types of error are totally avoidable (as opposed to measurement errors, which we are stuck with) and therefore it is best to keep all intermediate digits until the final answer.
· created 7/31/2000minor revisions 2/25/2006 by D Trapp






SIGNIFICANT FIGURES

Significant figures are the reliable digits in a number or measurement
which are known with certainty.
Significant figures show the accuracy in measurements. We can understand the precision of
a measurement if we know exactly the significant figures in the measurement.
A measurement that contains more number of significant figures is more accurate than a measurement that contains less number of Significant figures.
For example: Radius of a bob is 3.3679 cm and that of the other is 3.36 cm. In this situation the first measurement is the most accurate as it has more number of significant figures.
Rules Of Significant Figures www.citycollegiate.com
In order to determine significant figures in a number we must follow the following rules:
(1) All the non-zero digits are significant figures.
For Example:
3.456 has four significant figures.
12.3456 has six significant figures.
0.34 has two significant figures.
(2) Zeros between non-zero digits are significant.
For Example:
2306 has four significant figures.
20,0894 has six significant figures.
(3) Zeros locating the position of decimal in numbers of magnitude less than one are not significant.
For Example:
0.2224 has only one significant figures.
0.0000034 has two significant figures.
(4) Final zeros to the right of the decimal point are significant.
For Example:
3.0000 has five significant figures.
1002.00 has six significant figures.
(5) Zeros that locate decimal point in numbers greater than one are not significant.
For Example:
30000 has only one significant figure.
120000 has two significant figures.
Rule # 1:
If the digit to be dropped is greater than 5, then add "1" to the last digit to be retained and drop all digits farther to the right.
For example:
3.677 is rounded off to 3.68 if we need three significant figures in measurement. 3.677 is rounded off to 3.7 if we need two significant figures in measurement. Rule # 2:
If the digit to be dropped is less than 5, then simply drop it without adding any number to the last digit.
For example:
6.632 is rounded off to 6.63 if we need three significant figures in measurement. 6.632 is rounded off to 6.6 if we need two significant figures in measurement. Rule # 3:
If the digit to be dropped is exactly 5 then:
(A) If the digit to be retained is even, then just drop the "5".
For example:
6.65 is rounded off to 6.6 if we need two significant figures in measurement. 3.4665 is rounded off to 6.466 if we need four significant figures in measurement.
(B) If the digit to be retained is odd, then add "1" to it.
For example:
6.35 is rounded off to 6.4 if we need two significant figures in measurement. 3.4675 is rounded off to 6.468 if we need four significant figures in measurement. Remember: Zero is an even number
3.05 is rounded off to 3.0 if we need two significant figures in measurement.
Use of significant figures in
www.citycollegiate.com addition and subtraction
In addition and subtraction we consider the significant figures on the right side of decimal point. This
means that only as many digits are to be retained to the right side of decimal point as the number with fewest digits to the right of the decimal point.
For example:
4.345 + 23.5 =27.845 (actual answer by using calculator)
Answer after rounding off: 27.8
Use of significant figures in
multiplication and division
In multiplication and division , the number obtained after calculation of two or more
numbers must have
no more significant figure than that number used in multiplication or division.
For example:
4.3458 x 2.7 =11.73366(actual answer by using calculator)
Answer after rounding off: 12(because 2.7 has only two significant figures)



Solutions



1. The process of dissolving a solid in a liquid involves separating particles in the solid from one another. As a result this process is normally
a. dissociation. c. endothermic.
b. precipitation. d. exothermic.
2. Solvents consisting of polar molecules are more effective in dissolving solutes that are made up of
a. polar molecules.
b. nonpolar molecules.
c. nonionic particles.
d. covalent molecules.
3. An increase in temperature increases both the rate of dissolving and the amount that dissolves for most
a. solid solutes.
b. liquid solutes.
c. gaseous solutes.
d. all of the above.
4. The rate of solution of a particular substance is affected by
a. stirring.
b. the temperature.
c. the size of particles.
d. all of the above.
5. The amount of a substance that dissolves in another substance is affected by all of the following except
a. the nature of the substances. c. the pressure.
b. the temperature. d. stirring.
6. The specific gravity of a substance is really the
a. actual weight of the substance in air.
b. apparent weight of the substance in water.
c. ratio of the volume of the substance to the volume of an equal mass of water.
d. ratio of the mass of the substance to the mass of an equal volume of water.
7. When substances are dissolved in water the effect is to
a. raise the boiling point and lower the freezing point of the water.
b. raise both the boiling point and freezing point of the water.
c. lower both the boiling point and freezing point of the water.
d. lower the boiling point and raise the freezing point of the water.
8. The salt that is dissolved in water to make a solution is called the
a. sediment. c. solute.
b. tincture. d. solvent.
9. The properties of a solution include all of the following except that
a. it is a homogeneous mixture if it has been well stirred.
b. dissolved particles will settle out upon standing.
c. it is clear and transparent with particles too small to be seen.
d. dissolved particles will pass through a piece of filter paper.
10. A 1.0 molar solution of sucrose in water is prepared and divided into two equal volumes, A and B. More sucrose is added to volume A and more water is added to volume B. Which of the following statements is correct?
a. A is more concentrated than B c. A is less saturated than B
b. A is more dilute than B d. A and B have the same molarity
11. A drinking glass full of cold water is left standing on the table. After a few minutes bubbles appear in the water. This observation supports the generalization that
a. water is a polar solvent.
b. gases are attracted to glass.
c. water can be made to boil at any temperature.
d. gases are more soluble in cold water than in warm water.



12. When a gas dissolves in liquid water one would expect, since there is a phase change of the solute, that there would be an accompanying
a. temperature increase. c. pressure increase.
b. temperature decrease. d. pressure decrease.
13. The molality of a solution refers to the number of moles of the solute in a
a. mole of solvent. c. cubic decimeter of solvent.
b. kilogram of solvent. d. cubic decimeter of solution.
14. The number of moles of a solute in solution can be determined from the
a. product of the molar mass and the molarity.
b. quotient of the molarity divided by the volume.
c. quotient of the volume divided by the molarity.
d. product of the volume and the molarity.
15. The molarity of a solution refers to the number of moles of the solute in
a. a mole of solvent. c. one cubic decimeter of the solvent.
b. a mole of solution. d. one cubic decimeter of the solution.
16. If a substance is only slightly soluble in water, then which of the following terms could not be used to describe a solution of the substance in water?
a. concentrated b. saturated c. dilute d. unsaturated
17. A solution is saturated when the solute in solution is
a. no longer visible to the naked eye.
b. starting to come out of solution as a solid.
c. in equilibrium with the undissolved. solute.
d. no longer coming out of solution as a solid.
18. The most common types of solutions are
a. gaseous solutions. b. liquid. c. solid solutions. d. amalgams.
B. Problems
Solve the following problems in the spaces provided. Show all your work.
19. Express the solubility of potassium nitrate (KNO3) per 100 g of water at 50°C if 21.0 g of KNO3 dissolves in 25 g of water at 50°C.
20. What is the molarity of a solution that contains 5.85 g of sodium chloride (NaCl) in 0.50 dm3 of a water solution?
21. What is the molality of a solution that is prepared by dissolving 5.0 g of calcium sulfate (CaSO4) in 0.50 kg of water?
C. Essay Question
22. Describe the differences in procedures used to prepare a one molar solution compared to a one molal solution.
23. The freezing point depression depends on the concentration of the solute particles rather than on the kind of particles. This kind of property is called
a. numerative. b. transitive. c. colligative. d. intensive.
24. A 35.0-g sample of a nonelectrolyte is added to a kilogram of water and lowers the freezing point by 0.93°C. What is the approximate molecular mass of the substance?
a. 70.0 g b. 35.0 g c. 17.5 g d. 9.0 g
25. The boiling point elevation depends on all of the following except the
a. nature of the solvent. c. nature of the solute.
b. number of particles of the solute. d. amount of the solvent.
B. Problems
Solve the following problems in the space provided. Show all your work.
26. How can we prepare 0.25M NaOH solution?.
27. How many grams of CuSO4 are in 200-ml of 1.35M CuSO4 solution?
C. Essay Question
28. Suggest a possible reason why both the molal freezing point depression and the molal boiling point elevation are colligative properties.
29. Which of the following has components that are obviously different? (a) homogeneous mixture
(b) solution (c) colloid (d) heterogeneous mixture
30. Types of mixtures are classified according to (a) color; (b) particle mass; (c) relative quantities
of components; (d) particle size.
31. Molecules whose water solutions conduct current (a) dissociate in water; (b) ionize in water;
(c) do not dissolve in water; (d) decompose in water.
32. In order to conduct electricity in water solutions, a particle must be (a) charged and mobile;
(b) noncharged and mobile; (c) charged and nonmobile; (d) noncharged and nonmobile.
33. Which of the following mixtures contains particles that are in a dispersed phase and do not
settle out? (a) colloids (b) heterogeneous mixtures (c) solutions (d) suspensions
34. Colloids (a) are separable by filtration; (b) settle out on standing; (c) scatter light;
(d) are obviously heterogeneous.
35. Which of the following does NOT increase the rate of dissolving of a solid in water?
(a) raising the temperature (b) stirring (c) using large pieces of solid (d) crushing the solid
36. Raising the collision rate between solute and solvent (a) increases dissolving rate; (b) decreases
dissolving rate; (c) has no effect on dissolving rate; (d) can have any of the above effects.
37. If the amount of dissolved solute in a solution is greater than the amount that can
permanently remain in solution, the solution is (a) saturated; (b) unsaturated;
(c) supersaturated; (d) dilute.
38. Which of the following is likely to release crystals of solid from solution, if undisturbed?
(a) an unsaturated solution (b) a supersaturated solution (c) a saturated solution
(d) all of the above
39. In the expression "like dissolves like," the word "like" refers to similarity in molecular
(a) mass; (b) size; (c) energy; (d) polarity.
40. Hydrocarbons tend to be (a) polar; (b) nonpolar; (c) ionic; (d) hydrogen-bonded.
41. When the energy released by the forming of solvent–solute attractions is greater than the
energy absorbed by overcoming solute–solute and solvent–solvent attractions, a dissolving
process (a) has a negative heat of solution; (b) has a positive heat of solution; (c) occurs
rapidly; (d) does not occur.
42. The dissolving of gases in liquids is generally (a) endothermic; (b) exothermic; (c) rapid;
(d) impossible.
43. Henry's law relates (a) pressure to temperature; (b) pressure to gas–liquid solubility;
(c) temperature to gas–liquid solubility; (d) pressure to liquid–solid solubility.
44. As temperature increases, solubility of gases in liquids (a) increases; (b) decreases;
(c) can increase or decrease; (d) is not affected.
45. Which of the following is an expression of concentration? (a) molality (b) molarity
(c) percent concentration by mass (d) all of the above
46. Percent concentration by mass equals (a) moles solute per liter solution;
(b) moles solute per kilogram solvent; (c) moles solute per liter solvent;
(d) grams solute per 100 g solution, times 100.
47. The symbol M stands for (a) volume; (b) molality; (c) percent concentration by mass;
(d) molarity.
48. A saturated solution (a) is always concentrated; (b) is always dilute; (c) is neither concentrated
nor dilute; (d) may be either concentrated or dilute.







49. What is the solubility of silver nitrate if only 11.1 g can dissolve in 5.0 g of water at 20°C? 2.2 g
A) 100 g H20 at 20°C
45 g
B) 100 g H20 at 20°C
22.2 g
C) 100 g H20 at 20°C
222 g
D) 100 g H20 at 20°C
0.45 g
E) 100 g H20 at 20°C
50. What happens to the solubility of a gas, in a liquid, if the partial pressure of the gas above the liquid decreases?
A) The solubility decreases.
B) The solubility increases.
C) The solubility remains the same.
51. If the solubility of a gas in water is 4.0 g/L when the pressure of the gas above the water is 3.0 atm, what is the pressure of the gas above the water when the solubility of the gas is 1.0 g/L?
A) 0.75 atm B) 1.3 atm C) 4.0 atm D) 12 atm
52. Which of the following factors both affect the solubility of a particular substance?
A) temperature and the nature of solute and solvent
B) temperature and degree of mixing
C) particle size and degree of mixing
D) particle size and temperature
53. Which of the following operations usually makes a substance dissolve faster in a solvent?
A) agitation B) raising the temperature
C) crushing the substance to a powder D) all of the above
54. Increasing the temperature of a solution will generally
A) increase the rate at which a solute dissolves
B) increase the amount of solute that dissolves
C) both A. and B.
D) neither A. nor B.
55. Which of the following expressions is generally used for solubility?
A) grams of solute per 100 grams of solvent
B) grams of solute per 100 milliliters of solvent
C) grams of solute per 100 grams of solution
D) grams of solute per 100 milliliters of solution
56. Which of the following pairs of substances are miscible?
A) water and ethanol B) water and sodium chloride
C) water and oxygen D) water and gasoline
57. If a crystal added to an aqueous solution causes many particles to come out of the solution, the original solution was
A) saturated B) unsaturated C) supersaturated D) an emulsion
58. Which of the following are immiscible liquids?
A) ethanol and water B) acetic acid (vinegar) and water
C) gasoline and water D) sulfuric acid and water
59. Holding the temperature constant while adding more solute to a solution that already has solute crystals at the bottom of the container_______________________________
A) makes the solution more concentrated
B) causes the solution to become supersaturated
C) causes more solute crystals to appear at the bottom of the container
D) none of the above
60. Which of the following substances is less soluble in hot water than in cold water? A) CO2 B) NaCl C) NaNO3 D) KBr

61. What does not exist in a supersaturated solution? A) undissolved solute B) dissolved solute
62. What can be done to crystallize a supersaturated solution?
A) add almost any type of crystal
B) add a crystal of the solute only
C) add a crystal of the solute or scratch the glass
D) add almost any type of crystal or scratch the glass
C) solvent vapors

63. To increase the solubility of a gas at constant temperature from 0.85 g/mL, at 1.0 atm, to 5.1 g/mL, the pressure would have to be increased to .
A) 0.17 atm B) 5.0 atm C) 6.0 atm D) 4.3 atm
64. If the pressure of a gas above a liquid is increased (at constant temperature), the solubility of the gas in the liquid
A) remains unchanged B) increases
C) decreases D) would be impossible to calculate
65. The solubility of a gas in a liquid is ____
A) proportional to the square root of the pressure of the gas above the liquid
B) directly proportional to the pressure of the gas above the liquid
C) inversely proportional to the pressure of the gas above the liquid
D) unrelated to the pressure of the gas above the liquid
66. In general, as the temperature of a solution composed of a gas in a liquid is increased, the solubility of the gas_____________________
A) increases B) decreases C) remains the same
67. In a concentrated solution there is ____
A) no solvent B) a large amount of solute
C) a small amount of solute D) no solute
68. What is the molarity of a solution containing 9.0 moles of solute in 500.0 mL of solution? A) 4.5M B) 18M C) 0.45M D) 1.8M E) 0.18M
69. What is the number of moles of solute in 250 mL of a 0.4M solution?
A) 0.1 mol B) 0.16 mol C) 0.62 mol D) 1.6 mol E) 1 mol
70. What is the molarity of a solution containing 8 grams of solute in 500 mL of solution? (gram formula mass of solute = 24 g)
A) 1M B) 0.67M C) 0.1M D) 0.5M E) 0.05M
71. What mass of Na2SO4 is needed to make 2.5 L of 2.0M solution? (Na = 23 amu; S = 32 amu; = 16 amu)
A) 178 g B) 284 g C) 356 g D) 710 g
72. What is the molarity of 200 mL of solution in which 2.0 moles of sodium bromide is dissolved?
A) 2.0M B) 10M C) 0.40M D) 4.0M
73. How many mL of 3M HCl are needed to make 300 mL of 0.1M HCl?
A) 10 mL B) 100 mL C) 90 mL D) 9 mL E) 30 mL
74. If 2.0 mL of 6.0M HCl is used to make a 500.0-mL aqueous solution, what is the molarity of the dilute solution?
A) 0.024M B) 0.24M C) 0.30M D) 0.83M E) 2.4M
75How many mL of a 2.0M NaBr solution are needed to make 200.0 mL of 0.50M NaBr? A) 25 mL B) 50 mL C) 100 mL D) 150 mL