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Chemistry Periods 3 and 4

Chemistry Labs
Chemistry Scanned Assignments
Chemistry Periods 1 and 2
Chemistry Periods 3 and 4
Chemistry Periods 5 and 6
Chemistry Periods 7 and 8

August 30, 2010

Cover books (9-3-10)
Rules and Regulations Signed and Returned (9-3-10)
Lab Safety Contract (9-3-10)
Quiz Ions 1-18 (9-17-10)
Student Info Sheets (9-3-10)
Text book scavenger hunt (9-1-10)
Lab equipment sketches (9-7-10)
Introduction to Chemistry/ rules and regulations

August 31, 2010

Text book scavenger hunt (9-1-10)
Read and Outline 1.3 and 1.4 (9-1-10)
Rules and Regulations Signed and Returned (9-3-10)
Cover books (9-3-10)
Lab Safety Contract (9-3-10)
Student Info Sheets (9-3-10)
Lab equipment sketches (9-7-10)
Quiz Ions 1-18 (9-17-10)
Period 3:
Lab: Lab equipment sketches
Period 4:
Math Pre-Test

September 1, 2010

Rules and Regulations Signed and Returned (9-3-10)
Cover books (9-3-10)
Lab Safety Contract (9-3-10)
Student Info Sheets (9-3-10)
Lab equipment sketches (9-7-10)
Quiz Ions 1-18 (9-17-10)
Read and Outline 1.3 and 1.4 (9-2-10)
Period 3:
Go over math pretest
Period 4:
No class

September 1, 2010

Quiz Ions 1-18 (9-17-10)
2.2 read and outline/ possible quiz (9-8-10)
Understanding Equivalences odds (9-8-10)
Period 3:

Units of Measurement


-         All measurements must include both a unit and a number.

o       Without the unit, the number has no meaning.

-         English vs. Metric System:

o       English system – feet, inches, etc are not used in science.

o       Metric system- the international system of measurement is used

         Common language for all scientists

         Easy conversions


SI Base Unit






Physical Quantity


Unit Name and Symbol



kilogram, kg



meter, m



second, s

count, quantity


mole, mol



kelvin, K

electric current


ampere, A

luminous intensity


candela, cd




Derived Units Commonly Used in Chemistry




Physical Quantity


Unit Name and Symbol



square meter



cubic meter



newton, N



pascal, Pa



joule, J



watt, W



volt, V



hertz, Hz

electric charge


coulomb, C


The International System of Units (SI)

-         Seven base units (shown above)



o       Length: distance that light travels in a vacuum during a time interval 1/299,792,458 of a second

o       Mass and weight:

         Mass: amount of material- about 2.2 lbs at sea level

         Weight: influence of the force of gravity on mass

o       Area and Volume (derived units – combinations of base units)

         Area = Length x Width

5.0 m x 3.0 m = 15 m2

        Both units and numbers are multiplied in the answer

         Volume: amount of space that an object occupies


Non- SI Units Used Frequently in Chemistry

-         Volume: liter, L (there are exactly 1000 L in one cubic meter)

-         Pressure: atmosphere, atm; millimeters of mercury, mm Hg

-         Temperature: Celcius degree

-         Energy: calorie, cal


Metric Prefixes

-         Prefixes added to the base unit that make the units larger or smaller


Prefixes that make the Unit Larger

o       kilo (1 km = 1000 m)

o       mega (1Mm = 1000000 m)


Prefixes that make the Unit Smaller

o       deci (1 dm = .1 m   or   10 dm = 1m)

o       centi (1cm = .01m)

o       milli (1mm = .001 m)

o       micro (1mm = .000001 m)

o       nano (1 nm = .000000001 m)

o       pico (1pm .000000000001 m)


1-5 Uncertainty in Measurement


*** When making a measurement, you must give all the certain (or exact) digits that the instrument can give + one additional “uncertain” digit that you estimate***


-         Measuring instruments always have some built in flaws

-         Measurement always involves some estimation


On an electronic balance:

-         The measurement is done for you.  The last number on the screen is the uncertain digit

o       Be sure to include the unit as well as the measurement


On a scale:

-         Imagine a graduated cylinder- measures volume

o       Liquids curve in the cylinder (known as the meniscus)

o       Measurements are taken from the bottom of the bend

-         The certain digits are those given on the graduated cylinder.

o       The uncertain digit is found by reading between the lines of the instrument


The Uncertainty of a Measurment

-         Generally, the uncertainty of a measurement reflects the value of the uncertain digit:

Suppose the measurement was: 32.7 mL

         The seven in this measurement is uncertain.

         Since the seven is uncertain, it could possibly as high as an 8 or as low as a 6.

         Therefore the uncertainty is +/- .1


Period 4:

Problem Solving


-         Dimensional Analysis: a technique of converting between unit


-         Unit Equalities: an equation that shows how different units are related



 The unit equality between gallons and liters:

 1 gal = 3.774 L


-         Conversion Factors: derived from the unit equalities


Look at the above unit equation.  Divide both sides by 1 gallon


                        1 gal = 3.774 L


                        1 gal/ 1 gal = 3.774 L/ 1 gal


                        1 = 3.774 L/ 1 gal

A second unit factor can be made if both sides were divided by 3.774 liters.


                        1 gal = 3.774 L


                        1 gal / 3.774 L = 3.774L/ 3.774 L


                        1 gal / 3.774 L = 1


-         Using Conversion Factors:


Suppose you wanted to know how many liters were in 250 gallons


                        This is simple- ready???


                        V =250 gal ( 3.774 L / 1 gal)


                        V= 940 L


Notice that there were two possible conversion factors to use… only one conversion factor would cause the units to cancel


Dimensional Analysis Applied to Unit Conversions


Step 1:

Write down the unit equality that you need to convert


Step 2:

Write down the conversion factors corresponding to the unit equality


Step 3

Use a symbol to represent the unknown quantity on the left side of the equation


Step 4:

Write the known information on the right side of the equation.  Use the appropriate conversion factor (Units should cancel)


Step 5:

The unit on the right side of the equal sign should be the unit you are converting to.



More than one conversion factor may be used in any calculation to get to the desired unit.

September 1, 2010

Quiz Ions 1-18 (9-17-10)
Unit Analysis Wkst b and c (9-14-10)
Period 3:
Reading Quiz 2.3 and 2.4

2-2 Temperature:


-         Why can not the sense of touch be used to measure temperature?


-         A thermometer is an instrument that gives and accurate and precise reading of temperature.


-         Galileo Galilei (1564-1642) – invented the first temperature instrument

o       Modern thermometers have a bulb filled with mercury or colored alcohol attached to a stem

o       Heating causes liquid to expand and move up the stem

o       Cooling causes liquids to condense and move down the stem


The Fahreheit and Celsius Temperature Scales


-         Gabriel Fahrenheit- Made thermometers in the late 1600’s and early 1700’s- made up his own temperature scale.

-         Anders Celsius (1701-1744) developed a scale much more in tune with the metric system

o       Freezing point at sea level = 0 Boiling point at seas level = 100


The Kelvin Temperature Scale


-         The SI scale used to measure temperature is the Kelvin Scale

-         Lord Kelvin (English- 1824-1907) : unit K

o       A degree change of 1 K is the same as a degree change of 1 C

o       Zero point in the Kelvin scale corresponds to absolute zero (-273 C)

         Absolute zero is where molecular motion stops


-         Some Equations

o       C= K – 273


o       K= C + 273


o       (oF-32oF) x (100oC/180oF) = oC


o       (oC x 180oF/100oC) + 32oF = oF


2-3 Matter


-         Matter is the “stuff” of which things are made

o       Has mass (amount of stuff) and volume (amount of space)

-         Do not know where the “stuff” came from, but it is here and we have learned a lot about the “properties of matter”

o       Has been a philosophical issue for millennia


States of Matter:

-         Four States of Matter:

o       Solid

o       Liquid

o       Gas

o       Plasma


-         Properties of the different states (generalized):

o       Solid:

         High density

         Density affected little by changes in pressure

         Shape not affected by the shape of a container

         Orderly arrangement of particles (ie. Crystals)

o       Liquid:

         High density

         Density affected little by changes in pressure

         Adopts the shape of the container

o       Gas:

         Low density

         Density depends on the pressure

         Expands to fill the container

o       Plasma

         Low density

         Density depends on pressure

         Expands to fill the container

         Exists only at high temperatures


Changes in State:

-         Can observe changes in states by heating or cooling a substance


Ex:       Water at 0 C is changing from liquid to solid

            Water at 100 C is changing from liquid to gas

            Water from 0 C to 100 C is in the form of a liquid


Ex        Mercury at –39 C is changing from a liquid to a solid

            Mercury at 357 C is changing from a liquid to a gas

            Mercury from –39 C to 357 C is a slippery liquid


Properties of Matter:

-         A sample of matter can be identified by observing its characteristics or properties


-         Physical Properties: properties that can be observed without changing the identity of the substance. (density/ color/ melting point, etc)

1)      State: (at standard temperature and pressure): Liquid, Solid, Gas

2)      Quantity: mass, volume, density

3)      Color

4)      Texture

5)      Melting and boiling points

6)      Conductivity

7)      Solubility in Water

-         Chemical Properties: properties that cannot be observed without changing the identity of the substance (flammability, etc)

1)      Evolution of a gas

2)      Formation of a precipitate

3)      Absorb or Gives off heat

4)      Emission of light

5)      Color Change


Atomic Number: number of protons in a nucleus (is equal to the number of electrons)


Atomic Mass: average of all naturally occurring isotopes


Periodic Law: physical and chemical properties of an atom are periodic functions of the atomic number.



Changes in Matter

-         Physical Changes: changes that do not alter the identity of the substances

o       Crushing, tearing, and changes in state


-         Chemical Changes: changes that do alter the identity of the substance

o       Change in the chemical make up of a substance.


Conservation of Matter:

-         Antoine Lavoisier: “one may take it for granted that in every reaction there is an equal quantity of matter before and after”


- Antoine Lavoisier: 1800’s (1743-1794)


-         Mass of substances before a chemical change was always equal to the mass of substances after the change.

-         Conclusion:

o       Matter was neither created nor destroyed during a chemical reaction.


Became known as the Law of Conservation of Mass



                                    By mass, 1 g of H always binds with 8 g of O


                                                So 2 g of H will bind with 16 g of O


                                                            3 g of H will bind with 24 g of O


-         Since you know the mass of both reactants, you can figure out the mass of the products:

o       1g H + 8g O = 9g water


-         Knowing this, you can reverse the reaction:

o       Electrolysis- using electricity to break water


45.0 grams of water: broken via electrolysis yielded 5.0g H and how many grams O?


            45.0g water- 5.0g H= 40g O


** one of the most important principles in chemistry**


-         Lavoisier got the ax during the Reign of Terror that followed the French Revolution

2.4 Elements and Compounds



-         An element is a substance that cannot be separated into simpler substances by a chemical change

o       Over 100 known elements

o       Named for famous people, states, planets, countries etc.

-         Element Symbol: a one or two letter abbreviation

o       First letter is always capitalized; second letter is always lower case

o       Most abbreviations come from the English name, others come from the Latin origin.

         Copper – cuprum Cu

         Gold – aurum Au

         Iron – ferrum Fe

         Lead – plumbum Pb

         Mercury – Hydragyrum Hg

         Potassium – Kalium K

         Silver – Argentum Ag

         Sodium – Natrium Na

         Tin – Stannum Sn

         Tungsten – Wolfram W


-         Periodic table- simple, harmonic, rhythmic way of organizing the elements by innate properties



-         A compound is a substance that contains two or more elements combined in a fixed proportion

-         Chemical Symbols are used to represent compounds- merely putting the element symbols in a specific order which notes the number of each element present.


Distinguishing between elements and compounds

-         Elements and compounds are pure substances

o       Has a unique set of chemical and physical properties

o       Separation techniques like electrolysis help distinguish between the two

o       Careful measurements of mass help distinguish as well.


2-5 Mixtures


-         A mixture is a blend of two or more pure substances


Types of Mixtures:

-         A mixture that has visibly different parts is called a heterogeneous mixture

-         A mixture that does not have visibly different parts is called a homogeneous mixture


Separating the Components of a Mixture:

-         Filtration- separation of heterogeneous mixtures of liquids and solids

-         Distillation- separation of homogeneous mixtures of liquids based on different boiling points (one changes to gas form first)

-         Distillation may also be used to separate impurities from liquids- solids are left behind

-         Crystallization produces solids of very high purity by evaporating the liquid component.

-         Chromatography- separation by flowing along a stationary substance.]


Unit Equations and Unit Factors


Based on equivalent relationships

-         statement of a relationship between two quantities that are equal

-         will be used during unit conversions


Example: 1 dime = 10 pennies


Unit Equations:

            Is a series of two equivalent quantities


            Ex:       1 dime = 10 pennies

                        10 pennies = 1 dime


Unit Factor:

            Ratio of two equivalent quantities


            Ex: 1dime/ 10 pennies or 10 pennies/ 1 dime


            Both the quantity and the reciprocal are true.


Exactly Equivalent Equations:

These equations are equivalent by definition, such as 1 foot is equal to exactly 12 inches.


As a result, rules for significant figures do not apply to these quantities are not considered when rounding for significant figures when doing the calculation.


A three lined equal sign is used to express these quantities.


            Unit Analysis:

                        Also known as dimensional analysis or the factor label method.


                        A simple three step process:


                        Step 1:

Read the problem, determine the units needed in the answer

Step 2:

Read the problem, determine which measurements given relate to the answer.

                        Step 3:

Use Unit Factors and exact equivalents to convert units through the equation to reach the desired answer units.

September 15, 2010

Quiz Ions 1-18 (9-17-10)
Unit Analysis Wkst d (9-15-10)
Period 3:
go over unit analysis problems b and c

September 22, 2010

Lab: Making Measurements (9-24)
Lab equipment quiz (9-24)
Quiz ions 18-37 (10-8)
no class

September 30, 2010

Quiz Ions 30-end (10-8)
Read and Outline 3.3 (10-1)
Separation of a mixture lab (10-8)
Period 3:
Lab Separation of a Mixture
Period 4:
Percent Error, Percent Composition, Percent Yield

To see previous class notes, click below

Archived Notes

If you happened to have lost something from class, click below

Chemistry Scanned Images