Safety in the Laboratory 1) Study your lab assignment before you come to the lab. If you have any questions, be sure to ask your teacher for help. 2) Do not perform experiments without your teacher's permission. Never work alone in the laboratory. 3) Use the table on the inside front cover of this textbook to understand the safety symbols. Read all CAUTION statements. 4) Safety goggles and a laboratory apron must be worn whenever you are in the lab. Gloves should be worn whenever you use chemicals that cause irritations or can be absorbed through the skin. Long hair must be tied back. 5) Do not wear contact lenses in the lab, even under goggles. Lenses can absorb vapors and are difficult to remove in case of an emergency. 6) Avoid wearing loose, draping clothing and dangling jewelry. Bare feet and sandals are not permitted in the lab. 7) Eating, drinking, and chewing gum are not allowed in the lab. 8) Know where to find and how to use the fire extinguisher, safety shower, fire blanket, and first-aid kit. 9) Report any accident, injury, incorrect procedure, or damaged equipment to your teacher. 10) If chemicals come in contact with your eyes or skin, flush the area immediately with large quantities of water. Immediately inform your teacher of the nature of the spill. 11) Handle all chemicals carefully. Check the labels of all bottles before removing the contents. Read the label three times: 1) Before you pick up the container. 2) When the container is in your hand. 3) When you put the bottle back. 12) Do not take reagent bottles to your work area unless instructed to do so. Use test tubes, paper, or breakers to obtain your chemicals. Take only small amounts. It is easier to get more than to dispose of excess. 13) Do not return unused chemicals to the stock bottle. 14) Do not insert droppers into reagent bottles. Pour a small amount of the chemical into a beaker. 15) Never taste any chemicals. Never draw any chemical into a pipette with your mouth. 16) Keep combustible materials away from open flames. 17) Handle toxic and combustible gases only under the direction of your teacher. Use the fume hood when such materials are present. 18) When heating a substance in a test tube, be careful not to point the mouth of the test tube at another person or yourself. Never look down the mouth of a test tube. 19) Do not heat graduated cylinders, burettes, or pipettes with a laboratory burner. 20) Use caution and proper equipment when handling hot apparatus or glassware. Hot glass looks the same as cool glass. 21) Dispose of broken glass, unused chemicals, and products of reactions only as directed by your teacher. 22) Know the correct procedure for preparing acid solutions. Always add the acid slowly to the water. 23) Keep the balance area clean. Never place chemicals directly on the pan of a balance. 24) After completing an experiment, clean and put away your equipment. Clean your work area. Make sure the gas and water are turned off. Wash your hands with soap and water before you leave the lab.

 

1. See http://www.chemistryclasses.com/Lab/lab_safety.htm

Lab Safety

You are responsible for your safety and the safety of your lab partner(s). Before, during and after a lab, check that you and your partner(s) follow proper safety procedures.

I. Dress

• Put on safety goggles for the entire duration of the lab.
• Wear long pants or long skirts (below the knee). Shorts and short skirts are not allowed.
• Wear close-toed shoes.
• Tie back long hair.
• Tuck in shirts, tuck in ties, and restrict bulky clothing.
• Wear a lab coat or lab apron.
• Remove watches, rings and bracelets.
• Wear gloves any time you touch chemicals, glassware, or chemical equipment. It is especially important to wear gloves when you are cleaning glassware and equipment.

II. Lab Conduct

• Never eat or drink in the laboratory.
• Come to lab in proper dress and prepared to work.
• Store bookbags and books away from the lab benches. You should only use your lab notebook during experiments.
• Use common sense!
• Follow all instructions.
• Do not perform unauthorized experimentation.
• Ask questions before performing experiments.
• Never leave experiments unattended.
• LABEL EVERY PIECE OF GLASSWARE using a permanent marker. As soon as you put a chemical into a flask, label it. If you change the contents, change the label.
• Keep your lab space and lab drawer clean and organized.
• Horseplay is not allowed in the laboratory.

III. Handling Chemicals

• Treat all chemicals as if they are hazardous and toxic!
• Wear gloves when handling chemicals or chemical bottles.
• Read all labels carefully.
• Clean spatulas before each use.
• Take only as much reagent as you need.
• Do not return unused reagent to stock bottles.
• Always put the lids back on bottles after using them.
• Never mix chemicals unless instructed to do so!
• When preparing solutions, add water to glassware before adding chemicals.
• Always measure chemicals (by mass or volume) before using them.
• Never taste chemicals.
• When smelling chemicals, gently waft odors towards your nose.
• Dispose of chemicals according to proper procedures.

IV. Handling Glassware

• Carry glassware carefully.
• Check your glassware for cracks and chips each time you use it. Cracks could cause the glassware to fail during use and cause serious injury to you or lab partners.
• Do not apply force when working with glassware. If something (stopper, thermometer, etc.) is stuck in a piece of glassware, take it to you chemistry teacher.
• Allow time for hot glassware to cool before touching it. Handle hot test tubes with test tube clamps.

V. Heating Substances

· Never use open flames in laboratory unless instructed by Ms. Bennett.

· Exercise caution when using gas burners. Keep all chemicals and clothing away from flames.

· Do not turn your back on flames.

· Extinguish the flame of gas burners by turning off the gas.

· Never heat a closed container.

· Do not look directly into a container while it is being heated.

VI. At the Conclusion of a Lab: Do the following in the listed order.

· Dispose of chemicals in proper containers.

· Wipe glassware with acetone, removing all labels and marker residue.

· Clean your glassware with soap and water, and rinse thoroughly.

· Rinse glassware with de-ionized water.

· Clean your lab space. Use spray cleaner if necessary.

· Remove solid waste from the sink and rinse the sink with water.

· Dispose of gloves.

· Hang up your lab coat or apron.

· Put away safety goggles.

· Wash your hands.

VII. Emergency Procedures

· Know the location of the fire extinguisher, fire blanket, eyewash, safety shower, and first aid kit.

· Know how to use fire extinguisher, fire blanket, eyewash, safety shower, and first aid kit.

· Notify Ms. Bennett immediately after an injury, explosion, fire, or spill.

· In case of broken glass, pick up large pieces carefully, then sweep up small pieces. Dispose of broken glass in the broken glass container, not the trash can.

Worksheet attached titled- Laboratory Safety: Rules and Regulations for all Labs

2. Finding and Interpreting MSDS information to address first aid, clean up of spills, handling, storage, etc.
3. See any MSDS source sheet (e.g.) Flinn MSDS Collection at http://www.flinsci.com/homepage/sindex.html

MSDS sheet attached

Questions: 
1. What do the letters MSDS stand for?

      Material Data Safety sheet

2. What is the name of this chemical?

                  

Potassium Chlorate

3. What is its chemical formula?

      KCLO3

4. What safety equipment is needed when handling this substance?

Gloves, chemical safety goggles, faceshield, smock, apron, eye wash station, ventilation hood, proper gloves, fire extinguisher

5. What type of substance should be used to put out a fire involuving this chemical?

      Non-flammation

6. If a person swallows this substance, what are the procedures that should be followed?

Call physician or poison control center immediately; induce vomiting only if advised by appropriate medical behavior

7. Which procedures should be followed if this substance gets on your skin?

Remove contaminated clothing flush thoroughly with mild soap and water. If irritation occurs, get medical attention

8. If this chemical is spilled, what procedure should be followed to clean it up?

Sweep up material , then place into a suitable container for disposer. Wash spill area for soap and water.

9. What are the potential health effects of splashing this chemical into eyes?

Check for and remove contact lenses. Flush thoroughly with water for at least 15 minutes, lifting upper and lower eyelids occasionally. Get immediate medical attention.

10. What precautions should be followed when storing this chemical for future use?

Store in a cool, dry place away from flammable and combustible materials. Keep away from heat, sparks, and flame. Use with adequate ventilation.

11. What is the purpose of an MSDS?

      To provide precautions and procedures in case of an accident. 

4. Be familiar with common units of measurement.

Meter (m)-length

Liter (L)-volume

Gram (g)-mass

Scientific notation: move decimal until there is only one none zero in front of the decimal.

When reducing number [moving decimal to left] exponent is postive; vice versa 

 
 

5. Know how to read instruments ACCURATELY.

A graduated cylinder is used to measure liquid volume. The unit is the milliliter. Place the graduated cylinder on a flat surface and view the height of the liquid in the cylinder with your eyes directly level with the liquid. The liquid will tend to curve downward. This curve is called the meniscus. To get the most accurate reading, read the meniscus. Calibrations are the small line markings on the cylinder. Ex: markings every 1 mLà 1.0

Thermometer: read it to ½ of the smallest calibration. Estimate to one digit beyond the markings. Same on ruler and other laboratory equipment

Significant Rules:

All non-zero digits are considered significant.

If a zero is between two non-zero digits then it is significant.

Leading zeros are never significant.

Trailing zeros are only signiicant if there is a decimal present.

Percent error: the ratio of an error to an accepted value

Percent error=error/accepted value *100

When caculating ignore plus and minus signs

Error= difference between experimental value and accepted value

6. Be familiar with safe lab practices

1. List some common safe lab practices--ex--when to wear goggles, how to use a Bunsen burner

Mortar- vessel in which substances are crushed or ground with a pestle

Pestle- club-shaped, hand-held tool for grinding or mashing substances in a mortar

Stirring Rod- rod used for stirring

Evaporating Disk- disk used to set out stuff to evaporate

Watch Glass-  concave dishes that can be used as beaker lids; can hold protists and other invertebrates for viewing under a microscope; or to dissolve materials such as crystals and powders; very handy for making ice lenses

Well Plate- used for seeing reactions of different solutions

Beaker- glass preferably; can be used for routine mixing, measuring boiling

Erlenmeyer Flask- commonly used for simple measurign, storing and mixing of liquids

Test Tube Holder- used to hold test tubes

Striker- strikes flint using friction to light bunsen burner

Buret- used where accurate amounts of liquid must be added in small amounts

Volumetric Flask- can help measure solutions accurately

Graduated Cylinder- measures solutions

Tripod Stand- stand with three legs; used to hold substances being heated over a bunsen burner

Ring Stand- holds test tubes to be heated

Spatula- used to pick up items

Tongs- used for picking up items

Bunsen Burner: to make flame hotter [blue] open air mixture valve accordingly//cooler [more orange] , close air mixture valve accordingly 

2. Apply the principles of experimental design in laboratory and field investigations. (2A, 2D, 2E)

1. Scientific Methods as a Systematic Approach pp. 10-13

Scientific Method: systematic approach used in scientific study

Qualitative data- descriptive data

Quantitative data- numerical data

Hypothesis: prediction of outcome of data

Experiment: used to test hypothesis

Independent Variable: variable altered

Dependent Variable: changes depending on altered changes of independent variable

Control: group used for comparison

Theory: explanation supported by many, many experiments

Scientific Law: relationship in nature supported by many experiments

2. Proper Use and Selection of Lab Equipment and Correct Scale Interpretation

Measure to half a calibration

More calibrations= more accurate

3. See http://lhs.lps.org/staff/squiring/chemistry/intro/useoflabequipment.pdf

Printed out and attached

4. See http://www.hmpublishing.com/brown/chapter%204.pdf

Printed out and attached

II. Chapter 2--Data Analysis--TEKS Oj. 2B, 2C, 2D

3. Using the factor-label method, solve problems involving significant figures, SI units and scientific notation.

1. SI Units of Measurement- Base SI Units, pg. 26 Table 2-1, and Derived Units pp. 25-30

Base unit: defined unit in a system of measurement that is based on an object or event in the physical world

 

• 
  

Derived unit: unit that is defined by a combination of base  unites. Ex: for speed SI unit is meters per second (m/s). Volume and Density are derived as well.

Volume: space occupied by an object (cm^3)--two base  units are used to find it

• 
  

Density: ratio that compares mass of an object to its volume (g/cm^3)

2. Scientific Notation pp 31-33: pp 889-892 Math Handbook

• 
  

Scientific notation expresses numbers as a multiple of two factors: a number between 1 and 10; and ten raised to a power, or exponent

3. Dimensional Analysis / Metric Conversions pp 34-35/ pp. 900-902 Math Handbook

• 
  

conversion factor: ratio of equivalent values used to express the same quantity in different units

dimensional analysis: method of problem-solving that focuses on the units used to describe matter

Be able to make conversions with various metric units, including conversions involving squaring

Ex: area- 10m * 12m = 1.2 cm

Pg 51: problems 80, 85; page 52 problems 90, 93; page 53 problems 1-3

Page 51:

• 
  

80. A. 5700 B. 0.0437 C. 783000 D. 0.0453 E. 1000 F. .0375

85. A. 3.96 * 10^3 B. 82.0 * 10^-4 C. 74.8 dm D. 13.81 cm E. 17.2 mg F. 4.3 * 10^5 G. 8.097 km

• 
  

90. No, units do not cacel out

93. A. 301 cg B. 6.200 km C. 6 240 000 000 * 10^-7 D. 2 dm^3 E. 0.00013 kcal/g F. 0.00321 L

• 
  

Page 53:

1. C

• 
  

2. C

3. C

Accuracy and Precision of Measurement pp. 36-37

Accuracy: refers to how close a measured value is to an accepted value

• 
  

Precision: refers to how close a series of measurements are to one another

7. Percent Error Calculations pp 37-38

• 
  

Percent error is the ration of an error to an accepted value

Percent error=error/accepted value * 100

• 
  

Error=experimental value - accepted value

Plus and minuses do not matter

Significant Figures Determination pp. 38-42; pp 893-896 Math Handbook

Significant figures: include all known digits plus one estimated digit

• 
  

Rules for recognizing significant figures

Non-zero numbers are always significant.

2. 
   

Zeros between non-zero numbers are always significant.

All final zeros to the right of the decimal place are significant.

4. 
   

Zeros that act as placeholders are not significant. Convert to scientific notation to remove the placeholder as zeros. 0.0

253

   

432

0..only highlighted are sig figs

5. 
   

Counting numbers and defined constants have an infinite number of sig figs. Ex. 6 molecules.. 60 s= 1 min

Rules for Rounding Numbers

1. 
   

If the digit to the immediate right of the last sig fig is less than five, do not change the last sig fig

If the digit to the immediate right of the last sig fig is greater than five, round up

3. 
   

If the digit to the immediate right of the last sig fig is equal to five and followed by a nonzero digit, round up last sig fig

If digit to immediate right of last sig fig is five and is followed by zero, round up if odd, do not roudn up if it is even

• 
  

Addition and Subtraction/Multiply and Divide

answer must have the same number of digits to the right of the decimal point as the value with the fewest

• 
  

answer must have the same number of significant figures as the measurements with the fewest

9. Know the SI base units- p 26: Table 2-1

• 
  

Base unit is a defined unit in a system of measurement that is based on an object or event in the physical world

Time- second (s)

• 
  

Length- meter (m)

Mass- kilogram (kg)

• 
  

SI Base Units

• 
  

4. Determine density by interpreting data on a graph and by using problem solving strategies.

1. Density Calculations/Interpreting Density Graphs pp. 27-29

• Density: a ratio that compares the mass of an object to its volume. Unit are grams per cubic centimeter (g/cm^3)
• Density= mass(g) / volume(cm^3)

 
 
 
 
 

1. Set up a graph that would appropriately display density…which information would go on which axis? How would you determine the density from the graph? Page 52 do problem 87 using Table 2-7

• Mass- y axis, volume- x axis

• You would determine the density by dividing the mass by the volume
• Page 52, number 87 Table 2-7
• Slope is 2.7, graph on calculator equation y=2.7x

2. If you needed to determine the density of an irregularly shaped object (like in the density lab), outline a brief procedure to determine it. See

http://www.middleschoolscience.com/irregularvolume.htm

• 1) Obtain a beaker of colored water and pour some into the graduated cylinder. Record initial volume in mL.

2) Place object in graduated cylinder and record volume in mL.

3) Find volume of object by subtracting initial volume and final volume in mL.

4) Weigh object using triple beam balance in grams for the mass.

5) Divide the mass in grams by the density of mL or cm^3.

1. Page 29 problems 1, 3; page 30 problem 10

• Page 29 # 1,3

• 1) 7 g/mL            3) D= 5, it is not made of pure aluminum because density is an intensive property of aluminum with a density of 2.7
• Page 30 # 10
• 10) Oil floats on water because it is less dense.

2. Correct Set Up, Reading and Interpretation of Graphs pp. 43-45/ 903-907 Math Handbook

• Circle graph: sometimes called a pie chart; useful for shownig parts of a fixed whole; usually labeled as percentiles
• Bar graph: used to show how a quantity varies with factors such as time, location, or temperature
• Line graph: ind and dep variable
• If linear- variables directly related
• Interpreting graph: first identify independent and dependent variables
• Linear or nonlinear?
• Interpolation- you can read data from a graph that falls between measured points
• Extrapolation- you can extend the line beyond the plotted points and estimate values for the variables

III. Chapter 3--Matter--Properties and Change--TEKS Obj. 4A, 4B, 4C, 5A

4. Differentiate between physical and chemical properties of matter, as well as physical and chemical changes in matter. (4A)

1. Properties of Matter: Intensive vs Extensive/Physical vs Chemical pp 55-59; pp 61-62; pg 82: 39, 41, 70; page 85 #4, 10

• Substance: matter that has a uniform and unchanging composition; also known as a pure substance
• Physical property: characteristic that can be observed or measured without changing sample's composition
• Ex. Density, color, odor, taste, hardness, melting and boiling poing
• Extensive properties: dependent on amount of substance present
• Ex. Mass, length, volume
• Intensive properties: independent of amount of substance present
• Ex. Density
• Chemical property: ability of a substance to combine with or change into one or more other substances
• Ex. Ability of iron to form rust when combined with air; inability of substance to change into another substance
• Physical changes: changes that alters substance without changing composition
• Ex: fermentation, rust, explode, oxidize, corrode, tarnish, burn, rot

2. You could determine if it was a chemical or physical change by checking for other signs of a chemical change including ferment, rust, and oxidize.

3. 
   

1. Law of conservation of mass: mass is neither created nor destroyed during a chemical reaction--it is conserved
2. Law of Definite and Multiple Proportions pp. 75-77:

• Regardless of amount, a compound is always composed of the same elements in the same propportion by mass

• Percent by mass: total mass of compound as sa percentage
• Percent by mass (%) = mass of element/mass of compound * 100

5. Classify and describe matter as it relates to solid, liquid, and gas( 7B) and element, compound, and mixture

1. States of Matter: all matter that exists on Earth can be classified as one of these physical forms

• Solid: form of matter that has its own definite shape and volume

• Particles are tightly packed; when heated solid expands
• Liquid: form of matter that flows; constant volume
• Particles less closely packed than solids; are able to move past each other
• Like solids, tend to expand when heated
• Gas: form of matter that flows to conform to shape of container
• Particles very far apart but easily compressed
• Vapor: gaseous state of solid or liquid at room temperature

B.

• mixture: combination of two or more pure subsstances in which each pure substance retains its individual chemical properties
• heterogeneous mixture: does not blend smoothly throughtout and some individual substances remain extict. Ex: sand and water
• homogeneous mixture: constant change in compsotition throughout
• solutions: homogenous mixtures
• filtration: technique used as a porous barrier to separate a solid from a liquid
• crystallization: seperation technique that results in formation of solid particles under a substance
• chromatography: technique that separates components of a mixture (called moble phase) and on the basis of the tendency of each to travel ro to be drawn across surface of another material (stationary phasae)
• element: pure substance that cannot be separated into simpler substances by physical or chemical means
• periodic table: first version of modern table
• compound: isi a combination of two or more different elements that are combined chemically

 
 
 
 

• element-

 
 
 
 

• compound-

 
 

• Mixture

5. periodic table
6. 43. Physical 46.mixture: combination of two or more pure substances in which each pure substance retains its own individual chemical properties 48. C 49. C

 

IV. Chapter 4--Structure of the Atom--TEKS Obj. 3E, 6, 6B, 11B

4. Summarize the history of chemistry in terms of significant scientists' experiments and their development of the atomic model. (3E, 6A)

A. Early Theories (Democritus, Aristotle, Dalton, Crooke, Thomson, Millikan, Rutherford, Chadwick) pp. 87-97

 
 
 
 
 
 
 

  Aristotle 
•·  One of the most influential philosophers 
•·  Wrote extensively on many subjects, including politics, ethics, nature, phsycis, and astronomy. 
•·  Most of his writings have been lost through the ages. 
 

  Democrituss Ideas 
•·   Matter is composed of empty space through which atoms move. 
•·  Atoms are solid, homogeneous, indestructible, and indivisible. 
•·  Different kinds of atoms have different sizes and shapes. 
•·  The differing properties of matter are due to the size, shape, and movement of atoms. 
•·  Apparent changes in matter result from changes in the groupings of atoms and not from changes in the atoms themselves. 

1. Review yellow boxes on pp. 88, 89

 

  Daltons Atomic Theory 
•·  All matter is composed of extremely small particles called atoms. 
•·  All atoms of a given element are identical, ahving the same size, mass, and chemical properties. Atoms of a specific element are different from those of any other element.  
•·  Atoms cannot be created, divided into smaller particles, or destroyed.  
•·  Different atoms combine in simple whole number ratios to form compounds. In a chemical reaction, atoms are separated, combined, or rearranged. 
 
 
 

4. Distinguish between atoms and their isotopes including:

1. Characteristics and Properties of Subatomic Particles--including location, charge and relative mass pp. 92-97

• Electrons: negatively charged particles part of all forms of matter
• Found outside of nucleus
• 9.1 X 10^-28 g = (1/1840 of a hydrogen atom)
• Nucleus: tiny, dense region centrally located within atom that contained all of an atom's positive charge and virtually all of its mass
• Protons: postive charge of 1 +
• Found in nucleus
• 1.673 X 10^-24 = (1 of a hydrogen atom)
• Neutrons: mass nearly equal to that of proton but no charge

B. Atomic Number, Atomic Mass, Isotopes, Isotope stability (6A, 6B) pp. 98-104; pg 836 problem 44

• Atomic number: number of protons in an atom is referred to as the element's atomic number
• Isotopes: atoms with same number of protons but different number of neutrons
• Mass number: sum of number of protons and neutrons in the nucleus
• Atmoic mass  unit (amu): 1/12 mass of carbon-12 atom. Anearly equal to mass of a single proton or single neutron
• Atomic mass: weighted average mass of isotopes of that element
• Page 836 #44
• The most stable ratio from the ratio of neutrons-to-protons is 1:1. The larger the atom with this ratio, the more stable the atom is.

2. Radioactivity/Types of Radiation pp. 105-107; pp 806-809

• Nuclear reactions: reactions that involve a change in an atom's nucleus
• Radioactivity: process where some substances spontaneously emit radiation
• Radiation: rays and particles emitted by radioactive material
• Radioactive decay: unstable nuclei lose energy by emitting radiation in a spontaneous process (process that does not require energy)
• Alpha radiation: radiation deflected toward negatively charged plate; radiation made up of alpha particles which have two protons and two neutrons and a charge of 2+
• Nucleur equation: shows atomic number and mass number of particles involved
• 226/88 Ra à 222/86 Rn + 4/2He
• beta radiation: radiation deflected towards the positively charged plate. Radiation cosits of electronsà beta particles
• 14/6Cà 14/7 N +0/-1B
• gamma rays are high-energy radiation that possess no mass and are denoted by symbol 0/0 Y
• radioisotopes: isotopes of atoms with  unstable nuclei

3. Pg 112 problem 48, 49, 50; pg 113 problem 72; complete the table

• Page 112 #48-50
• 48) They differ in the nucleus with the number of neutrons and mass number but  have the same number of protons
• 49) numbers of protons neutrons is the mass number
• 50) superscript is the mass number and subscript is the atomic number
• Page 113 #72

• Complete the Table

 

V. Chapter 25--Nuclear Chemistry--TEKS Obj. 9A, 9B, 9C, 9D, 11B

10. Using nuclear reactions be able to calculate half lives (9B), write and balance nuclear equations. (11B, 11 C), and apply properties of radioactive particles.

1. Types of Nuclear Radiation- alpha, beta, gamma, alpha, positron pp. 806-809

• Radioisotopes: isotopes of atoms with unstable nuclei; emit radiation to attain more stable atomic configurations in process of radioactive decay
• Radioactive decay: unstable atoms lose energy by emitting one of several types of radiation, three most common types are alpha, beta, and gamma
• Ernest Rutherford- gold foil experiment that helped define modern atomic structure, identified alpha, beta, and gamma radiation when studying effects of an electric field [diagram pp. 807]
• Alpha
• Has same composition as a helium nucleus--two protons and two neutrons--given symbol 4/2 He
• 226/88 Ra  -à222/86Rn    +   4/2 He 
  Radium-226    Radon-222    Alpha Particle    
• slow moving b/c of mass and charge, so not very penetrating
• Beta
• Very fast moving that has been emitted from a neutron of an unstable nucleus
• Represented by symbol 0/-1 B
• 0= no mass, -1 = charge
• 1/- n à 1/1 p       +      0/-1 B 
  neutronà proton + beta decay [[leaves behind a proton]]
• 131/53  I   à    131/54 Xe   + 0/-1 B 
  Iodine- 131àXenon-131        Beta particle
• Gamma
• High energy (short wavelength) electromagnetic radiation
• Does not change atomic number or mass number of nucleus
• Almost always accompany alpha and beta radiation and account for most of the enrgy loss that occurs as a nucleus decays
• 238/92 U   à  234/90 Th     + 4/2 He    + 2 0/0 Y

2 in front of gamma means that two gamma rays of different frequencies are being omitted

• x rays: form of high-energy electromagnetic radiation; are not produced by radioactive srouces but emitted from certain materials that are in an excited electron state; only blocked by lead and concrete; highly penetrating

2. Page 819 # 17-19. Be able to interpret ½ life graphs--ex Fig 25-13 page 817; Page 839 # 2,3
3. Radioactive Decay/Nuclear Stability pp. 810-812

• Nucleons: protons and neutrons in nucleus

• Strong nuclear force: force that acts only on subatomic particles that are extremely close together and it overcomes the electrostatic repulsion between protons
• Band of stability: arrea on graph with stable nuclei
• Positron emission: radiactive decay process that involves emission of a positron from a nucleus
• Positron: particle with same mass as an electron but opposite charge; symbol = 0/1 B
• Electron capture: occurs when nucleus of an atom draws in a surrounding electron, usually one from lowest energy level
• 1/1 p + 0/-1 e--à 1/0 n

4. Writing and Balancing Nuclear Equations pp. 813-814

• in balanced chemical equations, numbers and kinds of atoms are conserved; in balanced nuclear equations, mass numbers and atomic numbers are conserved
• step by step pp. 813

5. Transmutation and Calculating Radioactive Decay Half-Life Rates pp. 815-818

• Transmutation: conversion of an atom of one element to an atom of another element
• Induced transmutation: process, which involves striking nuclei with high-velocity charged particles
• Transuranium elements: elements immediately folowing uranium on periodic table, with atomic numbers 93 and greater
• Have been produced in laboratory by induced transmutation and are radioactive
• Half-life: time recquired for one-half of radioisotope's nuclei to decay into its products
• Amount remaining= (Initial amount) (1/2) ^n            
• n=number of half-lives that have passed
• n=t/T
• t= elapsed time; T is the duration of the half-life

11. Differentiate nuclear fission and nuclear fusion in terms of masses, reactants, products, amount of energy, and their properties. (9A)

1. Fission and Fusion of Atomic Nuclei pp. 821-826

 

12. Evaluate the commercial use of nuclear reactions and the environmental implications. (3C, 9C, 9D)

1. Problems Associated with Nuclear Waste Disposal pp. 824-826

• CFCs

• Pollution,recquires high energies to initiate and sustain
• Confinement of reaction

2. Applications and Effects of Nuclear Reactions pp. 827-831
1. List several ways that radiation can be used in a positive way

• Detect trace amounts of elements present in sample

• Computer chip manufactors to annalyze composition of highly purified silicon wafers
• Important uses in medicine

 

VI. Chapter 5- Electrons in Atoms (OMIT section 5.1 and 5.2) NO TEK Obj.

4. Using electron configurations, identify atoms and ions

1. Ground state and ion electron configurations pp. 135-140

• Electron configuration: arrangement of electrons in an atom

1. Aufbau Principle, Pauli Exclusion Principle, Hund's Rule

• Aufbau: states that each electron occupies lowest energy orbital available
• Pauli: max of two electrons may occupy a single orbital if they have opposite spins
• Hund's: single electrons with same spin must occupy each equal-energy orbital before additional electrons with opposite spins

2. Orbital Diagrams

• Represents an atom's electron configuration
• Includes box for each of atom's orbials; arrows for electrons of different spins

 

VII. Chapter 6--The Periodic Table and Periodic Law--TEKS Obj. 3E, 4D, 6C