December 15, 2009

Today, we asked Dr. B questions about everything that we covered this semester. We discussed Oxidation numbers, VSEPR Theory, London dispersion forces, and some questions on the test we had returned to us yesterday. Remember that the exam will cover each chapter evenly so don't study a certain chapter more than others. Leave comments about what else we discussed in class.

Charlie Fisher's Post

I believe that your points should not be based on how many cards that are used, but the amount of atoms used overall in the compound. I think that the one card was not necessary as their as an understood one after every element/polyatomic ion anyway. I also believe that all the candy that is brought for the group should be put into one pile and then we should use the bingo pieces as chips. This eliminates the problem of someone using a Reese's Pieces as an ante while someone else uses a kit kat. I don't agree with Michael in that we should have to name the compound in order to receive points, because many of the compounds that I created worked according to charge, but they were too complicated to name.

December 11

Percent Composition- the percent of an element found in a compound

• The percent of an element in a compound can be found using:

(Mass of the element in the sample of the compound/Mass of the sample of the compound)* 100

• The mass percent of an element is the same in a compound regardless of the size of the sample
• This can also be used for the molar mass

Mr. Scully TAKE 4

So It's the morning, and Charlie still hasn't posted. I personally won't hold it against him, seeing how his account is messed up. But I need these blog points. So I would like to say that the "1" card is a bit meaningless. In a revised edition of the game, we wouldn't need it. Also, after playing your hand, I think the player should have to name the elements in the compound in order to receive points.

Mr. Scully TAKE 3

• The Formula Poker Project was assigned today
• IMPORTANT: When faced with the question: "paper or plastic?" Answer: "Neither!" Because both are considered food, and Dr. B will nail you with a fine, whether swallowed or not.
• If wondering who's in your group, just picture yourself in Chemistry Class (scary, I know.) Who's sitting near you in the vision? Chances are, they're in your group.
• Person #1 chose how you're going about making the cards. RESPECT HIS AUTHORITY!
• Stop talking during class. 1st of all, no one cares what you have to say. 2nd of all, Dr. B gets mad and takes out her anger on certain people.
• Also, You're welcome for doing these three blogs in a row. I know you all hate posting, so I guess consider it an early Christmas Present from me.
• And another Present: NO QUIZ TOMORROW! WHOOP WHOOP! But quiz friday... AWWWWWWW.
  
• When Looking for how to get the Project sheet, Just remember 3 steps:

1. HER WEBSITE
2. HONORS CHEMISTRY
3. FORMULA PAPER

It's as easy as 123456 Pokemon. A video which you should all check out on youtube ON YOUR OWN TIME. Because right now, it's time to focus on chemistry.
POKER CHIPS=CANDY! BRING CANDY! PLEASE PLEASE PLEASE!

Mr. Scully TAKE 2

• Hey guys. I'll keep the shenanigans to a minimum today, due to request by Dr. B

• One thing to remember: If doing multiple parts to one question, don't round your first answer when applying it to the second part of the problem. YOU COULD MAKE A MISTAKE. Not that I expect that to happen. You are all very bright students: Hence the honors chemistry.

• OH! Don't forget: Finish the packet, oh ye who hath suffered lack of knowledge of its due...ness.

• Also for homework: PAGES 252-253. NUMBAS 28,29,30,31,32,and 33

• Ok, back to the information: It is useful to know the percentage by mass of a particular element in a chemical compound.

• Why? I don't know. But here's how you find it=

• % of element = 100 * (Mass of that particular element / mass of the whole compound)

• Exams are coming up, and I know studying sounds lame... but it's important.

• Yes, Michael Scully is actually telling someone to study. Shocking, I know.

• But Even Though You Want to Give up, You must KEEP ON!

• I know what you're thinking. You think that the quarter is over.
  

Over? Did you say "over"? Nothing is over until we decide it is! Was it over when the Germans bombed Pearl Harbor? Heck no! And it ain't over now. 'Cause when the goin' gets tough... the tough get goin'! Who's with me? Let's go! Ace these quizzes, Master The Material, Go Out There, And Make Yourself Proud! You've worked too hard to let the end of the quarter kill your grade. We can do this, guys. We're smart. That's why we were chosen for the Honors Track. Because the work's harder, but we can handle it! Be smart. Be strong.

Mr. Scully TAKE 1

-The molar mass of a compound can be used as a conversion factor. We said this about twenty times at the end of class today, so it must be important.
WHAT IN THE WORLD IS THE MASS OF 2.50 mol of OXYGEN GAS?!?
Don't fret my friend. It's relatively simple. 1st, realize that Oxygen exists in its gaseous state as TWO oxygen bonded to each other.
SHOCKING, I KNOW.
Buy anywho... it's just simple multiplication from here. Oxygen (or oxy, as her friends call her) has a molar mass of 16.00, roughly. Since oxy and her twin sister are hanging out together, that means we have to take 16.00 and multiply it by TWO. Now multiply that by your number of mols, and you have your answer.

so, basically just 2.50*2*16.00

BUT WAIT! HOW MANY SIGNIFICANT DIGITS DO I GO TO?!?
That's enough fretting out of you, bud. When dealing with this problem, just understand that the 2.50 mols is your estimated amount. How many sig-digs does it have? 3. How many sig-digs does your answer have? 3.

=80.0 grams. Bam, Boom, See you later.

Now as most of you remember, Dr. B flashed something about Ibuprofen across the screen for about 2 seconds at the end of class
Perhaps an exaggeration, but I reserve my creative license.
Thankfully, she pointed out to me that the problem's in the book.
<3>And seeing how you're probably already bored with my ridiculous attempts to make 5 minutes of notes in chemistry entertaining, I'll just spell the problem out for you really simply.

Ibuprofen, C13H18O2 is the active ingredient in many nonprescription pain relievers. Its molar mass is 206.31 g/mol
a. If the tablets in a bottle contain a total of 33 g of ibuprofen, how many moles of ibuprofen are in the bottle?
b. How many molecule of ibuprofen are in the bottle?
c. What is the total mass in grams of carbon in 33g of ibuprofen?
Given: 33 g of C13H18O2, molar mass 206.31 g/mol
Unknown: a. moles C13H18O2
b. Molecules C13H18O2
C. total mass of C

Solution
a. 33g * (1 mol/206.31) = 0.16 mol (don't forget sig-digs)
b. 0.16 mol * (6.022*10^23) = 9.6*10^22
c. 0.16 mol * (13 mol) * 12.01 g = 25 g

Friday, December 4th 2009

Well not much to blog about. We played bingo and we went over the ion quiz. There will be an ion quiz Monday. Also the fall exam review as been posted on Dr. B's website. It is a very helpful study guide.

December 3, 2009

Today we played bingo, took the quiz over tables 1 and 2, and went over the test.
Begin studying for the next ion quiz
Homework: Next page in the Chapter 7 Review

December 2nd, 2009

Quiz tomorrow (Tables 1 & 2)

Formula Masses

• The mass of a water molecule can be reffered to as a molecular mass.
• The mass of one formula unit of an ionic compound such as NaCl is not a molecular mass.
• The mass of any unit represented by a chemical formula can be reffered as the formula mass.

Molar Mass

• The molar mass of a substance is equal to the mass in grams of one mole of the substance.
• The molas mass of a compound is calculated by adding the masses of the element's present in a mole of the molecules or formula units that make up the compound.
• One mole of H2O molecules contains exactly two moles of H and one mole of O atoms.
• A compound's molar mass is numerically equal to its formula mass.

Molar Mass as a Conversion Factor

• The molar mass of a compound can be used as a conversion factor to relate an amount in moles to a mass in grams for a given substance.
• To convert moles to grams, multiply the amount in moles by the molar mass.

Tuesday December 1, 2009

Tuesday December 1, 2009
For the first part of class today, we played bingo. Then we took a couple of notes:

We first did the example problem on the book on page 233-234. You can look for yourself for the work and the answers. The text does a good job at explaining it so you might wanna take a look at that.

USING OXIDATION NUMBERS FOR FORMULA AND NAMES
- Many nonmetals can have more than one oxidation number.
- These numbers can sometimes be used in the same manner as ionic charges to determine formulas.
- Suppose, for example, you want to know the formula of a binary compound formed between sulfur and oxygen. From the Common +4 and +6 oxidation states of sulfur, you could expect the sulfur might form SO2 or SO3. Both are know compounds.
- Using oxidation numbers, the Stock System, introduced in the previous section for naming ionic compounds, can be used as an alternative to the prefix system fir naming binary molecular compounds.
- An example of this is shown on page 235.

SECTION 3
- A chemical formula indicates:
- The elements present in the compound
- The relative number of atoms or ions of each element present in a compound
- Chemical formulas also allows chemists to calculate a number of characteristic values for a given compound.
FORMULA MASS
- The formula mass of any molecule, formula unit, or ion is the sum of the average atomic masses of all atoms represented in its formula.


That was it for today. Don't forget about the quiz tomorrow on table 1 and 2 and also some more info on the Chapter.
See ya

Monday, November 30th, 2009

Covalent-Network Compounds

• Some covalent compounds do not consist of individual molecules.
• Instead, Each atom is joined to all its neighbors in a covalently bonded, three-dimensional network
• The subscripts in a formula for a covalent-network compound indicate the smallest whole-number ratio of the atoms in the compound.

Ex: SiC-Silicon carbide

SiO2-Silicon dioxide

Si3N4-Trisilicon tetranitride

Acids and Salts

• Acid is a certain type of molecular compound.
• Most acids used in the laboratory can be classified as either binary acids or oxyacids.
• Binary Acids- Are acids that consist of two elements, usually hydrogen and one of the halogens-Fluorine, chlorine, bromine, iodine.
• Oxyacids- Are acids that contain hydrogen, oxygen and a third element (usually a nonmetal)
• In chemical nomenclature, the term acid usually refers to a solution in water of one of these special compounds rather than to the compound itself.
• For example, hydrochloric acid refers to a water solution of the molecular compound hydrogen chloride, HCl.
• Many polyatomic ions are produced by the loss of hydrogen ions from oxyacids.

Ex. Sulfuric acid- H2SO4 Sulfate-SO2-/4

Nitric acid- HNO3 Nitrate-NO-/3

Phosphoric acid- H3PO4 Phosphate- PO3-/4

• An ionic compound composed of a cation and the anion from an acid is often referred to as a salt. Table salt, NaCl, contains the anion from hydrochloric acid. Calcium Sulfate, CaSO4, is a salt containing an anion from sulfuric acid. HCO-/3, Hydrogen carbonate ion bicarbonate ion.

Oxidation Numbers

• The charges on the ions composing an ionic compound reflect the electron distribution of the compound
• In order to indicate the general distribution of elections among the bonded atoms in a molecular compound or a polyatomic ion, oxidation numbers, also called oxidation states, are assigned to the atoms composing the compound or ion.
• Unlike ionic charges, oxidation numbers do not have an exact physical meaning.

Assigning Oxidation Numbers

• As a general rule in assigning oxidation numbers, shared electrons are assumed to belong to the more electronegative atom in each bond.
• More specific rules for determining oxidation numbers are provided by the following guidelines:

1. The atoms in a pure element have an oxidation number of zero. For example, the atoms in pure sodium, Na, oxygen, O2, phosphorus, P4, and sulfur, S8 all have oxidation numbers of zero.
2. The more-electronegative element in a binary molecular compound is assigned the number equal to the negative charge it would have as an anion. The less-electronegative atom is assigned the number equal to the positive charge it would have as a cation.
3. Fluorine has an oxidation number of -1 in all of its compounds because it is the most electronegative element.
4. Oxygen has an oxidation number of -2 in almost all compounds. Exceptions include when it is in peroxides, such as H2O2 , in which its oxidation number is -1, and when it is in compounds with fluorine, such as OF2, in which its oxidation number is +2.
5. Hydrogen has an oxidation number of +1 in all compounds containing elements that are more electronegative than it; it has an oxidation number of -1 in compounds with metals.
6. The algebraic sum of the oxidation numbers of all atoms in a neutral compound is equal to zero.
7. The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion.
8. Although rules 1 through 7 apply to covalently bonded atoms, oxidation numbers can also be assigned to atoms in ionic compounds.

11/24/09

Hw assignment Monday night- fill out the bingo cards Dr. B is giving us, fill it in with the ions of charts 1 and 2 from chapter 7. We play bingo on Tuesday- winners get candy!!!!

Naming Binary Ionic Compounds

The Stock System of Nomenclature

• Some elements such as iron, form two or more cations with different charges
• To distinguish the ions formed by such elements, scientists use the stock system of nomenclature.
• The system uses a Roman numeral to indicate an ion's charge.

Compounds Containing Polyatomic Ions

• Many common polyatomic ions are oxyanions polyatomic ions that contain oxygen.
• Some elements can combine with oxygen to form more that one type of oxyanion
• Ex: nitrogen can form WO-3 or NO-2
• The name of the ion with the greater number of oxygen atoms end in -ate. The name of the ion with the smaller number of oxygen atons ends in -ite.

NO-3 NO-2

nitrate nitrite

• Some elements can form more than two types of oxyanions.
• Ex: Chlorine can form ClO-, ClO-2, ClO-3, ClO-4
• In this case, an anion that has one fewer oxygen atom than the -ite anion has is given the prefix -hypo.
• An anion that has one more oxygen atom thatn the -ate anion has is given the prefix per-.

ClO- ClO-2 Cl0-3 ClO-4

hypochlorite chlorite chlorate perchlorate

DO NOT SAY DIMERCURY, CALL IT MERCURY 1 OR 2.

Naming Compounds with Polyatomic Ions

• Name the cation
• Name the anion
• Name the salt- names of cation and anion

YOU WANT POLYATOMIC IONIC COMPOUNDS TO BE ELECTRICALLY NEUTRAL

YOU CANNOT MOVE ATOMS FROM 1 POLYATOMIC ION TO THE NEXT

Naming Binary Molecular Compounds

• Unlike ionic compounds, molecular compounds are composed of individual covalently bonded units, or molecules.
• As with ionic compounds, there is also a stock system for naming molecular compounds.
• The old system of naming molecular compounds is based on the use of prefixes.
• Ex: CCl4- carbon tetrachloride (tetra=4)
• Ex: CO-carbon monoxide( mon=1)
• Ex: CO2- carbon dioxide ( di=2)

Covalent -Network Compounds

• Same covalent............ (TO BE CONTINUED!!)

Movie

So I know Dr. B wasn't at school today, but this movie really interested me. I think it's cool how Orpheus crashed into the Earth back in the day. I never realized that the way planets collide cause such different outcomes. Where do you think is the next place humans will travel after Mars? Will we ever find another easily inhabitable planet? Anyone else have something to add?

11/22/09

Hey guys, sorry it took so long to get the blog up.

CH 7.
Significance of a Chemical Formula.

• a chemical formula indicates the relative numbers of each kind of atoms in a chemical compound.
• for a molecular compound, the chemical formula reveals the number of each element conrained in a single molecule of a compound.
• ex. octane - C8H18 (subscripts after element symbols express number of each atom in the molecule.
• the chemical formula for an ionic compound represents one formula unit - the simplest ratio of the compounds positive ions (cations) and its negative ions (anions.)
• ex. Al2(SO4)3
• ( ) around the polyatomic ion identifies it as the unit.
• note. no subscript = UNDERSTOOD TO BE 1.
• Monatomic Ions
• many main-group elements can gain or lose e- to form ions.
• ions formed from a single atom are known as monatomic ions.
• some main-group elements tend to form covalent bonds eather than form ions.
• Naming Monatomic Ions
• monatomic cations are identified simply by elements' name.
• ex. K+ is called potassium cation.
• Mg 2+ is called magnesium cation.
• for monatomic anions, ending of elements' name is dropped, and the ending "-ide" is added to the root name.
• ex. F- is called FLOURIDE anion.
• N3- is called the NITRIDE anion.
• STUDY THE CHAPTER 7 CHARTS GUYS.
• Binary Ionic Compounds
• compunds composed of 2 elements are known as binary compounds.
• in a binary ionic compound, the total number of positive charges and negative charges must be equal.
• the formula for a binary ionic compound can be written given the identities of the compound's ion.
• ex. magnesium bromide
• ions: Mg2+, Br-, Br-
• = MgBr2
• Al3+O2-
• Al2O3
• "CROSS OVER METHOD" = bleh.
• Naming Binary Ionic Compounds
• the nomenclature, or naming system, of binary ionic compounds involves combining the names of the compound's positive and negative ions.
• the name of the cation is given 1st; followed by name of the anion.
• ex. Al2O3 - aluminum oxide.

Post/Comment Tally

I have counted your posts for the quarter and I will let you know today as you turn in your test what your current count is. There are 15 days left in this quarter. That does NOT include today since we are having a test. Every other day after today is counted. To date, 2 people have posted or commented 20 or more times. 5 of you are in single digits. Everyone has at least one comment/post. One person can not get 20 by the end of the quarter and one person will have to comment every SINGLE day to get to 20. Comments on this post will NOT count for your total!

Direction of Dipoles is represented by an arrow toward negative poles and a cross at the positive pole. the positive dipole is indicated as follows

-I-->

H Cl

the negative region in one polar molecule attracts the positive region in adjacent molecules. the moleules are attracted to each other by opposite sides.

each forces of attraction between polar moleules are known as a dipole- dipole forces

dipole-dipole foreces act at short range only between nearyby molecules.

this explains the high boiling point in molcules with dipole dipole forcesex I-Cl (97 c) but Br-Br (59 c)

a polar molecule can induce a dipole in a non polar molecule by temporarily attracting its electrons

the result is a short range intermolecular force that is weaker than a dipole-dipole force.

this accounts for the oxygen's ability to be dissolved into water

some hydrogen containg compounds have high boiling points. explored by a particularly strong of dipole-dipole force

examples are phosphorous and sulfur

this gives a hydrogen atom a positive charge that is almost half that of a bare proton

the small size of the hydrogen atom allows it come close to an unshared pair in another atom

this forms a hydrogen bond

these are connected with a dotted line

excellent example is water

properties aquired from hydrogen bonding are surface tension, cohesion, solvent, and boiling point is raised

london dispersion theory

even noble gas atoms and non polar molecules can experience weak intermolceular bonding

in any atom or molecule the electrons are in constant motion

thus there exists a possibilty of there being a more positive and negative side which attracts the more negative or positive sides of a different atom or molecule

this is the weakest type of bond

suggested in 1930

11/17/09

HYBRIDIZATION(HB): the blending of bonding orbitals

• 2 types of bonding: Sigma(o-) and Pi(~)(I know these aren't the actual signs but i do not know how to make them on a computer so I will just use these)
• Sigma is ALWAYS the first bond
• Pi will be the second bond or any after that
• Single bond: 1 o-
• Double bond: 1 o-, 1 ~
• Triple bond: 1 o-, 2~

IN HYBRIDIZATION YOU ONLY PAY ATTENTION TO THE SIGMA BONDS AND LONE PAIRS(LP)!!!!!!!!!!

• Example: Methane, CH4

(Hydrogen does not form hybrids)

Because HB is o- bonds plus LPs, to hybrid Carbon, first find out the number of o- bonds, ~ bonds and LPs. (You may have to draw the Lewis structure for this). In C there are 4 o- bonds and no LPs. Therefore the HB=4=sp^3.

• If HB results in an sp orbital, then the molecular geometry is 180 degrees(linear)
• If HB results in an sp^2 orbital, then the molecular geometry is 120 degrees(trigonal-planar)
• If HB results in an sp^3 orbital, then the molecular geometry is 109.5 degrees(tetrahedral)

Intermolecular Forces (IMF)

• The forces of attraction between molecules are known as intermolecular forces
• The boiling point of a liquid is a good measure of the IMF but its molecules: the higher the boiling point, the stronger the forces between the molecules
• IMF varies in strength but are generally weaker between atoms within molecules, ions in ionic compounds, or metal atoms in solid metals
• Boiling points for ionic compoundsand metals to be much higher than those for molecular substances: forces between molecules are weaker than those between metal atoms or ions
• The strongest IMF is between polar molecules
• Because of their uneven charge distribution, polar molecules have dipoles. A dipole is created by equal but opposite charges that are separated by a short distance
• The direction of a dipole is from the dipole's positive pole to it's negative pole

TEST IS THURSDAY! MAKE SURE YOU STUDY! READ THE CHAPTER

HOMEWORK: PAGE 211-212, #'s 50-60. IT IS HOMEWORK 6

(If you couldn't tell, I decided to go Memphis colors for the game tonight because Memphis is the best team in the nation. Hope you can read the gray good enough. I don't think it should be a problem. GO TIGERS!!!!)

PS Ben Smith

(\,,(\

(=' ;')

(,(")(")

Woops

Hey, i left my notes at school so....

yeah, sorry class, but now is your opportunity to post class notes since i am inadequate

Friday, the 13th of November 2009 AD

[A quiz over 6.4 is scheduled for Tuesday, and the Ch. 6 test is scheduled for Thursday, along with the outline. We're supposed to do the aluminum lab Tuesday or Wednesday.]

Molecular Geometry

• properties of molecules depend not only on bonding, but also molecular geometry
• molecular geometry-3D arrangement of a molecule's atoms
• polarity of each bond + molecular geometry ---> molecular polarity
• --> the uneven disrtibution of electron density

1) strongly influences the forces that act between molecules in liquids and solids

• chemical formula reveals little about the molecular geometry, so we need something... better... Like...

VSEPR THEORY!!!

• diatomic molecules can only be linear because there is only 2 atoms
• use VSEPR to predict molecular geometry of complicated molecules

1) it takes into account the location of all the electron pairs surrounding the bonding of atoms

• Valence Shell Electron-Pair Repulsion
• VSEPR Theory- repulsion between sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible (the electrons want to have the greatest distance they can have between them)

ex. - BeF2 {only 2 electron pairs shared between the atoms, so they will be 180 degrees apart, making the molecule's geometry linear}

If A--central atom in a molecule
B--atoms bonded to A
Then according to VSEPR, BeF2 is an AB2 molecule (i.e. linear)

• AB3-3 A-B bonds stay farthest apart by pointing to the corners of an equilateral triangle {120 degree bonds}
• AB4- distance between electron pairs maxed if the A-B bonds point to the corners of a tetrahedron {109.5 degree bonds}

[past AB4, there are exceptions to the octet rule---the nonmetals 3rd period and below

• VSEPR also acounts for the unshared electron pairs

-->like ammonia (NH3) and water (H2O)

• Lewis structure of ammonia shows central atom has an unshared electron pair
• VSEPR postulates that the lone pair occupies space around the N atom just as bonding pairs do
• NH3 = AB4 molecule, but with a pyramid shape instead of a tetrahedron

KEY POINT--SHAPE OF MOLECULE CONCERNS THE POSITION OF ATOMS ONLY

• H2O with 2 unshared pairs has a "bent," or angular, geometry
• bonds and electrons take up different amounts of space (electrons take up more than bonds)
• unshared electron pairs repel other electron pairs more strongly than bonding pairs, which pushes the bonded atoms together
• same basic principles of VSEPR that have been described can be used to determine the geometry of several additional types of molecules -- AB2E, AB2E2, AB5, AB6

1) treat double and triple bonds the same as single bonds

2) treat polyatomic ions similar to molecules

And that is where we stopped.

Study for the quiz with your book, because we took only 2 lines of notes over sect. 4

November 12, 2009

We had a quiz today over Section 2 of Chapter 6. Then we took a couple notes:

• In a metal, the vacant orbitals in the atom's outer energy levels overlap.
• This overlapping of orbitals allows the outer electrons of the atoms to roam freely through the entire metal.
• The electrons are Delocalized, which means that they don't belong to any one atom, but move freely about the metal's network of empty atomic orbitals.
• they form a sea of electrons around the metal atoms, which are packed together in a crystal lattice.
• Metallic Bonding - results from attraction between metal atoms and the surrounding sea

Dr. B reminded us that we should be working on our outlines and reading along and ahead in the textbook.

Have a good night!

Wednesday, November 11, 2009

A Comparison of Ionic and Molecular Compounds

• force that holds ions together in ionic compounds is a very strong overall attraction between positive and negative charges
• in molecular compound, the covalent bonds of the atoms making up each molecule are also strong
• the forces of attraction between molecules are much weaker than the forces among formula units in ionic bonding which gives rise to different properties in the two types of compounds
• for ionic compounds--- high melting and boiling points, hard but brittle
• ionic compounds are hard but brittle because in an ionic crystal even a slight shift of one row of ions relative to another causes a large buildup of repulsive forces that make it difficult for one layer to move relative to another layer; but if one layer is moved the repulsive forces make the layers part completely causing ionic compounds to be brittle
• in the solid state ions cannot move so the compounds are not eletrical conductors
• in the molten state ionic compounds are eletrical conductors because the ions can move freely to carry eletrical current
• many ionic compounds are soluble in water
• when they are dissolved in water they do conduct eletrical current
• other ionic compounds do not dissolve in water because the attractions between water molecules and the ions cannot overcome the attractions between the ions
• polyatomic ion= a charged group of covalently bonded atoms
• polyatomic ions have characteristics of both molecular and ionic compounds
• the charge of a polyatomic ion results from either an excess or shortage of electrons
• use formal charge with polyatomic ions to figure out the most stable lewis structure
• in a polyatomic ion's lewis structure, often times brackets are drawn around the entire structure and the charge of the polyatomic ion is written on the outside of the brackets

Metallic Bonding

• chemical bonding is diferent in metals than it is in ionic, molecular, or covalent-network compounds
• thus they have special properties
• excellent eletrical conductors in solid state, strong reflectors and absorbers of light, luster
• metallic bonding= the chemical bonding that results form the attraction between maetal atoms and the surrounding sea of electrons
• sea of electrons----> in a metal vacant orbitals in atoms' outer energy levels overlap and this allows the outer electrons of the atoms to roam freely throughout the entire metal; the electrons are delocalized which means that they don't belong to any one atom but move freely about the metal's network of empty atomic orbitals; these mobile electrons form a sea of electrons which are packed in a crystal lattice
• malleability= the ability of a substance to be hammered or beaten into thin sheets
• ductility= the ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire
• ductility of metal is allowed because metallic bonding is the same in all directions throughout the solid which allows, when struck, one plane of atoms to slide past another without encountering any resistance

quiz tomorrow---> be prepared and read ahead

Dr. B might be posting on this blog tonight

Finally, do not forget to begin studying the tables that are in the next chapter so you won't have to wait till the night before to learn them all

November 10

Today, we went over homework # 4. It was filled with symbols with Lewis structures and whatnot and I don't really know how to make those on this blog, so if you weren't here to get the correct answers, see someone who was here.

Formal Change
Formal Change = Valence Electrons - (Bonding Electrons/2) - Lone Pair Electrons
F.C. = V.E. - (B.E./2) - L.P.E.
In Lewis Structures, you want to create the structure with the lowest Formal Charge and have the central atom having a positive formal charge and the surrounding atoms have a negative formal charge.
A Comparison of Ionic and Molecular Compounds
-The force which holds ions together in ionic compounds is a very strong electrostatic attraction
-In contrast, the forces of attraction between molecules of covalent compounds are much weaker
-This difference in strength of attraction between basic units of molecular and ionic compounds gives rise to different properties between the two types f compounds.

11/9/09

Today, we went over the test and took some notes. The outline for ch. 6 is due sometime at the end of the week. We will also have a QUIZ on ch.6 on THURSDAY.

Ionic Compounds

Most of the rocks and minerals that make up the Earth's crust consist of positive and negative ions held together by ionic bonding. (Ex. Table salt consists of sodium and chloride ions combined in a one to one ratio so that each pos. charge is balanced by a neg. charge.)

A ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal.

Most ionic compounds exist as crystalline solids. A crystal of any ionic compound is a three-dimensional network of positive and negative ions mutually attracted to each other. In contrast to a molecular compound, an ionic compound is not composed of independant, neutral units that can be isolated.

The chemical formula of an ionic compound represents not molecules, but the simplist ratio of the compound's ions. A formula unit is the simplist collection of atoms from which an ionic compound's formula can be established. The sodium ion has one valence electron and the chlorine atom has seven valence electrons.

Atoms of Sodium and other alkali metals easily loose one electron to form cations. Atoms of chlorine and other halogens easily gain one electron to form anions. In an ionic crystal, ions maximize their potential energy by combining in an orderly arrangement known as a crystal lattice. Attractive forces exist between oppositely charged ions within lattice.

November 7th Notes

Pairing Lewis Structures with Many Atoms
1) Gather information

• Draw a Lewis structure for each atom in the compound. When placing valence electrons around an atom, place one electron on each side before pairing ant electrons on each side before pairing an electrons
• Determine the total number of valence electrons in the compound

2)Arrange Atoms

• Arrange the Lewis structure to show how the atoms bond in the molecule
• Halogen and Hydrogen atoms often bond only to one other atom and are usually at the end of the molecule
• Carbon is often placed in the center of the molecule
• You will find that, with the exception of carbon, the atom with the lowest electronegative is often the central atom

3) Distribute the dots

• Distribute the dots so that each atom is satisfied (With exception of Beryllium, Boron, and hydrogen).

4)Draw bonds

• Change each pair of dots that represent a shared pair of electrons by a large dash

5) Verify Structure

• Count the number of electrons surrounding each atom. Except for hydrogen, beryllium, and boron all atoms most satisfy octet rule. Check that the numbers of valence electrons is still the same number you determined in Step 1.

Here is a site that has many of sample problems that you can practice with

http://www.colby.edu/chemistry/webmo/mointro.html

Thursday, November 5

-Lewis Structure

-can be drawn if you know composition of the molecule and how the atoms bond

-a single covalent bond, single bond, is a bond in which one pair of e- are shared b/t 2 atoms

EXAMPLE

http://www.bing.com/images/search?q=structural+formula+of+ethanol&FORM=BIFD#focal=33d84189786a894a145ce3585b7a5146&furl=http%3A%2F%2Fwww.green-planet-solar-energy.com%2Fimages%2Fethanol-ethanol.gif

***If Carbon is present, it is the central atom of the molecule. If it is not present, the lest electronegative atom is the central atom. Hydrogen is never the central atom***

-Multiple Covalent Bonds

-double bond--a bond in which two pairs of electrons are shared between two atoms

-double bonds often form in molecules containing carbon, nitrogen, and oxygen

-a double bond is written with two dashes between the two bonded atoms

EXAMPLE

http://www.indigo.com/models/gphmodel/molymod-alkene-molecule-structure.jpg

-triple bond--a bond in which three pairs of electrons are shared between two atoms

-diatomic nitrogen and ethyne are two examples of molecules with double bonds

EXAMPLE

http://lb.chemie.uni-hamburg.de/search/pic.php?170/Qx7FgJkL.png

Single covalent bonds are the longest and lowest in energy
double covalent bonds are shorter and higher in energy
triple covalent bonds are shortes and highest in energy

-Double and Triple bonds are multiple covalent bonds

-when writing Lewis structures for carbon, nitrogen, and oxygen, remember multiple bonds between these atoms is possible

11/3/09

Test Tomorrow

- 40 questions, no calculations
- Know where the groups in the periodic table are and know their names
- Be able to give examples of electron configuration

Exceptions to the Octet Rule

• Exceptions to the octet rule include those for atoms that cannot fit eight electrons, and for those that can fit more than 8 electrons, into their outermost orbital
• Hydrogen - forms bonds in which it is surrounded by only two electrons
• Boron - has just three valence electrons, so it tends to form bonds in which it is surrounded by six electrons
• Main group elements in periods three and p can form bonds with expanded valence, involving more than eight electrons

Electron-Dot Notation

• to keep track of valence electrons, it is helpful to use electron-dot notation
• electron-dot notation is an electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element's symbol. the inner shell electrons are not shown

Lewis Structures

• electron-dot configuration can also be used to represent molecules (ex. H:H)
• the pair of dots between symbols represents the shared electron pair of the Hydrogen-Hydrogen covalent bond
• for a molecule of fluorine, F2, the electron-dot notations of the two fluorine atoms are combined
• an unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom
• the pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash (ex. H-H)
• a structural formula indicates the kind, number, and arrangement, and bonds but not the unshared pairs of atoms in a molecule

Where's the blog for 11/2/09?

Creepy Covalent Bonds (please read the blog. I put a lot of work into it.)

HAPPY HALLOWEEN, EVERYONE... which calls for a special blog. hehehe

{in class: we went over the quiz and took notes. ON Tues.- Outline due and Ch. 5 test. ON Wed.- Ch. 6 quiz. START MEMORIZING THE TABLES 1,2,and 5 ON PGS. 221, 226, AND 230. Dr. B said, "if you want your grades to improve, start now!" Quizes will be taken by tables}

Formation of a Covalent Bond

• Most atoms (and zombies) have a lower Potential Energy when they are bonded with other atoms (or zombies) than when the atom/zombie is alone.
• The chart in our book on pg. 179 shows the P.E. changes during formation of H-H bond

(SHE ALWAYS ASKS QUESTIONS ABOUT THIS CHART!!!)

• Suppose you have two atoms: the electron of one atom ATTRACTS the proton of the other atom (like kids and candy)
• in these same atoms, the electrons in the different atoms REPEL each other (same with the protons) (like candy and evil dentists)

--> These forces cancel out to form a covalent bond where the P.E. is lowest (they are chained together for eternity!!!!!! or until something breaks them up...)

Characteristics of a Covalent Bond

(VOCAB ALERT!)

• bond length- the distance between two bonded atoms at the minimum P.E.

--> The average distance between two bonded atoms

• Atoms release energy forming covalent bonds (like a mummy released from the grave!)
• The same amount of energy must be added to separate the atoms (stuff that mummy back in the grave! haha Take that mummies!)

(VOCAB ALERT!)

• bond energy- the energy required to break a chemical bond and form isolated atoms

• Shared electrons of two atoms in a covalent bond form overlapping orbitals

--> atoms are jealous of noble gases, so they want to make their outer shells like a noble gas's (then they can overthrow the noble gases and rule the periodic table muahaha!)

ex.- two bonded H atoms (with the overlapping orbitals) can each have He's electron configuration

The Octet Rule

• the reason noble gases are unreactive is because their electron configuration is esp. stable

--> stability comes from full s and p orbitals

• Covalent bonding allows other atoms to reach this stability

(IMPORTANTE!!!)

• OCTET RULE- Chemical compounds tend to form sa that each atom, by gaining, losing, or sharing electrons, has an octet in its highest energy level

[THIS RULE APPLIES ONLY TO THE MAIN GROUP ELEMENTS IN THE 2ND PERIOD AND BELOW]

(ALSO IMPORTANTE!!!)

• There are exceptions to the octet rule!- any atoms that can't fit 8 electrons or can fit more than 8 electrons in its outermost energy shell
• the exceptions are- ARGGGGGGGGGGGGHHHHHHHHHHHHHHHHH! (the writer of this blog was dragged away by disgruntled zombies, an evil dentist, mummies, and jealous elements)

(we will cover what the exceptions are on Monday)

Chemical Bonding

Chapter VI: Chemical Bonding
10/29/09

NOTES:
Ionic bonding- chemical bonding that results from the electrical attraction between cations and ations

Covalent bonding- bonds that result from sharing of electron pairs between two atoms

• Nonpolar Covalent Bond- two atoms of same size bond
• Polar Covalent Bond- an atom bonds with an atom of a different size

Percentage Ionic Character-

• To find, subtract the electronegativity of two elements, then find the difference on the percentage ionic character.
• If 0-.3, then it's nonpolar covalent
• If .4-1.7, then it's polar covalent
• if 1.8-3.3, then it's ionic

Molecular Compounds

• a molecule is a neutral group of atoms that are held togethor by covalent bonds
• a chemical compound whose simplest units are molecules is a molecular compound
• the composition of compound is given by its chemical formula
• a chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.
• a molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound

Extra notes: Chapter 5 Test will be on Tuesday, November 3rd; The Chapter 5 outline will be due on the day of the test, Tuesday. Have a good afternoon.

10/28/09

The electrons in te both the cations and anions are in higher energy levels as one reads down a group.

There is a gradual increase of ionic radii down a group

Valence Elctrons

Chemical compunds for because electrons are lost, gained, orshared between atoms

The electrons that interacts in this manner are those in the highest energy level

The electrons available to be lost, gained, or shared in the formation of chemical compundsare referred to as valence elctrons

Valence electrons are often located in the incompletely filled main-enrgy levels

Example: The electron lost from the 3s sublevel of Na to form Na+ is a valence electon

Electronegativity

Valence electrons hold atoms togehter in chemical compounds

In many cmpounds, the negative charge of the valence eectron is concentrted closer to one atom than to another

Electronegativit is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound

Electronegatvities tend to increase across periods and decrease or remain about the same down a group

Need to Know on the Periodic Table

Groups 1 and 2

Inner T.M

Main Group(s&p)

Halogens(17)

Noble gases(18)

Metals(+ ions)

Nonmetals(-ions)

Metaloids

Atomic radius increase from bottom up and right to left

Ionization energy incrreases from bottom to top and left to right

Electron affinity increases bottom to top and left to right

Electronegativity increases bottom to top and left to right

ions!

(+) Cations smaller than original atom

(-) Anions larger than originl atoms

Chapter 3

Electron Affinity

• The energy change that occurs when an electron is acquired by a neutral atom is called that atom's electron affinity
• Electron affinity generally increases across periods
• Increasing nuclear charge along the same sublevel attracts electrons more strongly
• Electron affinity generally increases down groups
• The larger an atom's electron cloud is, the farther away its outer electron are from its nucleus

Ionic Radii

• A positive ion is known as a cation
• The formation of a cation by the loss of one or more electrons always leads to a decrease in atomic radius
• The electron cloud becomes smaller
• The remaining electrons are drawn closer to the nucleus by its unbalanced positive charge
• A negative ion is known as an anion
• The formation of an anion by the addition of one or more electrons always leads to an increase in ionic radius
• Cationic and anionic radii decrease across a period

10-26-09

Regarding the dart lab: It's due Thursday. Use a line graph for the graphs. Clarity is important; neatness isn't. If you had hits outside of the circle, add a row for >10cm, calculate the area of the sheet of paper (remember to convert inches to centimeters), and add your data. Hits are infinitely significant for calculations. The figure on the sheet is not the one from you book. Use Fig. 11 on page 107.

Atomic Radii

• The boundaries of an atom are fuzzy, and an atom's radius can vary under different conditions.
• To compare different atomic radii, they must be measured under specific conditions.
• Atomic radius may be defined as one-half the distance between the nuclei of identical atoms that are bonded together.
• Atoms tend to be smaller the farther to the right they are found across a period.
• The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus, which attracts electrons toward the nucleus.
• Atoms tend to be larger the farther down in a group they are found.
• The trend to larger atoms down a group is caused by the increasing size of the electron cloud around an atom as the number of electron sublevels increases.

Sample Problem E

Of the elements Mg, Cl, Na, and P, which has the largest atomic radius? Highlight for answer. Na. All of these elements are in the same period, and Na is right-most.

Ionization Energy

• An ion is an atom of group of bonded atoms that has a positive or negative charge.
• Na, for example, easily loses an electron to form Na+.
• Any process that results in the formation of an ion is referred to as ionization.
• The energy required to remove one electron from a neutral atom of an element is the ionization energy, IE (or first ionization energy, IE1).
• In general, ionization energies of the main-group elements increase across each period.
• This increase is caused by increasing nuclear charge.
• A higher level charge more strongly attracts electrons in the same energy level.
• Among the main-group elements, ionization energies generally decrease down the group.
• Electrons removed from atoms of each succeeding element in a group are in higher energy levels, farther from the nucleus.
• The electrons are removed more easily.

P.S. Since I'm such a nice guy, I'll give somebody a chance to volunteer to do the blog tomorrow. If nobody does, I'm going to pick somebody using a random number generator, so no whining if luck doesn't favor you.

10/22/09

We have a lab tommorow
-Each person in the group drops it 50 times at shoulder length with arm extended
-Record where it sticks in the target after each drop
-If it doesn't go through the paper, mark an x where it hit
-You decide whether it is in or out if it lands on a line
-Don't poke people with the darts (Ben).

Notes
-The 3d sublevel is higher in energy than the 4s sublevel, so they are filled in the order
4s3d
-D Block metals are typically have metallic properties and are often referred to as
transition elements
-P Block elements consist of all the elements of Groups 13-18 except helium.
-The p Block elements together with the s Block elements are called the main-group
elements
-Properties of elements of the p Block vary greatly.
-At it's right hand-end, the p Block includes all of the nonmetals except hydrogen
and helium
-All 6 of the metalloids are also in the p Block
-At the left-hand side and bottom of the block, there are 8 p Block metals.
-The elements of group 17 are known as the halogens
-flourine, chlorine, bromine, iodine, and astatine
-The halogens are the most reactive nonmetals
-They react vigourously with most metals to form examples of the type of
compound known as salts
-The metalloids, or semiconducting elements, are located between nonmetals and metals in
the p Block
-The metals of the p Block are generally harder and denser than the s Block alkaline earth
metals, but softer and less dense than the d Block metals
-In the periodic table, the f Block elements are wedged between Groups 3 and 4 in the
sixth and seventh periods
-Their position reflects the fact that they involve the filling of the 4f sublevel.
-The first row of the f Block, the lanthanides, are shiny metals similar in reactivity to the
Group 2 alkaline metals
-The second row of the f Block, the actinides, are between actinium and rutherfordium.
These actinides are all radioactive.

Ch.5 Sec 1 and 2

Ch. 5 Sec 1

1. Periodic Table

• An arrangements of the elements in order of their atomic number so that elements w/ similar properties fall in same column or group
• Elements are arranged vertically in the P.T.(periodic table) in groups that share similar chemical properties

Ch. 5 sec. 2

Periods

• Elemts organized horizontally in rows
• Length of each period is determined by # of electrons that can occupy the sublevels being filled in that period
• The P.T. is divided in 4 blocks, s p d & f; the name of each block is chosen by the electron sublevel being filled in that block (disregarding Hydrogen and Helium b/c they are too small)

Alkali metals

• Elements of group 1 on P.T.
• Lithium, Sodium, potassium, rubidium, cesium, and francium are the Alkali metals
• In alkali metals's pure state, they all have a silvery appearance and are soft enough to cut w/ a knife

Alkaline-earth metals

• Elements in group 2 of P.T.
• Beryllium, magnesium, calcium, strontium, barium, and radium (all Alkaline-earth metals)
• Less reactive than alkali metals, but are still too reactive to be found in nature in pure form

• Hydrogen has an electron configuration of 1s1 but despite the ns2 configuration, it doesn't share the same properties as elements of group 1
• Hydrogen is an unique elments

• Like the Group 2 elements, helium has an ns2 group configurtaion, yet it is part of group 18
• B/c its highest occupied energy level is filled by 2 electrons, helium posses special chemical stability

P.S. remember to do Pre-Lab for tomorrow

10/19/09

Test Tomorrow

• 43 questions
• 6 calculations the rest are conceptual questions
• 100% Scantron
• know the vocab
• know all the scientists and their experiments

Sample Problems from class (there is no option for super and subscript so I will separate with commas)

Ge- 1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p6,5s2,4d9; [Ar]4s2,3d10,4p2

Mn- 1s2,2s2,2p6,3s2,3p6,4s2,3d5;[Ar]4s2,3d5

Ag- 1s2,2s2,2s6,3s2,3p6,4s2,3d10,4p2;[Kr]4d10,5s1

K- 1s2,2s2,2p6,3s2,3p6,4s1;[Ar]4s1

B- 1s2,2s2,2p1;[He]2s2,2p1

Ca2+ --> 1s2,2s2,2p6,3s2,3p6

Br- ----> 1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p6

C4- -----> 1s2,2s2,2p6

Cu2+ ----> 1s2,2s2,2p6,3s2,3p6,4s0,3d10

Al3+ -----> 1s2,2s2,2p6

K ---> n=4; l=0; m=0; s=+1/2

Cu2+ ---> n=4; l=2; m=+1; s=-1/2

Mendeleev and Chemical Periodicity

Mendeleev noticed that when the elements are arranged in order of increasing atomic mass, certain similarities in chemical properties appeared at regular intervals

repeating patterns referred to as PERIODIC

Mendeleev created a table in which elements with similar properties were grouped together----> a periodic table (or chart) of elements

After Mendeleev placed all the known elements in the table there were several empty spaces

1871--> Mendeleev predicted the existence and properties of elements that would fill in 3 of the empty spaces

by 1886---> all 3 of those elements were discovered

Moseley and Periodic Law

In 1911, English scientist Henry Moseley discovered that elements fit into patterns better when they are arranged according to atomic #

Periodic Law-->states that physical and chemical properties of elements are periodic functions of the atomic #s

Chapter 3

Law of Conservation of Mass--Established in 1789 by French Chemist Antoine Lavoisier; States that mass is neither created nor destroyed in any ordinary chemical reaction.

Law of Definite Proportions--a given chemical compound always contains the same elements in the same fixed proportion by mass

Law of Multiple Propotions--statement that when two elements combine with each other to form more than one compound, the weights of one element that combine with a fixed weight of the other are in a ratio of small whole numbers

atomic number--the number of protons in the nucleus of an atom.

atomic mass unit (amu)--a unit of mass that is exactly 1/12 the mass of a carbon-12 atom; also 1 gram/ Avogadro's number

mole--the SI unit used to measure the amount of a substance whose number of particles is the same as the number of atoms of carbon in exactly 12 g of carbon-12

Avogadro's Number--6.0221415 x 10^23

Molar Mass--the mass in grams of 1 mol of a substance

Cathode Ray Experiment

When investigators passed a current through the Cathode tube, they noticed that the surface directly opposite the cathode glowed. They hypothesized that the glow was caused by a stream of particles. They also noticed that the ray was deflected by a magnet much like an electric current was. The rays were deflected to a positive charge and away from a negative charge. This led them to believe that there was a certain type of particle that was being emitted. This was supported when J.J. Thomson found that the particles mass in the cathode tube was always the same, even when the metals where changed. Thomson concluded that all of the cathode rays were composed of the same particles, electrons.

Oil-Drop Experiment

Milikan's experiment was used to find the charge of an e-. Using an atomiser, Milikan put oil droplets in a chamber with a hole at the bottom. Some of the droplets fall through a hole in the bottom. Once though this hole, the droplets are exposed to radiation and attach themselves to free e- in the air. Then, two plates at the top and bottom have current passed through them. The top plate is negative. Milikan determined the charge of an e- by the droplets' abilities to overcome gravity when the charge of the top plate was high enough. An e- has a charge of 1.60217646 × 10-19 coulombs.

Rutherford's Gold Foil Experiment

Ernest Rutherford and his associates bombarded a strip of gold foil with fast moving alpha particles. Geiger and Marsden assumed that the partles would pass through the foil with only a small deflection. However, 1 in 8000 particles was actually deflected directly back at the alpha source. This led Rutherford to conclude that the atom was mostly empty space with most of its mass concentrated at a central point, which he called the nucleus.

Chapter 4

Electromagnetic Radiation--the energy associated with electric and magnetic fields; it varies periodically and travels at the speed of light

Electromagnetic Spectrum--range of all possible frequencies of electromagnetic radiation

Wavelength--the distance over which the wave's shape repeats

Frequency--how many waves are made per time interval, measured in Hertz (Hz)

Photoelectric effect--when light shines on a metal surface, the surface emits electrons

Ground State--The condition of an atom, ion, or molecule, when all of its electrons are in their lowest possible energy levels

Excited State--A stationary state of higher energy than the lowest stationary state or ground state of a particle or system of particles

Line-Emission Spectrum--relative intensity of electromagnetic radiation of each frequency emitted by atoms or molecules of that element or compound when they are excited

Heisenberg Uncertainty Priniciple--The position and momentum of a particle cannot be simultaneously measured with arbitrarily high precision

Orbital-- the probability distribution of an electron in a atom or molecule

Principal Quantum Number--indicates the main energy level occupied by the e-

Angular Momentum Quantum Number--indicates the shape of the orbital

Magnetic Quantum Number--indicates the orientation of the orbital around the nucleus

Spin Quantum Number--indicates the fundamental spin states of the e-

Aufbau Principle--an e- occupies the lowest energy orbital that can receive it

Pauli Exclusion Principle--no two e- can have the same set of quantum numbers

Hund's Rule--orbitals of equal energy are each occupied by one e- before any orbital is occupied by a second e- and all e- in singly occupied orbitals must have the same spin

Noble Gas--a family of nonreactive monoatomic gases found on the far right of the periodic table

Noble Gas Configuration--shortened version of electron configuration in which you simply include the closest noble gas (of a smaller atomic number) and list the electron configuration from the next element from the noble gas to the element you are doing the configuration for

The Photo-Electric Effect

In the early 1900s, scientist conducted two experiments that had results that could not be explained by the wave theory of light. One was the Photo-Electric Effect. Scientists observed that even at high intesities, a metal would not emit an electron unless the frequency was correct. This contradicted the theory that light of any frequency could supply enough energy to eject an e-. Scientists began to think that light wasn't just a wave, that it also had mass.

10/15/09

10/13/09

Today we did a worksheet in class and prepared for our test next tuesday. TEST NEXT TUESDAY. Finish the worksheet for tomorrow

Low energy

• Radio waves
• Microwaves

Medium energy

• Infra red
• Visible light

High energy

• Ultraviolet ----> goes through the magnetic field and is absorbed in ozone field
• Gamma Rays ----> deflected by magnetic field
• Cosmic Rays----> deflected by magnetic field

1m= 1x 10^9nm

Hz= 1/s=s^-1

c=3.00x10^8m/s

h= 6.626x10^-34j(s)

1MHz= 1x 10^6 Hz

1m= 1x 10^6mm

Have fun fellas

Joe's Blog Oct. 13

we turned page 126 in today, it was HW #1

47)E=hv v=E/h = 1.55 x 10^-24 J
________________ = 2.39 x 10^9 Hz
6.626 x 10^-34 J(S)

42) c. they are the same speed

46) e. 3
f. 10

49) Schrodinger used math probabilities to find electrons in a certain area
Bohr said electrons orbit the nucleus much like planets around the sun



Selenium (Se) - configuration for the last electron
1s^2,2s^2,2p^6,3s^2,3p^6,4s^2,3d^10,4p^4

l = 0 ->s
l = 1 -> p
l = 2 -> d

n = 4
l = 1
m = -1
s = -1/2
----------------
Fr
n = 7
l = 0
m = 0
s = +1/2


--------------
Ra
n = 7
l = 0
m = 0
s = -1/2



we can use the TI calculators 83-89 , change the batteries, bring 2 #2 pencils, don't randomely guess.
1.A
2.C
3.D
4.D
5.B
6.B
------
1.C
2.B
3.C
4.B
5.C
6.C
7.c
8.D
9.E
10.B
11.C
12.B
13.D
14A
15.C

Elements

N

1s2 2s2 2p3

[He] 2s2 2p3

2/1s 2/2s 3 1/2p

n=2

l=1

m=1

s=+1/2

He

1s2

1s2

2/1s

n=1

l=0

m=0

s= -1/2

Zn

1s2 2s2 2p6 3s2 3p6 4s2 3d10

[Ar] 4s2 3d10

2/1s2 2/2s2 3 2/2p 2/3s 3 2/3p 2/4s 5 2/3d

n=3

l=2

m=2

s= -1/2

Cr

1s2 2s2 2p6 3s2 3p6 4s1 3d5

[Ar] 3d5 4s1

2/1s 2/2s 3 2/2p 2/3s 3 2/3p 1/4s 5 1/3d

n=3

l=2

m=2

s= +1/2

P

1s 2s2 2p6 3s2 3p3

[Ne] 3s2 3p3

2/1s 2/2s 3 2/3p 2/3s 3 1/3p

n=3

l=1

m=1

s=+1/2

Al

1s2 2s2 2p6 3s2 3p1

[Ne] 3s2 3p1

2/1s 2/2s 3 2/2p 2/3s 1 1/3p

n=3

l=1

m=-1

s=+1/2