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

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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?