Archived Notes
Week of September 14, 2009
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Chemistry Period 5 Archived Notes

 
September 14, 2009

 

Homework:


Ion Quiz 1-18 (9-17)
Element Icosohedron (10-13)
Quiz Lab Safety (9-15)
Quiz Lab Equipment (9-16)

 
Classwork:
 
Period 5: 
 
lab safety video
 
Period 6:
 
no class

September 15, 16,17, 2009

 

Homework:


Ion Quiz 1-18 (9-17)
Element Icosohedron (10-13)

Test Chapter 1/2 (9-25)

Density Worksheet 1 and 2 (9-18)

 
Classwork:
 
Period 5: 
 
 
 

The Percent Concept

 

To explain the concept of percent

To apply percent as a unit factor

 

Percents express the amount of a single quantity (a portion) as compared to the entire sample. A percent is:

           

A ratio of a single quantity, to a total sample, all multiplied by 100.

 

            A dime is 10% of a dollar.

            A gram is .001% of a kilogram

 

Percent Unit Factors:

            A percent can be listed as parts per 100. 

 

            So a dime can be written as 10% of a dollar or 10/100 of a dollar.

 

            Here’s an example:

Apollo 11 traveled 394,000 km from Earth to land on the Moon.  How many Km did it travel when it was 15.5% of the distance?

 

By definition: 15.5% is 15.5 parts per 100 or 15.5/100.  Since we are dealing with Km, then it would be 15.5km/100km.  This has now become a unit factor to solve the problem.

 Mathematics lesson- percent coposition.  Stop by to see examples of percent composition and percentages as a unit factor

 


 

When Reading a Piece of Scientific Equipment:

-         Be sure to note the uncertainty of the measurement

-         Be sure to give all significant figures in the measurement.

 

Examples:  Take a look at the following pieces of equipment:

 

A Centigram Balance:

 

So, if you look at the balance, you can see that the number of certain digits include:

 

o       The hundreds, the tens, the ones, the tenths, and the hundredths.

o       In other words, you can read for certain 128.19

§         That means 5 certain digits

o       Since all measurements need one unknown digit, the last thousandths digit can be estimated

§         Therefore there would be 6 significant digits in this measurement

§         I would estimate the thousandths place to be a 2

§         Ergo, the total measurement would be:

·        128.192 g

§         Do not forget to list the uncertainty:

·        128.192 +/- .001 g

 

Try This One:

 

 

-         I got: 3.5 +/- .1 C

-         Although there are more lines present, the divisions are not certain enough to allow to you estimate the hundredths place.

-         You could choose a higher uncertainty value here: +/- .5

 

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.]

 

 
Period 6:
 

 

-         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

 

Reliability in Measurement

 

-         Precision- several measurements that are close in value

-         Accuracy- how close a measurement is to the accepted value (a standard)

 

Reliable measurements have both high precision and high accuracy.

Can you have one without the other???