Assignment: Read Section 1.1



Assignment: Read Section 1.2



Assignment: Read Section 1.3


Dependent variable

Independent variable


Qualitative Data

Quantitative Data

Scientific law


Assignment Read Section 1.4

Pure research

Applied research

Safety Rules on p. 19 Table 1.2

Icebreaker (Theory, Substance, Mixture, Metric prefix and meaning, Something in common)

CHEMLAB Identify the Water Source

_Discussion Assignment: Visit our Wiki at

Pick out an article under the article tab from the selections
posted on the board. Read one and post a commentary on it under your name on our wiki. See the rubric presented below.

See below for the discussion post rubric.*

*Post and Response Rubric__

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Textbook assignment
  1. 31, 32 42, 43, 44, 45 pp 26-27
Funsie #1
Quiz #1


external image water.jpgElectrons in Atoms
Section #2 Unit #2

Light and Quantized Energy

Wave nature of light (Electromagnetic Spectrum)

Frequency (ν)

Wavelength (λ)


Visible Light 400- 750 nm. Which is red? Which is violet?

c = Speed of Light = 2.998 x 108 m/sec = 2.998 x 1017 nm/sec

c =λ x ν



photoelectric effect

The energy of each photon of light can be calculated by:

E = hν or E = hc/λ Where h is Planck’s constant.

h = 6.626 x 10-34 J.s/photon

atomic emission spectrum

Classwork: YCP worksheet Properties of Light

Lab: Flame Tests

Read and Post : Article on Motion Detectors

Textbook assignment : Read Section 5.1 Do #’s 5-7 p. 143, #13, 14 p. 145
#35, 44, 47-51, 55, 58 p. 166

Quantum Theory and the Atom

Bohr Model

Ground state vs Excited state:

Quantum Mechanical Model of the Atom.

The modern day model of the atom uses the concepts of quantum theory to best describe the behavior of electrons. We should not perceive of electrons as orbiting the nucleus like planets around the sun. They behave by a different set of rules.

deBroglie equation:

Heisenburg uncertainty principle

Schrodinger wave equation

To describe thProxy-Connection: keep-alive Cache-Control: max-age=0 energy, behavior and probable location of an electron we assign each one a unique set of 4 quantum numbers. These are designated with the symbols n, l, ml and ms
(See below).

Principal quantum number (n) This is the number of the level (ring) of electrons starting from the nucleus and going outward!

n = 1,2,3,4…!

[2n^2= number of electrons on that level!]


Second (azimuthal) quantum number (l) This designates the shape of the probability cloud!

l = 0, 1, 2…(n-1)!
These correspond to the shapes of the probability distributions(wave functions or orbitals). The first 4 l values are often designated s,p,d,f!

If... n=1 then l=0, n=2 then l=1 etc..!
s= sphere shape orbital, p= dumbell shape orbital, etc..!

Third (magnetic ) quantum number, ml. Used to designate the orientation of the orbitals in space.

ml = -l…0…+l

s is always 0!
p could be -1,0, Or +1!
so on and so forth!



Fourth quantum number, ms. Used to designate the spin of the electron.

ms = +1/2 or - 1/2

Classwork: Worksheets as assigned

Textbook Assignement: Read Section 5.2 Do # 15 p. 155; # 63, 67, 68, 70, 71, 72 p. 167; # 94, 97, 105, 106 p. 168

Electron Configurations, Orbital Diagrams and Abbreviated Electron Configurations

Electrons fill the energy levels from the bottom up.

pt04a.jpg :// FptsNFm17i6jD1ws1-8jFYfGEHE=&h=248&w=410&sz=90&hl=en&start=0&zoom=1&tbnid=9ydiMjmpe6_WEM:&tbnh=122&tbnw=201&ei=Qe9wTZuMC4P6lwecptHKDQ&prev=/images%3Fq%3Ds%2Bblock%2Bp%2Bblock%2Bd%2Bblock%26um%3D1%26hl%3Den%26safe%3Dactive%26client%3Dsafari%26rls%3Den%26biw%3D1267%26bih%3D611%26tbs%3Disch:1&um=1&itbs=1&iact=hc&vpx=774&vpy=105&dur=2329&hovh=174&hovw=289&tx=109&ty=76&oei=Qe9wTZuMC4P6lwecptHKDQ&page=1&ndsp=17&ved=1t:429,r:3,s:0

Aufbau principle
-"Electrons fill from lowest energy levels to the highest energy levels." Quoth Mr. Williams.
-There are exceptions: as the energy levels get higher, they can start to overlap.

Pauli exclusion principle
-No two electrons in an atom can have the same quantum numbers.
-Both electrons spin in opposite directions in each orbital (if two are even present in the orbital).

Hund’s rule
-Maximize the number of unpaired electrons (Fill up each orbital in an energy level before pairing them).
Electron-dot structure
-Shows number of valence (outer shell) electrons. Valence electron pairs are represented with dots around the atomic symbol.

Valence electrons
-Valence electrons are the total electrons in the highest level.

-Oxygen has 6 valence electrons.

Fluorine and Chlorine have the same amount of valence electrons because they are in the same group/ column in the periodic table

D's run one behind the period number
F's run two behind the period number
-ex. 4s 3d 4p
6s 4f 5d 6p

Exceptions to Normal Electron Configurations

Filled and half filled sublevels are exceptionally stable.

Electron Arrangement in Ions

In transition metals, "s" electrons are lost first

Cations= Positive
Anions= Negative

Classwork: Worksheets as assigned.

Textbook Assignment: Read Section 5.3 Do #’s 21-24 p. 160; # 29, 30, 33 p. 162; # 81, 86, 87, 88, 112-114 pp 167-169.

The Periodic Table and Periodic Law


Groups or family: Chemistry a set of elements occupying a column in the periodic table and having broadly similar properties arising from their similar electronic structure.

Ex: Alkali metals : any of the elements lithium, sodium, potassium, rubidium, cesium, and francium, occupying Group IA (1) of the periodic table. They are very reactive, electropositive, monovalent metals forming strongly alkaline hydroxides.

Alkaline Earth Metals: any of the elements beryllium, magnesium, calcium, strontium, barium, and radium, occupying Group IIA (2) of the periodic table. They are reactive, electropositive, divalent metals, and form basic oxides that react with water to form comparatively insoluble hydroxides.
Transition metals any of the set of metallic elements occupying a central block (Groups IVB–VIII, IB, and IIB, or 4–12) in the periodic table, e.g., iron, manganese, chromium, and copper. Chemically they show variable valence and a strong tendency to form coordination compounds, and many of their compounds are colored.

Halogens: any of the elements fluorine, chlorine, bromine, iodine, and astatine, occupying group VIIA (17) of the periodic table. They are reactive nonmetallic elements that form strongly acidic compounds with hydrogen, from which simple salts can be made.

Noble or inert gases: any of the gaseous elements helium, neon, argon, krypton, xenon, and radon, occupying Group 0 (18) of the periodic table. They were long believed to be totally unreactive but compounds of xenon, krypton, and radon are now known.

Metalloids : an element (e.g., arsenic, antimony, or tin) whose properties are intermediate between those of metals and solid nonmetals or semiconductors.

Boron, Silicon, Germanium, Arsenic, Antimony and Tellurium.

Periods: a set of elements occupying an entire horizontal row in the periodic table. #’s 1- 7.

Rare Earth elements: any of a group of chemically similar metallic elements comprising the lanthanide series and (usually) scandium and yttrium. They are not esp. rare, but they tend to occur together in nature and are difficult to separate from one another.


Textbook Assignment: Read Section 6.1 and 6.2 Do #’s 2, 4, 5 p. 181; #12, 13 p. 186; # 29, 30, 31, 32, 33 35, 36, 40, 43, 46, 47, 48 pp198-199.

Trends in the Periodic Table

Increasing Core Charge
-Core Charge ( Effective Nuclear Charge)= # of protons - the # of core electrons LS
-An increase in positive core charge causes a decrease in the size of the atom as the electrons are drawn closer to the nucleus.

Increased Shielding/Screening
-The core electrons push the valence electrons out, so the atoms increase in size LS
-High and tight in the upper right
-Large and loose in the lower left

Atomic radius:
-the radius, or distance from the center to the end, of an atom
-measured in picometers (pm)

Ionic Radius:
- When electrons are added the ion increases in size
- When electrons are stripped the ion decreases in size
-ionic radius is also measured in picometers (pm)

Ionization Energy:
-Energy required to remove an electron
1st I.E.
2nd I.E.
3rd I.E


Electronegativity: (EN)

- Measure of the amount of attraction an atom has for electrons in a chemical bond

- Highest is F (4.0)
-Lowest Cs (.7)

LS and VP

Textbook Assignment : Read Section 6.3 Do #’s 21,22, 23 p. 194; Do #’s 59-64 p. 199; # 78, 89, 94, 95 pp. 200-201.


Ion Formation

Cation-A positively charged ion. It forms when an atom loses one or more valence electrons

Metal ions- lose valence electrons easily
Group I, II, III (1, 2, 13)
Group 1: turns into 1+ when s1 is lost
Group 2: turns into 2+ when s2 is lost
Group 3: turns into 3+ when s3 is lost

Transition Metal Ions(d block)- Form 2+ ions. May cause 3+ ions or greater

Anion- An anion is a negatively charged ion. Anions end in -ide

Non-metal Ions- Nonmetals gain the number of electrons that when added to their valence electrons equals 8, so they can become a stable octet.

Group V, VI, VII (15, 16, 17)
Group 15: Gain three electrons to achieve an octet
Group 16: Gain two electrons to achieve an octet
Group 17: Gain one electrons to achieve an octet

Textbook Assignment: Read Section 7.1 Do #’s 5, 6 p. 209 # 48, 49, 52-55 p. 232

Ionic bonds- electrostatic force that holds oppositely charged particles together in an ionic compound
Ionic compounds are metal/non-metal

Ex. NaCl- Binary compound (Contain only two different elements)

Note: metal-nonmetal and large difference in electronegativity.

Physical Structure: Large numbers of positive and negative ions exist in a ratio together
  • Determined by the number of electrons transferred from the metal atom to the nonmetal

Physical Properties:
-Melting point, boiling point, and hardness.

Electrolyte - Ionic compound whose aqueous solution conducts an electric current

Lattice Energy Energy required to separate 1 mol of the ions of an ionic compound.

Crystal Lattice 3 dimensional geometric arrangement of particles each positive ion is surrounded by negative ions and vice versa. It varies in shape due to the sizes and relative numbers of the ions bonded.

What affects the strength?- Lattice energy is related to the size of the ions bonded. Smaller ions have greater lattice energy
-Value is affected by the charge of the ion
*Larger positive/negative charged have greater lattice energy
KING: Al +3 O -2= Al2 O3
Textbook Assignment: Read Section 7.2 Do #’s 16, 17 p. 217; # 63, 65, 67, 70, 71, 73, 74, 75 p. 233 # 100, 101 p. 234

Using the Periodic Table to Assign Ionic Charges(Oxidation Numbers)

external image clip_image002.png

Naming and Writing Ionic Formulas

Binary Ionic 1 metal and 1 nonmetal

Always cation(typically a metal); anion (typically a non-metal) ending in –ide.
Know when to go Roman!

Ex. NaCl - sodium chloride

CaF2 - calcium fluoride

AlBr­3 - Aluminum bromide

Magnesium chloride Mg2+ Cl 1- MgCl 2
Potassium oxide K1+ O 2- K2O
Calcium sulfide Ca2+ S2- CaS

Polyatomic Ions: An ion made up two or more atoms bonded together that acts as a single unit with a net charge

Ionic compounds that contain polyatomic ions_

"Sulfur called carbon crazy" : -2

-ite NItrite NO2 -1
-ate Nitrate NO3 -1
If skinny "i" ATE an O, it gets "a"

Hydrox Cookies minus the top biscuit = Hydoxide!

ClO ^- Hypochlorite (Hypo stands for UNDER)
ClO2^- Chlorite
ClO3^- Chlorate
ClO4^- Perchlorate (Per is OVER like a Periscope on a submarine)

NOOO ^1- (You'll cost me a dollar unless you call me at the NIGHT RATE!!)

Code 1
CHO232 : Officer Acetate got to the scene and found one he has to report that there's been a population change of CHO 232^-1 !

Ex. CuSO4 copper(II) sulfate

CuSO3 copper(II) sulfite

Magnesium nitrate Mg2+ NO31-


Ammonium phosphate NH4 1+ PO4 3-


Textbook Assignment: Read Section 7.3 Do #’s 36, 37, 38, 39 p. 224; 81-86 p. 233; #102, 111 p. 234 + worksheets as assigned.

The noble metals : metals that never fade, such as gold or platinum

Metallic bonding - based on electronic structure, the attraction between the protons of the nucleus and the electrons flowing between them (in the electron sea model) is what causes the bond

Electron sea model - the electrons are loosely held with the matrix of metallic ions and can flow within like a sea
external image mtlbond.gif

Delocalized electrons - The electrons flow, they are free to move, and are out of their location (hence why its called delocalized).

Ex. Na


Alloys - a mixture of different elements that have metallic properties

Ex. Steel - combination of carbon and iron
Brass - Copper and Zinc
Bronze - Iron and copper (and some tin)

Textbook Assignment: Read Section 7.4 Do # 44 p. 228; # 87, 91, 96, 98, 113 pp. 233-234; #


Covalent bond: A chemical bond that results sharing valence electrons

Molecule: Formed when two or more atoms bond covalently

Molecules/Molecular compounds: Two or more non-metals chemically combined.
Molecular formulas(examples)

Diatomic Elements: Elements that come in pairs, covalently bonded together. (There are 7. Hydrogen, oxygen, nitrogen, chlorine, bromine, iodine, and fluorine)
  • HONClBrIF or ClIF BrOHN (Which apparently sounds somewhat like James Bond...)
  • 6 of the 7 diatomic elements form the shape of a 7 and points to hydrogen

The Lewis structures show the different numbers of bonds. The maximum number is three bonds (no quadruple bonds or anything).

As you can see from the picture, these two nitrogen atoms have covalently bonded together, because they need to get stable. They both now "possess" eight valence electrons, and according to the octet rule, it means they are stable.

Below are examples of common covalent bonds. *One of these are MOLECULES. Molecules are only associtated
H­­2O Water
C2H6O – EthylAlcohol ( Ethanol)
C5H12 - Pentane
Structural Formulas
external image clip_image002.png
Condensed Molecular Formulas
Ex.CH3(CH2)3CH3- pentane

Textbook Assignment #1 Read Section 8.1. Do #’s 1-4 p. 244, # 10, 13 p. 247; # 78, 83, 84, 86 p. 274.

Naming Binary Molecular Compounds

Binary Molecular 2 non-metals

Use Prefixes :
1- mono
2- di
3 –tri
4- tetra
5- penta
7- hepta

8 - octa
9- nona
10- deca

CO- carbon monoxide
PCl3- phosphorus trichloride
N2O- dinitrogen monoxide

Acids :Molecular compounds that contain an ionizable hydrogen ion(H+)

HX- binary acidscontain H and one other element.

HXO- oxy acidscontain H, oxygen and one other element.


Hydro-root word -icacid

HCl---------hydrochloric acid
HBr---------hydrobromic acid
H2S--------hydrosulfuric acid
HI-----------hydroiodic acid
Oxy acids!!!!!!
“I took a bite and became nauseous.”
“I ate it and became sick.”

H2SO4 sulfuric acid- ate>>> -ic
HNO3 nitric acid
H3PO4phosphoric acid
HNO2nitrous acid-ite>>>>-ous
HClOhypochlorous acid
H2SO3sulfurous acid
Textbook Assignment #2 Read Section 8.2 Do #’s 19-30, p. 251 # 35,36, p. 252, #92-96, p. 274

Molecular Structures and Intermolecular Attractions

Review of Ionic and covalent bonding via an animation

Lewis Structures and IM Forces

Single, Double and Triple Covalent Bonds

Generally occurs between 2 or more non-metal atoms which share valence electrons in order to achieve a Noble gas configuration. (usually 8 electrons).Hence the Octet Rule.

Primo Example : HONClBrIF

external image diatomic.gif
Bonding Tendencies
Carbon (C) tends to bond 4 times.
Nitrogen (N) tends to bond 3 times.
Oxygen (O) tends to bond 2 times.
Halogens (F, Cl, Br, I) bond 1 time.
Hydrogen (H) Bonds 1 time ONLY!!!!!

Lewis Structures

Rules p 254

1. Draw a "skeleton" attaching all atoms with single covalent bonds.
2. Count all valence electrons.
3. Subtract 2 electrons for each single covalent bond in the skeleton.
4. Distribute remaining electrons as non-bonding pairs. Look for double and triple bonds.

Ex. NH3








Expanded Octets

Primo Example:

Non-metal Halides



Molecular Geometry and Bond Angles

VSEPR- Valence
Repulsion Theory

Central atom with 4 attachments.

Tetrahedral- A central atom with four bonding atoms on the corners of the central atom


Central atom with 3 attachments and 1 nonbonding pair.



Central atom with 2 attachments and 2 nonbonding pairs.



Central atom with 3 attachments.

Trigonal Planar


Central atom with 2 attachments.




Expanded Octet Shapes



Textbook Assignment #3 : Read Section 8.3, 8.4 Do #’s 44-46 p. 258; #’s 47,48,52, 54 p. 260; # 63, 64, 67(skip hybrid orbitals) p. 264, # 101, 102, 108, 111, 112(skip hybrid orbitals), 126 pp 274-276.

Molecular Polarity

A molecule is considered polar if it has a permanent positive region and a permanent negative region.Molecules that are polar are often called dipoles.

In order to be polar a molecule must have polar bonds(eg ∆ EN > 0.3 but < 1.7) and be asymmetrical. That is the effect of the unequal sharing of electrons can’t be cancelled by symmetry.

-Symmetry cancels polarity when the atoms are the same... Different atoms with different electronegativity do not cancel polarity, even if the molecule is symmetrical.

View this site and the animation found there for further explanations:
external image water.jpg

-Water is very polar.... Oxygen's electronegativity is greater than hydrogen's.

-There must be a difference in electronegativity between the atoms.

General rule- bent and pyrimidal molecules are generally polar.

Intermolecular Attractions

These are attractions between molecules. are weaker than the intramolecular attractions better know as covalent bonds that are within a molecule.


1.) Dispersion forces (London) or induced dipole.

2.) Dipole-dipole forces

3.) Hydrogen Bonds (The Mighty Mice)

Textbook Assignment # 4 Read Section 8.5 Do #’s 74, 75, 76, 114, 118, 119, 120, 121, 127, 128 138-140

Lewis Structure Exceptions:

1.) Electron Deficient Molecule
2.) Odd # of Electrons
3.) More than and Octet

"Watch out for BBe'sexternal image YellowJacket.gif
B- Boron
Be- Berrylium


external image
external image IMG00001.GIF

Resonance: The moving of electrons to other location
-Increases Stability

The story of Benzene:
Kekulé was half asleep when he dreamt of a snake biting its own tail, creating the hexagon shape instead of the chain.

external image Benz1.png Benzene lewis structure (C6H6) external image 20051120170851Heinrich_von_Angeli_-_Friedrich_August_Kekul%25C3%25A9_von_Stradonitz.jpg Kekulé

Intermolecular Attraction

-Weaker than intramolecular attractions AKA covalent bonds

Dispersion forces (London) or induced dipole
-caused by fluctuations in the electron field
- These occur between all molecules, but are the only force between non-polar molecules

-Can shift to interfere with other clouds around it
- Form partial partial positives, and partial partial negatives
- Dispersion forces are greater in larger molecules


Dipole-dipole forces
-The positive region of one polar molecule attracts the negative region of another polar molecule

Hydrogen Bonds
-If a tiny H is attached to F, O, or N in a molecule, huge polarities are created.
-Therefore, very strong H-bonds form between the molecules.


note: DO NOT confuse with hydrogen bombs

V.P. and L.S.

Balancing Chemical Equations

You must have an equal number of each type of atom on either side of the equation.
The equation must obey the Law of Conservation of Mass.
Law of Conservation of Mass: Matter is neither created or destroyed.

Use coefficients to balance equations. Never change subscripts!

Reactants ---Yield---> Products

H2 +O2 -> H2O

Δ - Heat is applied to the reactants.
↓ - A precipitate form.
↑ - A gas forms.
(l) - A chemical is in the liquid state.
(aq) - A chemical is dissolved in water.
⇄ - A reversible reaction occurs.

  • Save H then O till last.
  • Eliminate fractions or decimals by multiplication.
  • CO3's almost always become CO2's in reactions

Magnesium reacts with oxygen gas to form magnesium oxide.

2Mg(s) + O2(g) → 2 MgO(s)

Physical state symbols: gas(g) liquid(l) solid(s) aqueous(aq)

Watch out for Diatomic elements (elements that occur in pairs)

The Magic 7: H2, O2, N2, Cl2, Br2, I2, F2

Honclbrif HONClBrIF

Ever play Battleship?
Try this version:

Read Section 9.1 Do #’s 1-3 p. 284. 4-6, p. 287, 12, 13p. 288, 62, 64-66, 68, 71-75 p. 312.

Precipitate: Anything that falls out of a solution

Types of Chemical Reactions

Synthesis (also known as Direct Combination) : Two or more reactants combine to form one.
A + B → AB

Decomposition: One reactant breaks down into two or more products.
AB → A + B
(Normally requires heat)

Single Replacement: With one element and a compound, the element (if it has the power) can replace part of the compound.
A + BX → AX + B

Double Replacement: When two ionic compounds dissolved in water, the front two elements switch to make two new compounds
XY + AB → AY + XB

Combustion: When the reactant is exposed to oxygen (through burning)
CxHy + O2 → CO2 + H20

Think you know your stuff?

Try this quiz :

Read Section 9.2 plus pp 299-302 Do #’s 14-17 p. 291, 18-19 p. 292, 21-24, p. 295, 25- 28, p. 297, # 35- 39, p. 302 – Just chemical equations, 80, 81, 82, 85, 86-88, 100, 107, 112, 126, 128, 130, 133 pp 312-315.

The Mole - Its like a dozen only bigger

A mole is an an animal that burrows in the ground,
Or the spot on your chin that you gotta shave around.
But there's another kind of mole of interest to me,
That's the kind of mole they use in chemistry.

Atomic Mass and the Mole - molar mass calculator

Atomic Mass: The average of all naturally occurring isotopes.
(Recall Isotopic Calculations)

Ex. 1 atom Na = 22.98977 amu’s

C – 12 has a mass of exactly 12.000000…. amu’s All other atoms are compared to C-12. In fact 1 amu is defined to be 1/12 of a C- 12 atom.

Avogadro’s Number and the Mole
Mole Examples/Jokes
1 mol = 6.02 X 10 23
602,000,000,000,000,000,000,000,000 602 sextillion

Suppose a mole of marshmallows fell upon the planet,
Over each square inch of land and sea, think that you could stand
it?That layer would be twelve miles high and of course block out the sun,
We're talking close to five million trillion tons.

Celebrate Mole Day 6:02 am – 6:02 pm 10/23
external image polo98.gif

Textbook Assignment #1 Read Section 10.1 Do #’s 1-4 p. 323, # 12 p. 324, # 92-94, 99 p. 358,

Mole ← → Gram conversions


1 mol Na atoms = 6.02 x 10 23 atoms
1 mol Na Atoms = 22.98977 grams

1 mol equals an atom’s atomic mass in grams.

Si – 28 27.977 amu’s = 1 atom.

10.0 g --→ moles?

10.0 g x 1 mol/ 27.977 g = 0.357 mols Si x 6.02 x 10 23 atoms/ 1 mol = 2.15 x 10 23 atoms.

A mole is similar to a dozen – For example:

.333 dozen donuts .333 doz x 12 donuts/dozen = 4 donuts

Still confused? Check out this explanation:

Textbook Assignment #2 Read Sections 10.2, 10.3 Do #’s 19 a, c, 20 a, c, e, p. 331, #28 p. 332, #’s 43, 46, p. 339, # 103, 104, 108, 109, 112, 115, 117, 118, 128, 131, 132, 135, 136, 145, 148, 150 pp 358-360.

Percent Composition from Formulas

Cinnabar HgS mercury ore

% of Hg in HgS 200.6 g Hg + 32.1 g S = 232.7 g total

200.6g Hg/ 232.7 total g x 100 = 86.2 %
13.8 % S

Milk of Magnesia Mg(OH)2

% of Mg in Mg (OH)2

24.3 g Mg % Mg = 24.3g Mg/58.3 total g x 100 = 41.7 % Mg
32.0 g O
2.0 g H
total = 58.3 g

Simplest Formula from Chemical Analysis

Hydrate example!!!!!!!!!!!!!!!!!!!!!!!!:) :) :)

Na2CrO4 . X H2O is experimentally determined to contain 32. 1 % water. Find X.

X = mols H2O/mols Na2CrO4

Assume 100 g of hydrate

32.1 g H2O/ 18.0 g/mol = 1.78 mols

67.9 g Na2CrO4/ 162 g/mol = .419 mols

X= 1.78 mols/.419 mols = 4.25 ⇒ 4 = X=4

Molecular vs Empirical formulas

A molecular compound contains 92.3% carbon and 7.7% hydrogen by mass.

a.) Find its empirical(simplest) formula. The simplest whole # ratio of atoms in a compound.

Assume 100g of compound

92.3 g C/12.0 g/mol = 7.69 mols C

7.7 g H/1.0 g/mol = 7.7 mols H

7.7:7.69 ⇒ 1:1 So emp. Formula is CH

b.) If 0.050 mol of this compound has a mass of 3.90 g, what is its molecular formula?

Emp. Formula is CH Molecular formula could be CH,C2H2,C3H3,C4H4 etc….

Molar mass = 3.90g/.050mol = 78 g/mol

CH???? = 78 g/mol Noooo!!!! It equals 13 g/mol
C2H2 ???? = 78 g/mol Noooo!!!! It equals 26 g/mol

Divide 78 by 13 78/13 = 6 so the molecular formula is C6H6 yessss!!!!

Textbook Assignment # 3 Read Sections 10.4, 10.5 Do #’s 68, 69, 73 p. 350, #’s 159, 161, 64, 166, 172, 173, 173, 175, 179, 180, 181, 185, 187, 194, 196, 204, 105, 211 pp 361-363.

Online Quizzes


Try a few on empirical formula, % composition, Avogadro’s number and the mole.

Colorful letters made by GRK! ™®©

Reaction Stoichiometry
Using Balanced Equations to Make Predictions
Mol-mol relationships
Example: View the video found here: mol to mol conversion

Mol-mass relationships
Example: View the video found here: mass to mols

Mass-mass relationships
Example: View the video found here: Mass to mass calculation

Textbook Assignment # 1 Read Section 11.1-11.2, Do #2 p. 371; #9, p. 372; # 12 p. 375; # 14 p. 376; # 16 p. 377; # 21, 22 p. 378; # 50, 51, 53, 59, 60, 62, 64, 66, 69, 70 pp 392-394.

Limiting Reactants and Theoretical Yield/% Yield

Support: Go to the following website for examples and explanations: Limiting and Excess Reagents

Example: View the video found here: Limiting reagent. Finding theoretical yield and the dreaded amount leftover!

Quizzes: Try these on for size: Matching quiz to see if you know the terminology.

Multiple Choice

Extra Quizzes: Try the following here: Stoichiometry - Mole to Mole Problems, Mole-mass Stoichiometry Problems, Mass-mass Stochiometry Problems, Limiting Reagent, Percent Yield, Percent Yield with Limiting Reagents

Textbook assignment # 2 Read section 11.3-11.4 Do #’s 23, 24 p. 383; # 27 p. 384; # 34, 35 p. 388; # 75, 77, 79, 81, 89, 90, 94, 103, 109, 110, 111a, c, 113 pp 394-397.



Check out the animation: Gases under pressure.
How is a gas different from a solid or liquid? List characteristics of gases.
-They don't stick together; non-bonding. (Only intermolecular forces)
-They are in constant, random motion.
-Greatest motion within particles.
-Usually non-polar. (No dipole-dipole forces!!!)
-Low boiling point.
-Low density.
-No defined shape.
-Can be compressed or expanded. Solids and liquids can't do this!
-Can be diffused through each other.

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Check out this video: è

Measurements of Gases

4 Variables that affect gases

Pressure(P)- force/area... measured in psi (pounds per square inch)- sea level 14.7 psi, 14.7 psi= 101.3 K Pa (Kilo-pascals)= 1.013 bars= 1 atm (atmosphere)
-Pressure decreases as you move farther above sea level.

Torricelli- invented barometer. X mm of Hg = X torr

Volume(V)- usually expressed in liters

Amount(n)- number of moles

Temperature(T)- always expressed in Kelvin (K)... K= degrees Celsius+273.15

Kinetic Theory

Tenets p 403

K.E.= 1/2 mass X velocity^2

external image image_275637.jpg<----- Balloon at normal temperature.

external image 2154_picture_of_a_sad_bear_holding_deflated_balloons.jpg <------- Deflated balloon after days of air pressure loss... And a very sad bear.

external image gas_particles.gifGas particles are not attracted to each other. They also bounce forever and continue to move at all times.
external image images?q=tbn:ANd9GcRMR3MpvZmS5x3OcujQGAmR2pUpRdX57aQZUVCaK9VEAnGCFZmG&t=1Particle collisions are "perfectly elastic".
Dalton’s Law of Partial Pressures
external image images?q=tbn:ANd9GcT1RK2--TC3xVsJOPmEyzMdDZPWes9amXPYIk1RZvz3MZnw2lez&t=1 Which equals 740 Torr.
Ptot = P1 + P2 + P3 +…..
PA = XA x Ptot Where XA = the mol fraction of A

Subtracting vapor pressure of water when finding the pressure of a gas collected over water.

Textbook Assignment #1 Read Section 12.1 Do #’s 4, 5, 7 p. 409, # 39, 45, 47, 49, 50 p. 434.

Boyle’s Law:

P1V1 = P2V2

Sketch graph

Charles Law:

V1/T1 = V2/T2

Sketch graph:

Avogadro’s Law

V1/n1 = V2/n2

Sketch graph:

Gay- Lussac’s Law

P1/T1 = P2/T2

Sketch graph

The Combined gas law and the value for “R” for an ideal gas.

VP, LS and KO

Textbook Assignment # 2 Read Section 13.1 Do #’s 1, 2, 3 p. 443, 4, 6 p. 446, 8, 10 p. 448, 11, 12, 13 p. 450, # 18 p. 451, # 55, 57, 58, 60 p. 468.


PV = nRT
One variable calculations.

Using Ideal gas law to find MM and Density.

Textbook Assignment #3: Read Section 13.2 Do #’s 20, 21, 24, 25 p. 453, #’s 26, 28, 30, p. 455, #36 p. 459, # 65, 67, 68, 70, 72, 75, 76, 93, 100 pp 468-470.

Relevant Reading Article on SF6

Demos: Tank of Doom/Flames of Gehenna

Stoichiometry of gaseous reactions


Problems can be solved in the same way as mol-mol problems IF temperature and pressure are constant.

Mass – Volume
Use molar volume (22.4 L/mol) if the conditions are at STP.

If not use ideal gas law to find n (mols) and then mass (g).

Textbook Assignment # 4: Read Section 13.3 Do #’s 38, 41 p. 461, #’s 42, 44, 45 p. 463, # 82, 84, 87, 89, 90, 103, 108, 110 b, d, f, 112, pp 469-471