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Chemistry

Chemistry Periods 1 and 2

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Chemistry Periods 1 and 2
Chemistry Periods 3 and 4
Chemistry Periods 5 and 6
Chemistry Periods 7 and 8

August 30, 2010

Homework:
 
Cover books (9-3-10)
Rules and Regulations Signed and Returned (9-3-10)
Lab Safety Contract (9-3-10)
Quiz Ions 1-30 (9-17-10)
Student Info Sheets (9-3-10)
Text book scavenger hunt (9-2-10)
Lab equipment sketches (9-3-10)
 
 
Classwork:
 
Introduction to Chemistry/ rules and regulations

August 31, 2010

Homework:
 
Cover books (9-3-10)
Rules and Regulations Signed and Returned (9-3-10)
Lab Safety Contract (9-3-10)
Quiz Ions 1-30 (9-17-10)
Student Info Sheets (9-3-10)
Text book scavenger hunt (9-2-10)
Lab equipment sketches (9-3-10)
Read/ Outline/ Quiz 2.1 (9-3-10) 
 
 
Classwork:
 
Period 1:
 
Finish first day information
 
Lab: lab equipment identification/ sketches
 
Period 2:
 
Begin math pretest (to be completed and returned Thursday)
 
Begin text book scavenger hunt

September 1, 2010

Homework:
 
Cover books (9-3-10)
Rules and Regulations Signed and Returned (9-3-10)
Lab Safety Contract (9-3-10)
Quiz Ions 1-30 (9-17-10)
Student Info Sheets (9-3-10)
Text book scavenger hunt (9-2-10)
Lab equipment sketches (9-3-10)
Read/ Outline/ Quiz 2.1 (9-3-10) 
 
 
Classwork:
 
Period 1:
 
no class
 
Period 2:
 
no class

September 7, 2010

Homework:
 
Quiz Ions 1-30 (9-17-10)
Safety Skits (9-13-10)
2.3 and 2.4 read and outline/ possible quiz (9-10-10)
Lab Equipment Quiz (9-21-10)
Understanding equivalences odds (9-10-10)
 
Classwork:
 
Period 1:
 

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

mass

 

kilogram, kg

length

 

meter, m

time

 

second, s

count, quantity

 

mole, mol

temperature

 

kelvin, K

electric current

 

ampere, A

luminous intensity

 

candela, cd

 

 

 

Derived Units Commonly Used in Chemistry

 

 

 

Physical Quantity

 

Unit Name and Symbol

area

 

square meter

volume

 

cubic meter

force

 

newton, N

pressure

 

pascal, Pa

energy

 

joule, J

power

 

watt, W

voltage

 

volt, V

frequency

 

hertz, Hz

electric charge

 

coulomb, C

 

The International System of Units (SI)

-         Seven base units (shown above)

 

Definitions:

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 2:
 

Significant Digits

 

-         When combining numbers in calculations, the uncertainty of the measurements must be considered in the final answer- this is done by keeping tract of significant figures

 

-         Significant digits: The certain digits and the uncertain digit make up the significant digits

 

Non- Significant Zeros:

-         As a place keeper- not significant:

o       If the zero is not a known digit, or an estimated digit, the number is not significant

-         Zeros to the Right of a number without a decimal point:

1040 – 3 significant figures

1040. – 4 significant figures

104 – 3 significant figures

1000 – 1 significant figure

1000. – 4 significant figures

1000 – 2 significant figures (the line tells that that zero is the uncertain digit.

-         The Atlantic Pacific rule???

 

More Information on Zeros:

 

Zeros to the right of the decimal place before a whole number are not significant.  .001 has only one significant figure. .0000001 also only has one significant figure.

 

Zeros between two whole numbers are significant.  1.01, now has 3 significant figures.  .0101 also has 3 significant figures.

 

Numbers following a whole number:

If a decimal place is not present, the zeros to the right of the number are not significant unless marked with a bar.  So, 1000 has only one significant figure.

 

If a decimal place is present, the zeros to the right of the number are significant.  So 1000. has 4 significant figures.

 

Zeros following a whole number after a decimal are significant. So, in the number 14.000 the three zeros are significant so there are 5 significant figures.  .0002012300 has 7 significant figures.  Remember, those first three zeros are not significant, but the ones following the three are significant.

 

Significant Digits in Calculations:

 

            Rule for exact numbers

           

When an exact number appears in a calculation, it does not effect the significant digits of the calculations

           

                        **All Defined Value Unit Factors are Exact Numbers**

 

            Rule for addition and subtraction of measurements:

 

Find the measurement with the most uncertainty (fewest number of decimal places) and round the answer to that uncertainty.

 

Ex.  4.5 g + 3.221 g + 4.3232 g = ?

 

            4.5 has an uncertainty of .1

            3.221 has an uncertainty of .001

            4.3232 has an uncertainty of .0001

 

So, the number with the most uncertainty is 4.5 with an uncertainty of .1, therefore, the answer must be rounded off to that uncertainty (in this case, one decimal place)

 

The Math: 

12.0442 is the answer so far, but remember to correct for uncertainty the answer may only have one decimal place.  In this case, 0 should be the last significant figure.  Look to the right of zero, notice that it is a four.  Since it is a four, we need to round that four down.  The zero does not change, so the answer to the question would become 12.0 +/- .1 g.

 

            Rule for multiplication or division of measurements:

                       

Find the measurement with the fewest number of significant figures and round the answer to that number of significant figures.

                        Ex. 5.32 g x .01 g

                                   

5.32 has 3 significant figures and .01 has 1 significant figure, so the answer to the problem must have only one significant figure.

 

The answer is .0532.  Remember, now you need to correct for significant figures.  Since the answer can only have one significant figure, anything after the 5 must be rounded off.  Since 3 is less than five, it is rounded down, and the five does not change.  The answer to this problem then is .05 (one significant figure).

 

Scientific notation:

 

Scientific measurements usually appear in the form of scientific notation, written in the general form as:

 

t.tt x 10n  where t is any integer and n is the some whole number of times that 10 has been multiplied to reach a given answer.

 

Rule # 1

Find the first non-zero significant figure in a number and place the decimal place immediately to the right of that number.

                        Rule # 2

IF the decimal was moved to the left, then the power of 10 is a positive number

 

IF the decimal was moved to the right, then the power of 10 is a negative number

 

Remove all non-significant numbers (place holding zeros)

 

Example:

 

Write .00923 in scientific notation:

 

Ok, since the 2 zeros are place holders, they will not appear in the final answer.

 

The fist non-zero sig fig is the 9, so place the decimal after the 9

            9.23

Now, since the decimal was moved to the right 3 places, the power of ten is a – number

 

                                    9.23 x 10-3

            Write 5600 in scientific notation:

 

Ok, the first sig fig is the 5, so place the decimal immediately to the right of the 5

                                                            5.600

           

Since the two zeros are not significant, they should not appear in the final answer:

 

            5.6

 

Since the decimal was moved to the left three place, the power of 10 is a positive number:

 

5.6 x 103 (to check to see if you have done this correctly simply revert to the ordinary number)

 

Since 103 is 10 x 10 x 10, 5.6 is being multiplied by 1000.  when 5.6 is multiplied by 1000, the answer becomes 5600

September 13, 2010

Homework:
 
Quiz Ions 1-30 (9-17-10)
Safety Skits (9-14-10)
Lab Equipment Quiz (9-21-10)
Density Worksheet 1 (9-14-10)
Unit Analysis Review Worksheet (9-16-10)
Test Chapter 1/2 (9-20-10)
 
Classwork:
 
Period 1:
 

 

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

 

                        Ex:

                                    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

 

Elements:

-         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

 

Compounds

-         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 22, 2010

Homework:
 
Quiz Ions 30-end (10-8)
Read and outline 3-1 (9-24)
Lab: Making measurements (9-28)
Quiz Lab Equipment (9-23)
 
Classwork:
 
Period 1:
 
Lab Making measurements- data table typed with calculations
 
Period 2:
 
Computer lab and go over chapter 1/2 test

September 30, 2010

Homework:
 
Quiz Ions 30-end (10-8)
Read and Outline 3.3 (10-1)
Separation of a mixture lab (10-8)
 
 
Classwork:
 
Period 1:
 
Lab Separation of a Mixture
 
Period 2:
 
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