Fremd Chem - Recently Updated Pages

Craddock's Crack-Me-Ups

★jessie — Mon, 23 Nov 2009 04:09:05 CST

Add all those crazy, wacky, funny (or not so funny--just kidding) jokes, quotes, or random comments from the man himself: Mr. Craddock. :)

1. "And if you press this button, there's an elf downstairs that gets whipped!"
2. "Blah blah blah, blah blah blah...any questions?"
3. "...excess cat "MEOW" ion ..."
4. "...a salted battery..."
5. "...vortex of doom..."
6. "... good and what else?"
7. "... this is the correct way to waft.."
8. "that means you have a good seal erh erh(seal sounds...duh)"
9. Mr. Craddock, "What is slope?", Mr. Craddock being a crazed student, "RISE OVER RUN!!!!"
10. "If a fire starts or something, I’d have to go to court and the judge would ask, 'Is it true you were playing
Shakira in your classroom?'" --Shrugs shoulders-- “My hips don’t lie..”
11. "Today is going to be a mustard day...I mean ketchup."
12. "Carbon DIE!!! - oxide"
13. "We have this drying oven in the corner because we like to bake cookies on Fridays"
14. "EXTRA CREDIT! Good! I have your attention."
15. "So this roman guy walks into a bar and goes up to the bartender and says, 'I want 5 diet cokes'" *raises two fingers
16. "So this mushroom walks into a restaurant and the waiter says, 'sorry, we don't serve mushrooms' and the mushroom goes, 'why not? I'm a fun-gi'"
17."...puddle of death.."
18."..whirlpool of wonder..."
19."...GET MOLE, GET MOLE,GET MOLE..."
20."marco polo! marco.... MOL-O!"
21."C4 kids...BOOM!"
22. there will be an explosion and someone will be like AHHHHHHHHHHHHHH. After Yale says that...
23. "So I have a magic trick" (Starts swirling hands, covers one hand, nothing there) *gasp* "Look! It's gone"
24. "SUPER-oxide"
25. "Modified Super Soft Lab Report"
26. "'C' can be for calories. But 'c' is for cookies and that's good enough for me!"
27. (Craddock singing) "Element by itself..." (class:) "No charge!"
28. "No- this isn't silver...although it is Fremd...so why not?"
29. "Is this what this class is being reduced to?" "All this negativity!"
30. "It's a cappuccino, look at the bubbles"
31. " My personal favorite is.....OLE-!" (has to be done with the hand motion :] )
32. Mr. Craddock- "Why does this happen?" Class - "Specific Heat"...
-73h 1337 h4xx0rz
33. "Shift happens."
34. "...1.50 "FIVE OH!" moles...."
35. "Ratio" -arms moves towards self and away (alternating)- "Rate" -arms moves side to side (alternating)- aka equilibrium dance (:
36. " ICE ICE Table" (cue Vanilla Ice music)
37. Mr. C: "This is ELF blood and we know what happens when you spill Elf blod..", Jerry: "It heals you!"
38. Mr. C points to himself, "It's all muscle...just in a relaxed state"
39. Mr. Cradock tries to tell a joke but no one laughs and his response is that it will be funny next hour...
40. "I wondered...if I stapled my finger, would it go through?"

Ch. 16 and 18 Equilibrium and Solution Equilibrium

carlita.shen — Mon, 23 Nov 2009 00:09:44 CST

Equilibrium: a state where both forward and reverse reactions continue to occur at equal rates to that no net change is observed.

Kc = [Products]/[Reactants]
If you have xA + yB = zC + w D
Kc = ([A]^x)([B]^y) / ([C]^z)([D]^w)

Solids and water are not included in the equilibrium constant expression

In an equilibrium constant expression:
All concentrations are at equilibrium values
Each concentration is raised to the power of its coefficient
K depends on the reaction and temperature
Units are never given with K

ICE "ICE" Table (cue music)
I = Initial Concentration (M)
C = Change in Concentration (M)
E = Equilibrium Concentration (M)

K > 1; product favored
K< 1; reactant favored

Le Chatelier's Principle
(A change in any of the factores that determine the equilibrium conditions of a system will cause the system to change in such a manner as to reduce or counteract the effect of the change.)

Disturbance	Change in Mixture (as is returns to equilibrium)	Effect on Equilibrium	Effect on K
Heated	Heat Energy consumed by Reaction	Exothermic-shifts "left" Endothermic-shifts "right"	Changes K
Cooled	Generate Heat Energy by Reaction	Exothermic-shifts "right" Endothermic- shifts "left"	Changes K
Reactact Added/Product Removed	Reactants consumed/more product formed	Shifts "right"	No Effect
Product Added/Reactant Removed	Products consumed/Reactants formed	Shifts "left"	No Effect
Decreased Volume/Increased Pressure	Decrease in Pressure	Shift into direction of fewer moles of gas	No Effect
Increased Volume/Decreased Pressure	Increase in Pressure	Shift into direction of greater moles of gas	No Effect
Adding Pure Liquids or Solids	no effect	----------------------------------------	---------------------
Adding Inert Gases	no effect	----------------------------------------	---------------------

http://www.youtube.com/watch?v=4-fEvpVNTlE&feature=SeriesPlayList&p=166048DD75B05C0D
^^ explaining Le Chatelier's Principles

Comparing Q and K
Relative Magnitude	Direction of Reaction
Q > K	Reactants -> Products
Q = K	Reaction at equilibrium
Q > K	Reactants <- Products

Comparing Q and Ksp
Q>Ksp (over saturation point, solid will precipitate)
Q=Ksp (@equilibrium/saturation point, if you add more, will precipitate)
Q<Ksp (unsaturated, more can still dissolve, will NOT precipitate)

Regarding the general equation:
A-->B+C
We may use the formula
K=(x)(x)
A+x
If K x 100 is less than A, then we may ignore x in the denominator!

Reaction Quotient (Q): same as constant equilibrium expression but helps determine whether a system is at eqilibrium or not

Equilibrium Constant for Pressure (Kp): Constant is the product of the partial pressure of the product over the product of the partial pressure of the reactions:
xA + yB <---> zC + wD

Kp= [(partial pressure of C^z)(partial pressure of D^w)] / [(partial pressure of A^x)(partial pressure of B^y)]

In addition:

Total Pressure = Sum of partial pressures (of all gasses in the container at a given time)

Thus, total pressure may consist of the partial presssures of both reactants and products.

Ksp= solubility constant for slightly soluble solutions= [product A] x [product B]

Dalton's Law of Partial Pressures: The total pressure of a mixture of gases at a given volume is equal to the sum of th eindividual gas pressures.

Ptotal = PgasA + PgasB

Ideal gas law: PV = nRT
where n/V = mol/L or M (concentration)

Formation Constant (Kformation) : an equilibrium constant for the formation of a complex ion

pH Rules
pH = -log[H+]
[H+] = 1 x 10^(-pH)
[H+][OH-] = 1 x 10^-14
pOH = - log [OH-]
pH + pOH = 14

WHEN COMPARING Ksp VALUES:
The largest Ksp is the most soluble compound, only if they have the same number of ions.
AgIO3= 1.5 x 10^-16
CuIO3=5 x 10^-12
CuSO4=6.1 x 10^-5
Because they all have the same number of ions, CuSO4 is the most soluble!

Converting Kc into Kp
Kp=Kc(RT)^change in n
Where change in n is moles product gas- moles reactant gas
And T is measured in Kelvin (Celcius + 273)

Different K values:

Keq / Kc : concentrations are in molarities, standard K
Kp : for gasses, uses partial pressures
Ksp = solubility product constant, uses molarities of slightly soluble compounds in aqueous solutions

Ch. 6 Energy

Vudoo23 — Wed, 04 Nov 2009 11:59:15 CST

Energy

The Ability to do work
Bonds store chemical energy (break bond=release, make bond=use energy)
Kinetic Energy: energy in motion
Potential Energy: stored energy

Heat

The amount of the thermal energy
Extensive

Temperature

Measurement of average kinetic energy of a substance
Intensive

Law of conservation of energy: energy can be neither created or destroyed, only transferred into different forms
1st law of thermodynamics: the energy of the universe is constant
Heat (total thermal energy/depends on amount) vs. Temperature (measure of average kinetic energy)

Exothermic-heat is transfered from a system to the surroundings. (Positive q)
Endothermic-heat is transfered from the surroundings to a system. (Negative q)

(when graph above has a slope)
q = mC(ΔT)
q is the heat transferred (J)
C is the specific heat capacity (J/g*K)
m is the mass of substance (g)
delta "t" is the change in temperature (K)

(when graph above slope is 0 )
q= mL
m is the mass of the substance (g).
L is either the latent heat of fusion or vaporization, depending on if the substance is melting or freezing. Can also be notated as ΔH (J/g)

· Sum of the heat content changes within a given system is zero

Energy Units

1 Joule = 1 kgm^2/s^2
1 calorie = 1 cal = 4.184 J
1 Calorie = 1000 cal = 1 kcal = 4184 J = 4.184 kJ (food calorie)

HH
Specific Heat Capacity (c): the amount of energy required to raise the temperature of one gram of the substance 1 K (Units: J/gK)

Energy is transferred between system and surroundings only as heat or work

system expands --> does work and is negative
system contracts --> work is done upon it and is positive

*Standard State-Most stable form of a substance in the physical state that exists at a pressure of 1 bar and at a specified temperature (usually 25 degrees celcius or 298 K).

Enthalpy Change - state function (Not dependent on path taken, only final/starting conditions)

1) When you reverse an equation, the magnitude of enthalpy change is the same but the sign is opposite.

2) Depends on number of moles (Ex. # of moles x 2 = enthalpy x 2)

3) State of matter specific

Hess's Law - If you add all of the delta H's up you get the delta H of the given equation.

Chem Demo Videos

jchiang — Fri, 30 Oct 2009 19:56:35 CDT

In-Class Demos

Special Occasion Demos:

Mole Day:

Ch. 5 Reactions In Aqueous Solutions

NicoleGut1 — Thu, 22 Oct 2009 00:55:53 CDT

Different Types of Reactions
1. Two Uncombined Elements (Synthesis Reactions)

Combine the elements

2. Single Reactant (Decomposition Reaction)

Break up the reactant

3. A reactant is water

A. A pure metal or a metal hydride
A base and hydrogen gas are created
B. Metal oxide is in water
A base is created
C. A nonmetal is in water
An acid is made

4. Acids and Bases

A. Acid and a base B. An acid and a basic salt C. A base and an acid salt D. A base and an acid anhydride

5. Two salt solutions are mixed

Using solubility rules, discover which salt precipitates

6. Carbon Compound Combustion

Generally, carbon dioxide and water are created

7. Solution and a solid transition metal are combined

A. A neutral metal salt solution is the solution (Single Replacement)
The metals will switch places
B. The solution is a strong oxoacid solution
An oxide gas and water forms from the anion of the oxoacid

8. Transition Metal Ions are in a solution with ammonia, hydroxide, cyanide, or thiocyanate

The ions and the other compound combines to make a complex ion

VOCAB:
solution: homogeneous mixture of two or more substances
solvent: the medium in which another substance, the solute, is dissolved
electrodes: conductors of electricity such as copper wire
electrolytes: compounds whose aqueous solutions conduct electricy (includeds all ionic compounds that are soluble in water)

strong electolytes: good electrical conductors owing to the presence of ions
weak electrolytes: poor conductors of electricy (i.e. acetic acid in water)

non-electrolytes: substances that dissolve in water but don't ionize

Tips for making a solution:
*Remember, concentration does not equal strength! (0.1M HCl is a dilute, strong acid)
-Volumetric flask to make solutions, fill up below the line
-Do as you oughta, add acid to wata!
-Use the M1V1=M2V2 equation (M=moles, V=volume)

Types of transfer reactions:

Atom/Group transfer (HCl + H2O > Cl- + H3O+)
Electron transfer (Cu(s) + 2Ag+(aq) > Cu2+ + 2Ag)
- oxidation reduction (redox): can result in generation of an electric current or by imposing a current

Oxidation-Reduction Reactions
These reactions occur when there is a transfer of electrons between the reactants in a reaction.

Terminology:
Oxidation- loss of electrons by a species (increase in oxidation #)
Reduction- gain of electrons (decrease in oxidation #)
Oxidizing agent- electron acceptor (species is reduced)
Reducing agent- electron donor (species is oxidized)

Ways to remember which reactant is being oxidized/reduced:

OIL RIG -->Oxidation is losing(electrons), Reduction is gaining (electrons)
LEO (the lion) goes GER-->Losing electrons = oxidation, gaining electrons = reduction
OLE'--> Oxidation losing electrons

Balancing Equations:

Divide the reaction into half-reactions, one for oxidation and the other for reduction.
Balance each for mass.
Balance each half-reaction for charge by adding electrons.
Multiply each half-reaction so electron on both sides are equal.
Add half-reactions up. Cancel things that are the same on both sides. You will always cancel the electrons from both sides.
Re-write overall reaction.

Tips for balancing:

If a reaction says that it occurs in an acid solution, you will be adding hydrogen and water molecules to the half-reactions to balance each for mass.
If a reaction says that it occurs in a basic solution, you will be adding hydroxide and water molecules to balance the half-reactions for mass.
The hydrogen, hydroxide, and water molecules may be eliminated or reduced in the final equation.
Never add O2, O atoms or O2- to balance O
Never add H2 or H atoms to balance H

Oxidation Number Rules:

1) All net molecules have a net charge of zero; furthermore, the sum of the positive charges equals the sum of the negative charges
2a) The charge of all free (uncombined) elements is zero.
b) The charge of all metals of Group IA is +1
c) The charge of all metals of Group IIA is +2
d) The charge on aluminum is +3
e) The charge on hydrogen is +1 (EXCEPT hydrides which = -1 when combined with alkai metals
f) The charge of fluorine is -1
g) The charge on other halogens is -1. The charge is positive when it is linked with a higher electornegative element.
h) The charge on oxygen is -2 EXCEPT peroxides (H2O2 & Na2O2 - -1 charge)

Molarity (Concentration) = amount of solute (mol) / volume of solution (L)

Colors of Ions in Aqueous Solutions (found this on Wikipedia)

Alkali metals	M+	None
Alkaline earth metals	M 2+	None
Scandium (III)	Sc 3+	None
Titanium (III)	Ti 3+	Violet
Titanyl	TiO 2+	None
Vanadium (II)	V 2+	Lavender
Vanadium (III)	V 3+	Dark grey/green
Vanadyl	VO 2+	Blue
Pervanadyl	VO2 1+	Yellow
Metavanadate	VO3 1-	None
Orthovanadate	VO4 3-	None
Chromium (III)	Cr 3+	Blue-green
Chromate	CrO4 2-	Colorless or Yellow(sometimes)
Dichromate	Cr2O7 2-	Orange/colorless
Manganese (II)	Mn 2+	Light pink/clear-colorless
Manganate (VII) (Permanganate)	MnO4 1-	Deep violet
Manganate (VI)	MnO4 2-	Dark green
Manganate (V)	MnO4 3-	Deep blue
Iron (II)	Fe 2+	Light blue
Iron (III)	Fe 3+	Yellow/brown
Cobalt (II)	Co 2+	Light red
Cobalt-ammonium complex	Co(NH3)6 3+	Yellow/orange
Nickel (II)	Ni 2+	Light green
Nickel-ammonium complex	Ni(NH3)6 2+	Lavender/blue
Copper (II)	Cu 2+	Blue
Copper-ammonium complex	Cu(NH3)4 2+	Royal Blue
Tetrachloro-copper complex	CuCl4 2-	Yellow/green
Zinc (II)	Zn 2+	Bluish-white/colorless
Silver	Ag+	None

Solubility Rule

Remember that all strong acids and strong bases fully dissociate in water.

Common Strong Acids:
Hydrochloric acid (HCl)
Nitric acid
Sulfuric acid

Common Weak Acids
Phosphoric
Carbonic
Acetic
Oxalic

Strong Bases:
Lithium Hydroxide LiOH
Sodium Hydroxide NaOH
Potassium Hydroxide KOH

Weak Base:
Ammonia NH3

Learning from Labs

joannechung — Tue, 20 Oct 2009 17:37:09 CDT

I made this page so that we have a place to put everything we learned from labs that we'll need to know for our AP exam.
Ex: compound colors

	Co +2 Br2	Pink Brown
	Mn +7 (MnO4-)	Dark Purple
	Ni +2	Green
	S	Yellow
	I2	Violet
	Cu +2	Blue
	Pn	Red
	F2	Yellow
	P4	White
	Cl2	Green
	Mn +2	Pink

Ch. 20.1 Oxidation-Reduction Reactions

Vudoo23 — Fri, 16 Oct 2009 12:57:01 CDT

Balancing Equations

Acidic Solutions
1. Find the which element is being reduced or oxidized.
2. Half reactions
3. Balance by mass
4. Balance the oxygens (if any) by adding H20
5. Balance the hydrogens (if any) by adding H+
6. Balance the charges by adding electrons
(DOES NOT HAVE TO EQUAL ZERO!!! As long as the charges are equal.)
7. Multiply the needed number to have equal electrons
8. Cancel any factors that you can.
9. Add the final balanced equations together
10. Check the masses and the charges. (BE SURE TO CHECK YOUR ANSWER!!!)

Basic Solutions

Ch. 4 Stoichiometry and Basic Equations

Kristin17 — Mon, 28 Sep 2009 22:02:10 CDT

Chemical Equations
reactants: substances combined in a reaction
products: substances produced in a reaction
law of conservation of matter: matter cannot be created or destroyed
combustion: a reaction with oxygen in which all of the products are combined with oxygen
limiting reactant: the reactant in the compound which sets the amount of product which can be formed
- How to find the limiting reactant Step 1. Find the amount of each reactant (g into mole) Step 2. What is the limiting reactant? Examine the ratio of amounts of reactants Step 3. Calculate the mass of product Step 4. Calculate the mass of excess reactant.

(s) = solid
(g) = gas
(l) = liquid
(aq) = aqueous (substance dissolved in water)

Basics of Stoichiometry: Mass Relationships

It is imperative to obtain a well-balanced equation of a chemical reaction to be able to see the mass/molar relationships between the reactants and the products.

Here is an example problem:

Phosporous reacts with chlorine gas and yields phosporous trichloride.

1) get the balanced equation.

P4 (s) + 6 Cl2 (g) -> 4 PCL3 (l)

2) if the problem asks for the mass of chlorine required for the reaction, you must set up a stoichiometric factor, but first...

a) calculate the moles of P4 in the equation

if we had 1.45g of P4, then there are .0117 moles of P4 (1 mol of P4 has 123.9 g)

b) then set up the factor.

basically, we want to relate the amount of P4 available to the amount of chlorine (based on the balanced equation)

.00117 mol P4 (6 mol Cl2/ 1 mol P4) = .0702 mol Cl2.

This is the amount of chlorine required. Notice that in the balanced equation there are six moles of chlorine gas to one mole of phosphorous. This is how you set up the stoichiometric factor.

3) You can then convert the amount of chlorine to mass.

.0702 mol Cl2 (70.91 g Cl2 / 1 mol Cl2) = 4.89 g Cl2

4) There are two ways to calculate the amount of product. One way is to keep in mind the fact that matter is neither created nor destroyed. If there are 4.88 g Cl2, and 1.45 g of P4, then the product should have a mass equal to their sums (6.43 g). The other way involves setting up another stoichiometric factor.

.0117 moles of P4 (4 moles of PCl3 / 1 mole of P4) = .0468 moles of PCl3

.0468 moles of PCl3 (137.3 g PCl3 / 1 mol PCl3) = 6.43 g

Notice that you get the same mass with both methods.

Balancing equations: The coefficients of a balanced equation represent the numbers of moles or molecules of reactants and products.

Hints for balancing equations:

Equations are balanced by adjusting coefficients in front of formulas
It is best to start with an element that appears in only one species on each side of the equation
Usually the simplest whole number coefficients are used in balancing an equation

Balancing equations with organic compounds:

Balance the Carbons
Balance the Hydrogens
Balance the Oxygens

To calculate the molecular mass of a hydrate:

calculate the mass of the compound without the water
multiply the mass of water by the coefficient in front of the water molecule
add the two masses

Percent Yield

Percent Yield = Percentage Recovered X 100
Theoretical Amount Expected

Theoretical Yield- maximum quantity of product we calculate can be obtained from a chemical reaction

Actual Yield- the quantity of material that is obtained in the laboratory

Percent Yield- percentage of how much of the theoretical yield was actually obtained

Don't forget the Hydrocarbons!
(You really only need to memorize the first four. the rest are common sense)

CxH(2x+2)
where x is the number of carbons

::::::::::::::::::::::
Methane - 1
Ethane - 2
Propane - 3
Butane - 4
Pentane - 5
Hexane - 6
Heptane - 7
Octane - 8
Nonane - 9
Decane - 10
::::::::::::::::::::::

Chapter 1 - Matter and Measurement

ⓙⓐⓔ — Sat, 12 Sep 2009 17:46:06 CDT

1. Matter

Solid, Liquid, Gas

2. Kinetic-Molecular Theory

The theory assists us in determining the properties of the states of matter
KMT states all matter contains tiny particles that are constantly in motion (Atoms, ions, molecules)
- Solids: Particles are packed closely and vibrate slightly in place.
- Liquids: Random arrangement of particles where the particles can move freely and are not confined.
- Gas: Also a fluid. Particles move extremely fast, colliding with others and the walls.
The higher the temperature, the faster the particles move
The energy of the motion overcomes the force of attraction between the bonds, which eventually translates it into a different state of matter

Solid/Liquid/Gas

/ \
homogeneous mixture heterogeneous mixture
/ \.............

pure solution ........
/ \ ....................
compound element..........................

Homogeneous Mixture: substances in the mixture are in the same phase, also called a solution, i.e gasoline
Heterogeneous Mixture: Uneven texture of the material can be detected, i.e chicken noodle soup
Chemical Compound- a pure substance, like salt or water, which is composed of 2 or more different elements held together by a chemical bond
Element: Composed of only one type of atom
Atom: The smallest particle of an element that retains the characteristic chemical properties of that element.
Ions: Electrically charged atoms/group of atoms
molecules: smallest unit that retains the composition and characteristic of a compound

Density: Density= mass/volume

Extenstive & Intensive Properties

Extensive properties: dependent on the amount of a substance (ex: density)

Intensive properties: not dependent on amount of a substance (ex: temperature)

Qualitative & Quantitative Observations

Qualitative Observations: No measurements and numbers (i.e. color, appearance)
Quantitative Observations: Have numerical information (i.e. time, mass, volume, length)

Precision & Accuracy

Precision: how similar several determinations of the same quantity are
Accuracy: the closeness of a measurement with the accepted value of the quantity

Significant Figures

All digits that are not zero are significant
For numbers greater than 1: all numbers are significant after the first nonzero number if there is a decimal point; in the case of trailing zeros , none of the zeros are significant figures unless they are followed by a decimal point (i.e. 1000 has 1 significant figure while 1000. has four)
For numbers less than 1: zeros to the right of the first nonzero number are significant
Defined quantities like 1000 grams per kilogram, do not limit the amount of significant figures
Addition/ Subtraction: keep the same number of decimal places as the factor with the least amount
Multiplication/ Division: keep the same number of sig figs as the factor with the least number of sig figs.

Chapter 3 - Molecules, Ions, and Their Compounds

xX_ThE_uSeRnAmE_Xx — Sat, 12 Sep 2009 12:38:49 CDT

3.1 Molecules, Compounds, and Formulas

-Molecule: smallest unit of an element or a compound; pure substance which can be divided and still have its physical and chemical properties
-Elements do not retain their properties when in a compound
Formulas
-Molecular Formula: the formula of a molecule which states how many atoms of each atom are in a molecule of the said compound
ex: C2H6O
-Condensed Formula: a formula of a molecule which explains how the individual atoms are grouped
ex: CH3CH2OH
-Structural Formula: a formula which specifically states how each atom is bonded to the other atoms in the said molecule
ex:

-Ball-and-stick Model: a model which uses spheres to represent atoms and sticks to represent bonds
Carbon: black/gray
Hydrogen: white
Oxygen: red
Nitrogen: blue
Chlorine: yellow
ex:

Nomenclature:

Polyatomic ions have their own unique names that must be memorized in order to facilitate the naming of various compounds.

AsO4-3 arsenate BO3-3 borate B4O7-2 tetraborate BrO3- bromate BrO- hypobromite CO3-2 carbonate CN- cyanide C2O4-2 oxalate C2H3O2- acetate C4H4O6-2 tartrate ClO4- perchlorate ClO3- chlorate ClO2- chlorite ClO- hypochlorite CrO4-2 chromate Cr2O7-2 dichromate IO4- periodate IO3- iodate IO- hypoiodite HCO3- hydrogen carbonate (bicarbonate) HSO4- hydrogen sulfate (bisulfate) HSO3- hydrogen sulfite (bisulfite) HC2O4- hydrogen oxalate (binoxalate) HPO4-2 hydrogen phosphate H2PO4- dihydrogen phosphate HS- hydrogen sulfide MnO4- permanganate NH2- amide NH4+ ammonium NO3- nitrate NO2- nitrite OH- hydroxide O2-2 peroxide PO4-3 phosphate PO3-3 phosphite SCN- thiocyanate S2O3-2 thiosulfate SO4-2 sulfate SO3-2 sulfite SeO4-2 selenate SiF6-2 hexafluorosilicate SiO3-2 silicate

Chapter 2 - Atoms and Elements

Jeni033 — Thu, 10 Sep 2009 23:03:50 CDT

Electricity

2 Types of Electrical Charges
-Positive/Negative

Radioactivity

Discoveries Dealing with Radioactivity
-Becquerel discovered uranium ore emitting rays and darkening photographic plate even though covered with black paper to protect from light.
-Curie isolated polonium and radium, which emitted rays like uranium did in Becquerel's example.

3 Types of Radiation
-Alpha Rays

Helium 2+ ions

-Beta Rays

Electrons

-Gamma Rays

Photons / Light

Cathode Ray Tube

-Ray went from cathode (negative) to anode (positive) so the beam was composed of electrons
-Cathode Rays travel in straight lines, cause gas to glow, can heat metal objects, be deflected in a magnetic field
-Thompson determined the charge to mass ratio of a beam of cathode rays
-Milllikan found the measurement for an electron's charge using the oil-drop experiment

Equations

Percent abundance= (# of atoms of a given isotope/# of atoms of all isotopes of that element) x 100%

Atomic Weight= (% abundance of isotope 1)/100 x (mass of isotope 1)
+ (% abundance of isotope 2)/100 x (mass of isotope 2) +....

AP Chemistry 09-10

shahrk — Thu, 10 Sep 2009 21:43:28 CDT

< 5/11/2010.

Fremd Chem fb Page

Intro Materials Ch. 1-3

shahrk — Fri, 28 Aug 2009 20:02:53 CDT

Click on links to follow:

AP Chemistry 08-09

MrCraddock — Wed, 12 Aug 2009 22:43:48 CDT

This page is for you to show what you know and to ask questions of the group!

I would like you to actually add content by hitting the easy edit button at the top more than just adding threads. The threads are for discussion. The pages are for content. It will take a while, but we will get it.

GO AP CHEM!

Classroom

MrCraddock — Mon, 22 Jun 2009 12:37:46 CDT

Welcome to the FremdChem wiki!

This wiki is for Craddock's students to show what they know and discuss it with others.

	AP Chemistry
		[add a new page and link to it here]

Need more? Add graphics and flourishes that match your wiki design (they're free from Paint Splatters)

Calorimetry

reenapatel21 — Mon, 11 May 2009 16:33:12 CDT

Calorimetry: Measuring heat transferred by chemical or physical properties

“Coffee Cup” Calorimetry:
heat gained = - heat lost
q system + q surroundings = 0

Bomb Calorimetry:
q bomb + q system + q surroundings = 0

the specific heat of the bomb is a given number. the q of the bomb is the mass of the bomb times the c of the bomb.
Usually, the q of the bomb and the q of the surroundings are together because the heat lost by the system is gained by both the bomb and the surroundings.

Electron Transfer Reactions

dr.whoo — Fri, 08 May 2009 08:37:51 CDT

I (amp) = C (coulumbs)/T (seconds)

Voltaic/galvanic cells - devices that use chemical reactions to produce an electric current

batteries are example of voltaic cells

Electrolysis- use of electric energy to effect a chemical change
Anode- the electrode at which oxidation occurs; negative ("an ox")
Cathode- electrode at which reduction occurs; positive ("red cat")
*the anode and cathode for electrolytic systems is the opposite those of a galvanic
Electrolytic Cell- electric current used to cause chemical change (reactant favored reactions)
Electromotive Force (emf)- drives the electron from the anode to the cathode
Standard Cell Potential-a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 degrees Celsius

Using Standard Cell Potientals:
-all half reactions listed in the table can occur in either directions
-the more positive the reduction potential, the more likely the reduction of a substance will occur and the less likely the oxidation of the substance will occur
-when presented w/ 2 redox-active substances, the species reduced witll be the one w/ the more positive value for E red

Oxidation- loss of electron(s) by a species; oxidation number increases
Reduction- gain of electron(s) by a species; ox. number decreases

Direct redox reaction- both agents in direct contact. Single replacement. No current produced (no potential difference)
Indirect redox reaction- electrons transferred through external wire (from red. agent --> ox. agent)

Nernst Equation
E= Eo - (RT/nF)ln(Q)
E= potential of a cell under nonstandard condtions
Eo= standard cell potential
n=moles of electrons transferred
F= Faraday's constant (9.64853 x 10^4 J/V(mol))
R= Gas constant (8.3145 J/K(mol))
T= temperaute (K)
Q= Reaction quotient ([Product]/[Reactant])(M)

Eo and its relation to change of free energy

∆G= -nFEo
n= the number of moles(electrons) transfered
F= Faraday constant (9.65853 x 10^4 J/Vmol)
Eo= standard cell potential

Product favored reactions yield ∆Go < 0 and so Eo > 0
Reactant favored reactions yield ∆Go > 0 and so Eo < 0

Determining which way a reaction proceeds:
(#1) Cd --> Cd(2+) + 2e- or (#2) Cd(2+) +2e- --> Cd

and

(#1) Fe --> Fe(2+) + 2e- or (#2) Fe(2+) + 2e- --> Fe

Since Fe is the most bottom right on our chart of standard reduction potentials (in aqueous solution at 25 degrees Celsius), it is the most likely to be oxidized (and also the best reducing agent). Therefore, the reaction will start here and we use the first Fe half reaction
Fe --> Fe(2+) + 2e-
Then we pick the Cd half reaction that balances out the two electrons on the right side, and so we use the second Cd half reaction:
Fe + Cd(2+) + 2e- --> Fe(2+) + 2e- + Cd

Principles of Reactivity: Electron Transfer Reactions

Maysen — Sun, 03 May 2009 19:30:08 CDT

anode - electrode at which oxidation occurs
cathode - electrode where the reduction occurs
salt bridge - allows cations and anions to move between the two half-cells

Entropy and Free Energy

Maysen — Mon, 27 Apr 2009 21:18:42 CDT

=ΔH= + / ΔS= -ΔH= + / ΔS= -
system-part of the universe under study
surroundings- the rest of the universe
state function- a quantity whose value is only determined by the initial and final states of the system
Standard conditions- Pressure of 1 bar (0.98692 atm) and solution calculations of 1 molal

Thermodynamically favored reactions are product favored. Many, but not all are exothermic.

Entropy (S) is the measure of dispersal or disorder. CHAOS!!!!!!!!

ΔS =∑S(products) -∑S(reactants)

units for S = J/(K x mol)

*each S value needs to be multiplied by its coefficient of moles in the balanced equation*

A pure crystalline solid at absolute zero has an entropy of zero (3rd law of thermo). Therefore, all entropies are POSITIVE under standard conditions.

Entropy is higher for:

when heat is added
gases go to liquids, and liquids go to solids
more complex molecules
mixtures of substances than for pure substances
when a pure liquid or solid dissolves in a solvent
when an ionic solid has a lower attractive force (NaF (+1, -1) has more entropy than MgO (+2, -2))

ΔS of universe= ΔS of surrounding + ΔS of system
ΔS of surrounding = q of surrounding/ T(temperature in Kelvin) = -ΔH of system/ T (temperature in Kelvin)

Second law of thermodynamics - in a spontaneous process, the entropy of the universe increases.

*in better terms, the universe favors increasing entropy, allowing reactions favoring disorder to become spontaneous
The Final State of a System can be more probable than the initial state by either or both of two ways:
1. The atoms and molecules can be more disordered
2. Energy can be dispersed over a greater number of atoms and molecules

- If energy and matter are both dispersed in a process, it is spontaneous
- if only matter is dispersed, quantative information is needed to decide whether a process is spontaneous
- if energy is not dispersed after a process occurs, then that process will never be spontaneous
ΔG=ΔH-TΔS
-spontaneous when ΔG is negative and vice versa

Gibbs free energy (G) - amount of energy that is availabe to do work
G = ΔH-TS

ΔG>0 => reactant favored/ not spontaneous/ is spontaneous in reverse
ΔG<0 => product favored/ spontaneous/ not spontaneous in reverse
ΔG=0 => Equilibrium

Third Law of Thermodynamics- there is no disorder in a perfect crystal at 0 K; S=0
ΔS= qrev/T
- all substances have positive entropy values at temperatures above 0 K
- when comparing the same or similar substances, entropies of gases are much larger than those for liquids, and entropies of liquids
are larger than those for solids.
-larger molecules have a larger entropy than smaller molecules, and molecules with more complex structures have larger entropies
than smaller molecules

∆G = ∆G° + RTlnQ
where ∆G = under nonstandard conditions
∆G° = under standard conditions
R = 8.31 J/mol K
T = temperature in K
Q = reaction quotient ([products]/[reactants])

Equilibrum
∆ G= -RTln(K)
if k>1, rxn=spontaneous
if k<1, rxn= not spontaneous
- when ΔG(rxn)<0, Q<K => reaction proceeds spontaneously to convert reactants to products until equilibrium.
- when ΔG(rxn)>0, Q>K => reaction proceeds spontaneously to convert products to reactants until equilibrium.
- when ΔG(rxn)=0, Q=K => equilibrium C:

Situation	Signs on ΔH & ΔS	Signs on ΔG	Comments
1	ΔH= - / ΔS= +	ΔG= - (all temps)	Both factors favor a spontaneous reaction. There is no temperature at which the reaction would not occur.
2	ΔH= + / ΔS= -	ΔG= + (all temps)	Both factors are against a spontaneous reaction. There is no temperature at which the reaction can be made to occur! [Hopeless case!]
3	ΔH= - / ΔS= -	ΔG= + @ high temps ΔG= - @ low temps	The reaction will not occur at high temps, but will occur at low temps.
4	ΔH= + / ΔS= +	ΔG= - @ low temps ΔG= + @ high temps	The reaction will occur at high temps, but will not occur at low temps.

SPONTANEITY OF REACTIONS
kdldkjf
when ΔG is -, a reaction is spontaneous; when ΔG is +, a reaction is nonspontaneous.
when the signs of ΔS and ΔH are different, the reaction will ALWAYS be either spontaneous or nonspontaneous.
when the signs of ΔS and ΔH are the same, the reaction may either be spontaneous or nonspontaneous, depending on the Kelvin temperature.
When a reaction is spontaneous, the reaction will occur.
When a reaction is not spontaneous, there is no reaction.

Gibb's Free Energy:
@ equilibrium ---
ΔGnot = -RT*Ln(K)

Acid/Base Equilibrium

Sammy_Adhikari — Fri, 17 Apr 2009 08:32:39 CDT

Acids and Bases, or more specifically, H+ ions and OH- ions always are proportionate within the solution of water. Unfortunately for us and the rest of the people that must deal with the numbers involved, the proportion is not at all linear. the concentration of H+ ions [H+] in any given solution is 10^(-pH) M.

Arrhenius definition of acid/base:
-acids increase H+ or hydronium concentrations in water
-bases increase OH or decrease hydronium concentrations in water

Bronsted-loury definition:
-Acids=proton donors
-Bases=proton acceptors

*All Arrhenius acids/bases are Bronsted-Lowry acids/bases, but not all Bronsted-Lowry acids/bases are Arrhenius acids/bases.

Monoprotic-donates 1 proton
Amphiprotic-can be either Bronsted acid or base
Amphoteric-can be either Lewis acid or base
Polyprotic- donates more than 1 protons (i.e:sulfuric, phosphoric, and oxalic acid)

*adduct = product of acid/base reaction

How to Tell if a Substance is an Acid or a Base:

Anions that are conjugate bases of strong acids are such weak bases that they have no effect on the pH (it is neutral)
Alkali metals and alkaline earth metals cations have no effect on the pH (it is neutral)
There are many basic anions - All conjugate acids of a weak acid.
Anions from polyprotic acids can be either acids or base depending on the KA
Acidic cations will be 2+ or 3+ charges in the transition metals. They act as Lewis Acids and hydrate in water.
When given something like HCO3 and you have to figure out if it acts as a base or an acid, react it with water as both an acid and a base. Then compare the Ka and the Kb of both reactions. If the Ka>Kb then the solution would be acidic.

Formulas:

pH= -log [H3O+]
pOH= -log [OH-]
pH +pOH=14
[H3O] = 1x10^-pH
[OH-]= 1x10^-pOH
Kw= [H30+][OH-]= 1x 10^-14
Ka * Kb = Kw

8. Ka= [H30+][A-]/[HA]

9. Kb=[BH+][OH-]/[B]

Henderson-Hasselbach Equation:

pH = pKa + log [A-]/[HA] pKa= -log Ka

pOH = pKb + log [HB+]/[B] pKb = -logKb

This equation is helpful when creating a buffer solution. A buffer causes solutions to resist a change in pH when a strong acid or base is added. If given a specific pH, put the number in for pH and Ka of the weak acid/base and you can determine the ratio of conjugate : weak acid.
The ratio of conj:acid cannot exceed 1M :0.1M. If exceeds, no longer buffer.

In the Henderson-Hasselbach equation, if the ratio of the conjugate to its subsequent acid/base is 1, the log of 1 is 0 so pH=pKa but when ratio>1, pH>pKa and vice versa.

IF KA <10^-5 or so (really small), you can ignore x. So instead of KA= [x^2]/[HB-x], you can write it as KA= [x^2]/[HB].
Ex) What is the pH of a 0.5 M solution of acetic acid. (1.8 x 10^-5= KA at 25 degrees Celsius)
1.8 x 10^-5 =[x^2]/[0.5]
x^2= 0.000009
x= 0.003, which is [H+]
pH= -log [0.003]= 2.522

Lewis Acid/Base
electron acceptor = acid (AAE--acids accept electrons)
electron donor = base

if Ka > 1 = strong acid
if Ka < 1 = weak acid

if Kb > 1 = weak acid
if Kb < 1 = strong acid

STRONG ACIDS: completely disassociate in water into H+ and cations

HCl, hydrochloric acid
HBr, hydrobromic acid
HI, hydriodic acid
HNO3, nitric acid
HClO4, perchloric acid
HClO3, chloric acid
H2SO4, sulfuric acid

STRONG BASES: completely disassociate in water into OH- and anions

NaOH, sodium hydroxide
KOH, potassium hydroxide
CsOH, cesium hydroxide
Ca(OH)2, calcium hydroxide

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Neutral solution: [H3O+] = [OH-] = 1x10^-7 M

equal quantities of strong acid + strong base = neutral

Acidic solution:[H3O+] >1x10^-7 M and [OH-] <1x10^-7 M, so [H3O+] > [OH-]

equal quantities of strong acid + weak base = acidic

Basic solution:[H3O+] < 1x10^-7 M and [OH-] > 1x10^-7 M, so [H3O+] < [OH-]

equal quantities of strong base + weak acid = basic

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How to know if an acid or base is stronger:

The basicity of an anionic base increases substantially as the negative charge of the anion increases.
+3 > +2 <--- is this right? i wish i knew what was going on so i could correct it. -HS

The addition of an oxygen atom in an oxoacid group increases the acid strength because it changes the anion structure to make it more stable if the H+ is removed.

O3 > O2

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Acid-Base Reactions

Strong Acid + Strong Base
- net ionic = H30+ + OH- ~> H20
- very large K value = all product
- pH depends on excess reactant, if any. the remaining reactant then dissociates into either its H+ or OH- ion which you can then plug into the pH/pOH equation
Strong Acid + Weak Base
- the strong acid dissociates into the hydronium ion
- the hydronium then reacts with the OH- ions after the weak base undergoes hydrolysis
- if at the equivalence point, it will form an acidic solution (NOT BECAUSE OF STRONG ACID THOUGH)
- the pH at the eqivalence point is driven by the conjugate acid of the weak base undergoing hydrolysis and reforming weak base and H+
  - the degree at which the reformation of weak base and H+ happens depends on the Ka of the conjugate acid.

Eg: NH3 + H20 <~> NH4+ + OH- Kb = 1.8 x 10^(-5)
H30+ + OH- <~> 2 H20 K = 1.0 x 10^(-14)
H30+ + NH3 <~> H20 + NH4+ Knet = 1.8 x 10(-9)

3. Weak Acid + Strong Base

same as Strong Acid + Weak Base but opposite
Important parts to see on the curve

1. the pH before titration begins
2. the pH at the equivalence point

driven by the conjugate base of the weak acid undergoing hydrolysis and reforming weak acid and OH-

3. the pH at the halfway point (half-equivalence point)

at this point, the concentration of weak acid and conjugate base are equal. This makes the ratio in the Hans equation 1 SO pKa = pH
This is buffer zone

4. Weak Acid + Weak Base

whichever K is greater determines the pH (if at equivalence)
double buffer situation: one before equivalence point and one after

Solving the pH of an Acid-Base Solution

Calculate the Stoich
Calculate the Concentrations
Find the pH

*Polyprotic Acids depend primarily on the first acidic hydrogen's Ka
Equivalence Point- point at which moles of acid = moles of base

- pH is not necessarily 7 at equivalence point
strong acid/strong base pH = 7
strong acid/weak base pH < 7
weak acid/strong base pH > 7
weak acid/weak base Depends on Ka and Kb...the larger one influences pH

Buffer- keeps pH at a constant level
Buffer Solution- An aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid.
Buffer Capacity: A measure of how much hydronium or hydroxide can be added to a buffer solution before the buffer can no longer control the pH.
Common Ion Effect: 1) Adding the conjugate base (or conjugate acid) to a solution of an acid (or base) increases the concentration of the acid (or base). 2) Adding a common ion to a saturated solution of a salt will lower the salt solubility.

To create a buffer solution:

1.) You need a weak acid or base and its conjugate.

2.) The amounts of both should be at the most a 1 M: 0.1 M ratio.

3.) pKa should be close to the desired pH.

Buffers are common ion problems. When you have a common ion, it lowers the overall solubility.

Titrations between a weak acid and a weak base are usually not done, because the equivalence point is hard to judge correctly.

Polyprotic Acids:
A titration graph between a diprotic acid with a strong base has a little bump in the curve. This differs from the curves of a strong acid titrated with a strong base and a weak acid titrated with a srong base, a strong acid with a weak base and a weak acid with a weak base. Those curve do not have a extra bump in the curve.

Inductive Effect: extent to which adjacent atoms or groups of atoms attract electrons from another part of a molecule. It comprises the attraction of electrons from adjacent bonds by more electron-negative atoms.

*oxides of nonmetals are acid anhydrides
*oxides of metals are basic anhydrides
*adding water reacts to form acid/base

the more oxygens an oxoacid has, the stronger the acid [H2SO4 > H2SO3.. HClO4 > HClO3 > HClO2 > HClO] because the electronegative oxygens draw electron density away from the central atom and therefore from the H-O bond, making the substance more polar so the H leaves more easily.
the more electronegative the central atom, the stronger the oxoacid [H3PO4 < H2SO4 < HClO4] because the more electronegative atom in the center makes the O-H bond more polar.
binary acid strength depends on the size of the atom [HF < HCl < HBr < HI] because greater distance means weaker attraction.