CHEMISTRY: SYLLABUS CHOICE 1

CHEMISTRY COURSE CHOICE #1

Use DVD Classes from the Chemistry DVD's from Netflix.

The plan:

1. Watch the 1st DVD and fill out TEST PACKET #1 as you listen.
2. Study Test Packet #1 for test.
3. Become acquainted with what a Chemical Element is,
4. The Periodic Table

**still under construction**
5.Complete the 2nd DVD Test Packet
6. Complete the 3rd DVD Test Packet

Accomplish a basic understanding.


Plan a few experiments for later weeks.

CHEMISTRY: Solutions, Colloids

SOURCE

Solutions and mixtures

Types of Mixtures

Homogeneous mixtures are uniform throughout. The composition is the same in all directions in the substance. The types of particles observed in one direction are the same that are observed in all others. A solution of sugar in water is homogeneous.



Heterogeneous mixtures are not uniform. There are pockets of one substance surrounded by pockets of different substances. A mixture of soil in water to make "mud" is heterogeneous. Likewise a mixture of oil and water in salad dressing is heterogeneous.



Types of Homogeneous Mixtures

Solutions

Particle sizes distinguish one homogeneous mixture from another. Solutions are mixtures with particle sizes at the molecule or ion level. The particles have dimensions between 0.1 to 2 nanometers. Typically solutions are transparent. Light can usually pass through the solution. If the solute is able to absorb visible light then the solution will have a color. A blue liquid transmits blue light and absorbs the other colors of the spectrum.

A mixture of water H2O and ethanol CH3CH2OH is homogeneous. The particles are individual molecules of H2O and CH3CH2OH. The two molecules are spread uniformly throughout the solution.

A mixture of water and sodium chloride is homogeneous by chemistry standards. The particles in the mixture are molecules of H2O and hydrated sodium cations, Na+, and chloride anions, Cl1-.

Solutions are transparent. You can see through them. The mixture remains stable and does not separate after standing for any period of time. The particles are so small they cannot be separated by normal filtration.

A solution may have a "color" but it will still be transparent.

Colloids

Colloids are mixtures with particle sizes that consist of clumps of molecules. The particles have dimensions between 2 to 1000 nanometers.

The colloid looks homogeneous to the naked eye. Fog and milk are examples of colloids.

Colloids frequently appear "murky" or "opaque". The particles are large enough to scatter light. You have experience with the way fog interacts with the light from car headlights. Colloids generally do not separate on standing. They are not separated by filtration.


Suspensions

Suspensions are homogeneous mixtures with particles that have diameters greater than 1000 nm, 0.000001 meter. The size of the particles is great enough so they are visible to the naked eye. Blood and aerosol sprays are examples of suspensions.

Suspensions are "murky" or "opaque". They do not transmit light. Suspensions separate on standing. The mixture of particles can be separated by filtration.



Examples of solutions

Solutions are mechanical combinations of materials. The physical state for the materials in the solutions is surprising varied. The combinations of states of matter that form solutions are listed below.

States of matter in solution
Example

gas in gas
air ( N2, O2 , Ar, CO2 , other gases)

gas in liquid
soda pop (CO2 in water)

liquid in liquid
gasoline (a mixture of hydrocarbon compounds)

solid in liquid
sea water ( NaCl and other salts in water)

gas in solid
H2 in platinum or palladium

liquid in solid
dental amalgams ( mercury in silver)

solid in solid
alloys ( brass, (Cu/Zn), solder (Sn/Pb)

Online Introductory Chemistry

Dr. Walt Volland all rights reserved revised March 29, 2005

CHEMISTRY SYLLABUS

NOTE TO SELF:

HERE ARE THREE SYLLABUS CHOICES.

PICK ONE AND RUN WITH IT OR MIX AND MATCH TO MAKE YOUR OWN.

THIS YEAR I LIKED SYLLABUS 3 AS THERE WERE A FEW TOPICS TO RESEARCH AND KAZZ WAS ABLE TO GO IN DEPTH OR SKIM AS HE PLEASED WITH THOSE TOPICS USING GOOGLE.

-------------------------------------------------------------------------------------------------------------

SYLLABUS CHOICE 1

Standard Deviants 3 DVD's

Fun and informative.

TEST PACKET FOR FIRST DVD

-----------------------------------------------------------

SYLLABUS CHOICE 2

SOURCE

Experiments

------------------------------------------------------------

SYLLABUS CHOICE 3

SOURCE

*still under construction

CLICK ON THE WORDS FOR THE LESSON

(If it isn't "clickable" use the internet/google for

information"

Measurements and Calculations

Intro. to Chemistry and Matter

Atomic Structure
Electron Configuration
Periodicity
Chemical Bonding
Molecular Shapes
Polarity

Second Quarter:

Nomenclature / Formulas & Composition
Reactions Equations
Stoichiometry

Third Quarter:

Kinetic Molecular Theory
The Gas Laws
Gas Stoichiometry
Liquids and Solids
Solutions, Colloids
Colligative Properties
Chemical Equilibrium
Acids, Bases

Fourth Quarter:

Acid-Base Titrations and pH
Reaction Rates Energy 17
Oxidation & Reduction, Electrochemistry
Nuclear Chemistry 22
Organic Chemistry: Carbon, Organics Compounds

Lab Activities

CHEMISTRY: Atomic Structure

Early Atomic Theory
-a little bit of history


Democritus (400 BC)


Thought that the world was made of empty space and tiny particles which he called "atoms." He proposed "atoms" as a concept but had no experimental proof.

Aristotle (384-322 B.C.)


Proposed that matter was continuous and not made up of smaller particles. He called this continuous substance "hyle." This belief was held till the 17th century.

Isaac Newton and Robert Boyle (1627-1691)(17th century)


Proponents of the idea of atoms, but could offer no proof

John Dalton (1766-1844)


Studied the observations of others concerning chemical reactions

Antoine Lavoisier (1743-1794)


Found that when a chemical change occurred in a closed system, the mass after the reaction equaled the mass before the reaction. He proposed that in ordinary chemical reactions, matter can be changed in many ways, but it cannot be created or destroyed. This became known as The Law of Conservation of Mass.

Joseph Proust (1754-1826)


Observed that specific substances always contain elements in the same ratio by mass. This became known as The Law of Definite Proportions.

Dalton's Atomic Theory

Formed the basis of our present atomic theory while trying to explain Lavoisier and Proust.


Dalton's Atomic Theory
1. All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass and other properties. Atoms of different elements differ in size, mass and other properties.
3. Atoms cannot be subdivided, created or destroyed.
4. Atoms of different elements can combine in simple, whole-number ratios to form chemical compounds.
5. In chemical reactions, atoms are combined, separated, or rearranged.


Dalton's ideas explained the "Law of Conservation of Mass" and the "Law of Definite Proportions."


Dalton's ideas are not entirely correct, but they did form the basis of our current thinking.





Dalton's Law of Multiple Proportions


The ratio of mass of one element that combine with a constant mass of another element can be expressed in small whole numbers. (SO2, SO3; NO, NO2; SnO, SnO2)


J.L. Gay-Lussac (1778-1850) observed that gas reactions (under conditions of constant temperature and constant pressure) the volumes of reacting gases and gaseous products are in the ratio of small whole numbers.





Avogadro's (1776-1856) Hypothesis (see Ch. 11)

Equal volumes of gases, under the same conditions, have the same number of molecules





Structure of the Atom

Early Research on Atomic Particles

In the mid-1800s, experiments showed that the atom is made up of smaller particles


Discovery of the Electron

J.J. Thomson (1897)


Worked with the cathode ray tube and is credited with the discovery of the electron.

Robert Millikan (1909)


Measured the charge of an electron using his "oil drop apparatus"

Discovery of the Nucleus

Lord Rutherford (1912-13) Gold Foil Experiment


In the work he conducted with Hans Geiger and Ernest Marsden, Ernest Rutherford bombarded a very thin sheet of gold foil with a stream of positively-charged subatomic particles (alpha particles). A few of the particles "bounced back."

He concluded that most of the atom is empty space. He proposed that all the mass was concentrated in the center of the atom and this center was positively charged. This center we now call the NUCLEUS.

E. Goldstein


Discovered positively-charged particles (protons) using a modified cathode ray tube.

James Chadwick (1932)


Discovered the neutron, a particle with the same mass as a proton but having no charge.

The discovery of these subatomic particles led to a major revision of Dalton's hypothesis.


Composition of the Nucleus

The nucleus consists of protons and neutrons and has an overall positive charge which is balanced by the negative charge of the electrons. The mass of a proton is 1836 times the mass of an electron. A neutron has slightly more mass that the proton.


The number of protons determines the identity of the atom. The number of protons and neutrons combined determines the mass.


Isotopes


Number of protons in an atom determines its identity.


Number of electrons must equal the number of protons in a neutral atom. The electrons determine an atom's chemical reactivity (how it behaves).


Atoms of the same element can differ in mass (same number of protons but a different number of neutrons). The number of neutrons (plus the number of protons) determines the mass of an atom. These different forms of the atoms are called ISOTOPES.


Moseley


Worked with X-ray tubes and discovered that the wavelength of the generated X-rays was dependent on the number of protons.

(Later, Moseley proposed organizing the Periodic Table according to the atomic number rather than the atomic mass)



Properties of Subatomic Particles Particle Symbol Relative Electric Charge Mass No. Relative Mass (amu)
electron e- -1 0 0.0005486
proton p+ +1 1 1.007276
neutron n0 0 1 1.008665


Dalton's theory changed again. It now states that "all atoms of an element contain the same number of protons but may contain different numbers of neutrons."


Nuclide


a particular kind of atom which contains a definite number of protons and neutrons

Nucleons


a general term pertaining to the particles which make up the atomic nucleus. These are the protons and the neutrons

Atomic Number (Z) - the number of protons


Atomic Mass (A) - The total number of nucleons in an atom:


atomic mass = number of protons + number of neutrons


number of neutrons = atomic mass - atomic number


number of neutrons = A - Z





Atomic Mass

The proton and the neutron have essentially the same mass.


An atomic mass unit (amu) is used to measure the mass of an atom. One amu is the mass of one proton or one neutron. However, we need a reference standard.


Carbon-12 was chosen as the standard reference for atomic mass. The nucleus has 6 protons and 6 neutrons. All other elements are compared to Carbon-12.


Carbon-12 atom is defined as having a mass of 12 amu


\ Atomic Mass Unit (amu) = 1/12 the mass of a Carbon-12 atom = 1.661 x 10-27 kg





Average Atomic Mass

The mass of each type of a specific atom (each nuclide) is different due to the differing numbers of neutrons. However, most elements have more than one isotope and each isotope is present in nature at some percentage. The average atomic mass (listed on the Periodic Table) is the weighted average of the various isotopes for that element. The average atomic mass can be calculated for any element if the masses of the individual isotopes and their percent abundances in nature are known.


Mass of Isotope #1 x % abundance =


Mass of Isotope #2 x % abundance =


Mass of Isotope #3 x % abundance = ____________


Average Atomic Mass S =








THE MOLE

Chemical Symbols and Formulas


Represent elements and compounds


Elemental symbol


one atom of an element

Formula of a compound


one molecule or one formula unit

Groups of Items


We often represent groups of items using such terms as:


Dozen (12)


Gross (144)


Acre (43,560 ft2; 4840 yd2; 160 rod2)


Yard (3 ft)


Yard of concrete (cubic) = 3x3x3 = 27 ft3





Mole

A term chemists use to refer to a specific number of atoms, ions, molecules, or formula units

The SI base unit representing the chemical quantity of a substance

1 mole = 6.02 x 1023 items

1 mole of particles has a mass in grams equal to that of one particle in amu.

If a single atom has a mass of 14 amu, then a mole of the same substance has a mass of 14 g.




amu/atom = grams/mole





1 mole atoms = 6.02 x 1023 atoms


1 mole molecules = 6.02 x 1023 molecules


1 mole ions = 6.02 x 1023 ions


1 mole formula units = 6.02 x 1023 formula units





Molar Mass

The mass of one mole of atoms, ions, molecules, (anything)


Convert grams to # atoms and grams to moles and vice versa


Convert grams to moles, and then convert moles to # molecules (see the worksheet on this procedure under the Resource section)


Problems:

Problem 1. You have 6.02 x 1023 atoms of Ca (calcium). How many moles do you have?


6.02 x 1023 x 1 mole/6.02 x 1023 atoms = 1.0 mole





Problem 2. How many moles do you have in 11.5 g ethanol (C2H5OH)?


1 mole ethanol has a mass of 46 g


(2x12) + (6x1) + (1x16) = 46 g


11.5 g C2H5OH x 1 mole C2H5OH/46 g C2H5OH = 0.250 g





Problem 3. You have 1.2 x 1025 molecules NH3, how many moles do you have and what is its mass?


Process: convert molecules to moles to grams


1.2 x 1025 molecules NH3 x 1 mole NH3/6.02 x 1023 x 17 g/mole = 338 g NH3





Additional Problems:


Convert 0.638 mole Ba(CN)2 to grams


Convert 50.4 g CaBr2 to moles


Convert 1.26 moles NbI5 to grams


Convert 86.2 g C2H4 to moles

CHEMISTRY: Intro. to Chemistry & Matter

Introduction to Chemistry &
Classification of Matter

Chemistry - the study of the composition and structure of materials, and the changes they undergo


Branches of Chemistry

organic the study of carbon-containing compounds

inorganic the study of all substances not classified as organic

physical the study of properties and changes involving matter and energy

analytical the identification of components and composition of materials

biochemistry the study of substances and processes occurring in living things

theoretical the use of mathematics to predict the properties of new compounds

Research
Basic research for the sake of increasing knowledge; knowledge for the sake of knowledge
Applied research carried out to solve a specific problem

Definitions:
Matter anything that has the property of inertia

Inertia a property of matter that shows itself as a resistance to any change in motion

Mass the amount of matter an object possesses

Weight the gravitational force of attraction of the earth (or other celestial body) for an object

Atom the smallest unit of an element that maintains the properties of that element

Element a pure substance consisting of only one type of atom

Compound a pure substance consisting of more than one type of atom

Extensive Properties Properties of a substance that depend on the amount of material present (mass, volume)

Intensive Properties Properties of a substance that are independent of the amount of material (color, density, m.p.)

Physical Property A characteristic that can be observed without changing the identity of the substance (melting point, boiling point, specific gravity, solubility)

Physical Change A change in a substance that does not involve changing the identity of the substance (melting, boiling, dissolution)

Chemical Property Refers to a substance's ability to undergo a change that transforms it into a different substance (reacts with acid to produce hydrogen, flammable)

Chemical Change A change in which one or more substances are converted into different substances (combustion)

Solid matter has definite shape and definite volume

Liquid matter had definite volume but takes the shape of its container

Gas matter has neither a definite volume nor a definite shape

Classification of Matter

Matter anything that has mass and occupies space

Mixture a blend of two or more kinds of matter with each component retaining its own identity

Phase any part of a system that has uniform composition and properties

Heterogeneous a mixture in which the components can be separated by physical means (e.g., filtration)

Homogeneous a mixture in which the components cannot be separated by physical means - the mixture is totally
uniform in composition

Substance matter that has a fixed composition (oxygen, table salt, sugar); element or compound

Element a pure substance consisting of only one type of atom

Compound a pure substance consisting of more than one type of atom

Classification of Elements

Metal An element that is a good conductor of heat and electricity;likes to donate electrons (aluminum, copper, magnesium)

Nonmetal An element that is a poor conductor of heat and electricity;likes to accept electrons (fluorine, nitrogen, phosphorus)

Metalloid Element possessing the properties of both metals and nonmetals (silicon, arsenic, antimony)

Noble Gases The elements of Group 18; are generally unreactive; (helium, argon, krypton)



The Periodic Table

Group/Family Vertical column of elements
Period Horizontal row of elements


Energy

Law of Conservation of Mass Mass is neither created nor destroyed during normal physical or chemical changes

Law of Conservation of Energy Energy is neither created nor destroyed during normal physical or chemical changes

Law of Conservation of Matter and Energy The sum of matter and energy is neither created nor destroyed

Energy a property of matter of which there are two types:

Kinetic Energy (KE) Energy of motion: KE = 1/2 mv2

Potential Energy (PE) Energy of position: PE = mgh

CHEMISTRY: Measurements and Calculations

SOURCE

The Scientific Method
A logical approach to solving problems by observing and collecting data, formulating hypotheses, testing hypotheses, and formulating theories that are supported by data.

System - that portion of matter in a given region of space that has been selected for study during an experiment or observation.

Scientific Method Activities

Observe
Collect Data
Formulate Hypotheses

Test Hypotheses

Theorize

The System Internationale (SI) - Metric System

1. Base Units (the MKS system: meter-kilogram-second)
length meter (m)
mass kilogram (kg)
time second (s)
temperature Kelvin (K)
amount of a substance mole (mol)
electric current ampere (amp)
luminous intensity candela (cd)


2. Derived Units

density mass / volume
velocity length / time
volume length x length x length


Conversion Factor - a ratio derived from the equality between two different units that can be used to convert from one unit to the other.

feet to inches 12 in = 1 ft
miles to feet 5280 ft = 1 mi.
inches to centimeters 1 in = 2.54 cm


Making Measurements

When making measurements with a measuring device (ruler, balance, burette,...), always estimate one place beyond the smallest calibration increment.

Accuracy and Precision (Error and Deviation)
See handout on Accuracy and Precision (on WebPage)Accuracy is a measure of how close the measurement is to the true or accepted value.Precision is a measure of how close individual measurements are to each other.
In this diagram, the X's are the individual measurements, M is the mean, d is the deviation of the measurement from the mean, E is the error of the individual measurement from the true value (or target) and Eavg is the average error.



Significant Figures
Rules

All non-zero digits are significant

All final zeros after the decimal point are significant

Zeros located between significant figures are significant

Zeros used solely for spacing the decimal point are not significant

*** The above rules can be reduced to only two rules:


All non-zero digits are significant

Zeroes are always significant unless they are used to space the decimal point (i.e., indicate the order of magnitude).

Math involving significant figures (Rules)

Rounding off -
the last significant digit is unchanged if the next digit is less than 5
the last significant digit is increased by one if the next digit is 5 or greater

Addition/Subtraction - the number of significant digits depends on the number with the largest uncertainty.

Multiplication/Division - the measurement with the smallest number of significant digits determines the number of significant digits in the answer.


Scientific Notation
Format: M x 10n


where: M is a number between 1.00 and 9.9999.....; and n is an integer (can be positive or negative)


Math involving scientific notation

addition/subtraction - ensure that the numbers have the same exponent, then add or subtract the numbers. If the exponents are different, change one of them so that they are the same, then add or subtract.

multiplication - multiply the numbers and add exponents

division - divide the numbers and subtract exponents

roots - taking the nth root of a number, divide the exponent by n

When in doubt, use three significant digits (x.xx, 0.00xxx, xxx00). Never just report all the digits that you get on your calculator.


Using Your Calculator - correctly! -

Always use EE (for exponents).


Problem-Solving Techniques

Analyze


read the entire problem carefully

define the answer required

identify the known facts and organize the information into a table or list

sketch a picture to help clarify the problem

Plan


Is the problem similar to any you have seen before?

write down any equations linking the unknown and the given information together

develop a plan for solving the problem (rearrange an equation, ??)

estimate the answer (What would be a reasonable answer?)

Compute - solve the problem - perform the mathematical steps you outlined above in your plan. Check your units and significant digits


Evaluate - does the result make sense???? Compare your answer with your estimate above.



Proportions
Direct Proportions

y is proportional to x

y/x = k or y=kx

if x increases, y also increases; when plotted, you get a straight line with a slope = k


Inverse Proportions

y is proportional to 1/x

xy = k

if x increases, y decreases; when plotted, the curve shape is a hyperbola.

Graphing

Always use graph paper! Linear or logarithmic. Select the appropriate graph paper such that you will be using about 2/3 - 3/4 of the paper for your data. This enables the reader to extrapolate.

Use a ruler to draw lines for the axes

Label the axes with appropriate descriptions, and units

Select a descriptive title for your graph.

In most cases you will draw the "best fit" for the data points you have. If the curve should be a straight line, then with a ruler draw the best fit. However, the best fit may be a curve. In such a case, draw a curved line - using French curves if they are available.


Always identify your data points clearly. If only one set of data is being plotted on the curve, use a dot with a circle around it. If additional data sets are also plotted on the same graph, select a different set of symbols for each data set. When other symbols are used, you must include a legend to explain what each set of symbols

CHEMISTRY: TABLE OF ELEMENTS

SOURCE

CLICK HERE FOR INTERACTIVE PERIODIC TABLE

OR HERE

Memorization Tips: Elements/Symbols
Unique ways to help us remember the symbols for the elements. Be warned - some are a little out there!

Silver - Ag If a person who is expecting a present of a gold necklace receives a
silver one. he might say, “Ag, I didn’t want silver!”

Gold - Au “Hey you, I want that gold necklace!” Said with “Hey you” sounding
like Au.

Aluminum - Al All cans are made from aluminum.

Bromine - Br That brother of mine - Bro of mine!

Calcium - Ca “Caws give milk!” Pronounced with an accent to make cows sound like
it’s spelled with an A.

Chlorine - Cl “You Clean with chlorine!”

Fluorine - F You need to Floss after using fluorine!

Iron - Fe “Fe, Fi, Fo, Fum, I’m an iron man!”

Helium - He If you breathe in helium, you will laugh! He, He, He!

Mercury - Hg Greek mythology - Hg stands for Helmet guy!

Potassium - K You will get Kicked out of school for the double nasty! You can’t do
the first three letters and cannot say the next three!

Sodium - Na “Naw, I don’t want any sodium!”

Nickel - Ni “Nick owes me a nickel!”

Oxygen - O “Open your mouth wide to take in oxygen!”

Lead - Pb Pencil broke!

Silicon - Si Silly con!

Tin - Sn A tin roof gets hot in the Sun.

Manganese- Mn Take first three letters - Man

Magnesium- Mg Take first three letters - Mag

CHEMISTRY TEST PACKET #1

CHEMISTRY TEST PACKET #1
(syllabus)


Teacher Note:
This packet is approximately 16 pages. (students may need to donate printer paper for this class)
Print this out and have the students fill it out while watching DVD Class #1



CHEMISTRY
TEST PACKET #1


WHAT IS MATTER?





3 STATES OF MATTER

1.

2.

3.


Can the words “Mass” and “Weight” always be used interchangeably? Why or Why not?



What is a “Change of State?”



What are the 2 kinds of property?

1.

2.


Physical Properties do what?




Give 6 examples of physical attributes or changes.












Explain in detail what Chemical Properties are..
























How does one find chemical properties?





When water decomposes it turns into what?






What are 2 types of matter?











Define both types of matter:

1.



2.





What is a substance?






What makes up homogeneous substances?







What is a solution?






What makes up heterogeneous mixtures?






What does distillation do?









What do you get when you separate a pure substance?





How many known elements are there?






What makes one element different from another?




NAME THE 3 CATEGORIES OF ELEMENTS

1.

2.

3.



Which element is on the left side of the table of elements but is NOT a metal?



NAME THE NOBLE GASES

1.

2.

3.

4.

5.

6.



Where are the Noble Gases on the Table of Elements?



Name the 3 “Inert” elements.

1.

2.

3.



Why are they considered “Inert”?



What is a compound?



Do atoms change during chemical reactions? Explain.




What makes a molecule?




What actually happens during a chemical reaction?






















FILL IN THIS CHART NEATLY…

METRIC SYSTEM PREFIXES
1
2
3
4
5
6
7
8
9
10



What is a molecular formula?



Define Precision:




Define Accuracy:




Define significant figures:









What are the rules of Sig Figs?



Addition:




Multiplication & Division:



What does Scientific Notation do?



Define dimensional analysis:



What does one use to convert units?




What does one use to convert units?




UNIT FACTOR METHOD

1.




2.



What is Stoichiometry?









What is a chemical equation?




LABEL THIS:

A + B à C + D

_____________ ____________ _______________





What is The Law of Conservation of Mass?




Balance these equations…


H2O + O2 à H2O













Fe + O2 à Fe2 O3








Iron oxide


What is the first & simplest element in the periodic table?





Draw it and Label the parts.















What is Atomic Mass?









Define Mass Number:






What is Atomic Number?











What does “z” stand for?





What does “A” stand for?






What are groups & Periods on the table of elements?






What are Isotopes?





How do you find the number of neutrons in an element?





What is amu?





What are the masses of all atoms measured by?






What is a Dalton?





What is “average mass”?






Explain the steps in detail for finding the average mass in AMU.
























Define Natural Abundance:






What is The Mole?












What is Avogadro’s Constant? Did I spell it correctly? If not, how is it spelled?








How many particles are in a mole?






What does the atomic mass of a mole equal?





What is molecular mass?






What is molar mass?





Use the Periodic Table…Find the AMU for CO2 (work it all the way out)
















Convert the AMU for CO2 to grams per mole. (work it all the way out)















Molar Mass is always expressed in what terms?









Use the Periodic Table to find the molar mass for H2O (work it all the way out)








What is the Percent Composition?





Write the percent Composition Formula?








What is the percent composition of glucose? (work it out)















Define Empirical Formula?










What is the Empirical formula for water?








What is a formula weight?





What is the difference between a Crystal and a Molecule?








Find the formula weight of Sodium Chloride. (work it out)












What is a solute?







What is a solvent?








What is molarity (M)?







What is the formula to find molarity?






What are titrations?







What happens when you mix a base solution with an acidic solution?







What is the equivalence point?






What are acid-base indicators?





What is the limiting reagent?





What is a yield?



1. Theoretical:





2. Actual:






3. Percent Yield:

CHEMICAL PERIODICITY

source

Chemical Periodicity

--------------------------------------------------------------------------------
Chemistry is made intricate and beautiful by the fact that each element reacts in a different way.
Even the ancient chemists saw evidence of quantum structure in the properties of the elements. They didn't see these properties directly, nor did they have the atomic numbers right until Moseley came along, but we know what the sizes of the atoms look like as a function of Z





What are the atomic numbers of the relatively smallest sized atoms? What do these atomic numbers (number of electrons) correspond to?

We know an even more direct measure of the atoms energy levels than that of the size of the atom. (How are size and energy related?) The Ionization Energy (or Ionization Potential) can be thought of as the negative of the energy of the highest occupied atomic orbital. The IP should increase as Z2 within a shell, if electron-electron repulsion were completely ignored.

The IP is really the Heat (Enthalpy change) of the following reaction



X(g) = X(g)+ + e- DH = IP


The Ionization Potential roughly parallels the non-metallic character of the elements, in that the easier it is to remove an electron from an element, the more metallic it is.




The first ionization potential is the energy required to remove one electron from the atom, but this process can be repeated until all the electrons are gone.



X(g)+(y-1) = X(g)+y + e- DH = IPy
The IP increases with the degree of ionization simply because of the extra energy it takes to remove an electron from an increasingly large positive charge. But, the structure of the atomic energy levels can be clearly seen upon multiple ionization of a single element or examining the periodic properties of ions of the same charge.


The Electron Affinity is defined as DH for the process


X(g) + e- = X(g)- DH = EA

Here are some E.A.'s for the main group elements. Note, some elements do not support stable negative ions, i.e. have positive EA's. (All elements have positive ions)





--------------------------------------------------------------------------------
Pauling introduced a concept that combined the IP and the EA of an element into something he called the electronegativity, which reflects a measure of how greedy for electrons an atom will be in chemical interactions with other elements.


The small 'tight' atoms have large electronegativity and the big 'squishy' atoms have low electronegativity as is shown with the dots sizes on the bargraph below:

We will use the electronegativity in our discussion of how polar a chemical bond is between unlike elements because it is a good measure of how unequally bonding electrons will be shared between atoms.

CHEMICAL BONDING

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Chemical Bonding

Though the periodic table has only 118 or so elements, there are obviously more substances in nature than 118 pure elements. This is because atoms can react with one another to form new substances called compounds. Formed when two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms.

Let's look at an example. The element sodium is a silver-colored metal that reacts so violently with water that flames are produced when sodium gets wet. The element chlorine is a greenish-colored gas that is so poisonous that it was used as a weapon in World War I. When chemically bonded together, these two dangerous substances form the compound sodium chloride, a compound so safe that we eat it every day - common table salt!

In 1916, the American chemist Gilbert Newton Lewis proposed that chemical bonds are formed
between atoms because electrons from the atoms interact with each other. Lewis had observed that many elements are most stable when they
contain eight electrons in their valence shell. He suggested that atoms with fewer than eight valence electrons bond together to share electrons and complete their valence shells.

While some of Lewis' predictions have since been proven incorrect (he suggested that electrons occupy cube-shaped orbitals), his work established the basis of what is known today about chemical bonding. We now know that there are two main types of chemical bonding; ionic bonding and covalent bonding.

Ionic Bonding

In ionic
bonding, electrons are completely transferred from one atom to another. In the process of either losing or gaining
negatively charged electrons, the
reacting atoms form ions. The oppositely
charged ions are attracted to each other by electrostatic forces, which
are the basis of the ionic bond.

For example, during the reaction of sodium with chlorine:
sodium (on the left) loses its one valence electron to chlorine resulting in

a positively charged sodium ion (left) and a negatively charged chlorine ion (right).

The reaction of sodium with chlorine

Concept simulation - Reenacts the reaction of sodium with chlorine.

Notice that when sodium loses its one valence
electron it gets smaller in size, while chlorine grows larger when it gains an
additional valence electron. This is typical of the relative sizes of ions
to atoms. Positive ions tend to be smaller than their parent atoms while
negative ions tend to be larger than their parent. After the reaction takes place, the charged Na⁺ and Cl^- ions are held together by electrostatic forces, thus forming an ionic bond. Ionic compounds share many features in common:

• Ionic bonds form between metals and nonmetals.
• In naming simple ionic compounds, the metal is always first, the nonmetal second (e.g., sodium chloride).
• Ionic compounds dissolve easily in water and other polar solvents.
• In solution, ionic compounds easily conduct electricity.
• Ionic compounds tend to form crystalline solids with high melting temperatures.

This last feature, the fact that ionic compounds are solids, results from the intermolecular forces (forces between molecules) in ionic solids. If we consider a
solid crystal of sodium chloride, the solid is made up of many positively
charged sodium ions (pictured below as small gray spheres) and an equal number of negatively charged chlorine ions (green spheres). Due to the interaction of the charged ions,
the sodium and chlorine ions are arranged in an alternating fashion as
demonstrated in the schematic. Each sodium
ion is attracted equally to all of its neighboring chlorine ions, and likewise
for the chlorine to sodium attraction. The concept of a single molecule
becomes blurred in ionic crystals because the solid exists as one continuous
system. Forces between molecules are comparable to the forces within the
molecule, and ionic compounds tend to form crystal solids with high melting points
as a result.

Covalent Bonding

The second major type of atomic bonding occurs
when atoms share electrons. As opposed to ionic bonding in which a
complete transfer of electrons occurs, covalent bonding occurs when two (or more) elements
share electrons. Covalent bonding occurs because the atoms in the compound
have a similar tendency for electrons (generally to gain electrons). This
most commonly occurs when two nonmetals bond together. Because both of the
nonmetals will want to gain electrons, the elements
involved will share electrons in an effort to fill their valence shells. A
good example of a covalent bond is that which occurs between two hydrogen
atoms. Atoms of hydrogen (H) have one valence electron in their first
electron shell. Since the capacity of this shell is two electrons, each
hydrogen atom will "want" to pick up a second electron. In an effort to
pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form
the compound H₂. Because the hydrogen compound is a combination
of equally matched atoms, the atoms will share each other's single electron,
forming one covalent bond. In this way, both atoms share the
stability of a full valence shell.

Covalent bonding between hydrogen atoms

Concept simulation - Recreates covalent bonding between hydrogen atoms.

Unlike ionic compounds, covalent molecules exist as true molecules. Because electrons are shared in covalent molecules, no full ionic charges are formed. Thus covalent molecules are not strongly attracted to one another. As a result, covalent molecules move about freely and tend to exist as liquids or gases at room temperature.

Multiple Bonds: For every pair of electrons shared between two atoms, a
single covalent bond is formed. Some atoms can share multiple pairs of
electrons, forming multiple covalent bonds. For example, oxygen (which has
six valence electrons) needs two electrons to complete its valence shell.
When two oxygen atoms form the compound O₂, they share two pairs of
electrons, forming two covalent bonds.

Lewis Dot Structures: Lewis dot structures are a shorthand to represent the valence electrons of an atom. The structures are written as the element symbol surrounded by dots that represent the valence electrons. The Lewis structures for the elements in the first two periods of the periodic table are shown below.

Lewis structures can also be used to show bonding between atoms. The bonding electrons are placed between the atoms and can be represented by a pair of dots or a dash (each dash represents one pair of electrons, or one bond). Lewis structures for H₂ and O₂ are shown below. Polar and Nonpolar Covalent Bonding
There are, in fact, two subtypes of covalent bonds. The H₂ molecule is a good example of the first type of covalent bond, the
nonpolar bond. Because both atoms in the H₂ molecule have an equal attraction (or affinity) for electrons, the bonding electrons are equally shared by the two atoms, and a nonpolar covalent bond is
formed. Whenever two atoms of the same element bond together, a
nonpolar bond is formed.

A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule.

Water molecules contain two hydrogen atoms (pictured in red) bonded to one oxygen atom (blue). Oxygen, with six valence electrons, needs two additional electrons to complete its valence shell. Each hydrogen contains one electron. Thus oxygen shares the electrons from two hydrogen atoms to complete its own valence shell, and in return shares two of its own electrons with each hydrogen, completing the H valence shells.

Polar covalent bonding simulated in water

The primary difference between the H-O bond in water and the H-H bond is the degree of electron sharing. The large oxygen atom has a stronger affinity for electrons than the small hydrogen atoms. Because oxygen has a stronger pull on the bonding electrons, it preoccupies their time, and this leads to unequal sharing and the formation of a polar covalent bond.

The Dipole
Because the valence electrons in the water molecule spend more time around the oxygen atom than the hydrogen atoms, the oxygen end of the molecule develops a partial negative charge (because of the negative charge on the electrons). For the same reason, the hydrogen end of the molecule develops a partial positive charge. Ions are not formed; however, the molecule develops a partial electrical charge across it called a dipole. The water dipole is represented by the arrow in the pop-up animation (above) in which the head of the arrow points toward the electron dense (negative) end of the dipole and the cross resides near the electron poor (positive) end of the molecule.

ELECTRON CONFIGURATION

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Electron Configuration

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The electron configuration of an atom is
a form of notation which shows how the electrons are distributed among the various atomic
orbital and energy levels. The format consists of a series of numbers, letters and
superscripts as shown below:

1s²

Here we see the
electron configuration for the element helium. This electron configuration provides
us with the following information:

• The large number "1" refers to the
  principle quantum number "n" which stands for the energy level. It tells
  us that the electrons of helium occupy the first energy level of the atom.

  
  



• The letter "s" stands for the angular
  momentum quantum number "l". It tells us that the two electrons of the
  helium electron occupy an "s" or spherical orbital.

  
  



• The exponent "2" refers to the total
  number of electrons in that orbital or sub-shell. In this case, we know that there
  are two electrons in the spherical orbital at the first energy level.

  
  

Before we continue, we must
review some information that you will need to fill in electron configurations correctly.

I. Principle Quantum Number (n) and
Sublevels

The number of sublevels that an energy level can
contain is equal to the principle quantum number of that level. So, for example, the
second energy level would have two sublevels, and the third energy level would have three
sublevels. The first sublevel is called an s sublevel. The second sublevel is
called a p sublevel. The third sublevel is called a d sublevel and the fourth
sublevel is called an f sublevel. Although energy levels that are higher than 4
would contain additional sublevels, these sublevels have not been named because no known
atom in its ground state would have electrons that occupy them.

II. Sublevels and Orbitals

An orbital is a space that can be occupied by up to
two electrons. Each type of sublevel holds a different number or orbitals, and
therefore, a different number of electrons. s sublevels have one orbital, which can
hold up to two electrons. p sublevels have three orbitals, each of which can hold 2
electrons, for a total of 6. d sublevels have 5 orbitals, for a possible total of 10
electrons. f sublevels, with 7 orbitals, can hold up to 14 electrons. The
information about the sublevels is summarized in the table below:

Table 3-6a - Orbital and Electron Capacity for the Four Named Sublevels
Sublevel	# of orbitals	Maximum number of electrons
s	1	2
p	3	6
d	5	10
f	7	14

III. Total Number of
Orbital and Electrons per Energy Level

An easy way to calculate the number of orbitals
found in an energy level is to use the formula n². For example, the third
energy level (n=3) has a total of 3², or nine orbitals. This makes sense
because we know that the third energy level would have 3 sublevels; an s sublevel with one
orbital, a p sublevel with 3 orbitals and a d sublevel with 5 orbitals. 1 + 3 + 5 =
9, so the formula n² works!

IV. Total Number of Electrons per
Energy Level

An easy way to calculate the total number of
electrons that can be held by a given energy level is to use the formula 2n².
For example, the fourth energy level (n=4) can hold 2(4)² = 32
electrons. This makes sense because the fourth energy level would have four
sublevels, one of each of the named types. The s sublevel hold 2 electrons, the p
sublevel holds 6 electrons , the d sublevel holds 10 electrons and the f sublevel holds 14
electrons. 2 + 6 + 10 + 14 = 32, so the formula 2n² works! We can
summarize this information in the table below:

Table 3-6b Orbitals and Electron Capacity of the First Four Principle Energy Levels
Principle energy level (n)	Type of sublevel	Number of orbitals per type	Number of orbitals per level(n2)	Maximum number of electrons (2n2)
1	s	1	1	2
2	s	1	4	8
	p	3
3	s	1	9	18
	p	3
	d	5
4	s	1	16	32
	p	3
	d	5
	f	7

V. Order of Filling Sublevels with
Electrons

Remember to start at the beginning of each arrow,
and then follow it all of the way to the end, filling in the sublevels that it passes
through. In other words, the order for filling in the sublevels becomes; 1s, 2s, 2p,
3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d,7p.

ELECTRON CONFIGURATION WORKSHEETSWORKSHEET 1WORKSHEET 2WORKSHEET 3WORKSHEET 4