Lab 1: Stoichiometry

NaHCO3 (aq) + CH3COOH (aq) ----> CO2 (g) + H2O (l) + CH3COONa (aq)

What is the molar ratio of NaHCO3 to CH3COONa?
What is the molar mass of each of the reactants and products?
How many moles of H2O will be produced from 2 moles of CH3COOH?
how many grams of CH3COOH?
How many grams of CO2 is produced from 3.2 grams of NaHCO3?

In this lab you will measure out the required amount and perform the combine the ingredients in a ziploc bag. You will measure the mass of the bag with the entire contents before and after the reaction.

Does your experiment abide by the law of convervation of mass? Why or why not?

Stoichiometry Review Problems

For the questions on this worksheet, consider the following equation:

Ca(OH)2(s) + 2 HCl(aq) ---> CaCl2(aq) + 2 H2O(l)


1) What type of chemical reaction is taking place? _____________________

2) How many liters of 0.100 M HCl would be required to react completely with 5.00 grams of calcium hydroxide?


3) If I combined 15.0 grams of calcium hydroxide with 75.0 mL of 0.500 M HCl, how many grams of calcium chloride would be formed?


4) What is the limiting reagent from the reaction in problem #3? __________

5) How many grams of the excess reagent will be left over after the reaction in problem 3 is complete?


Solve the following stoichiometry grams-grams problems:

1) Using the following equation:

2 NaOH + H2SO4 ---> 2 H2O + Na2SO4

How many grams of sodium sulfate will be formed if you start with 200 grams of sodium hydroxide and you have an excess of sulfuric acid?


2) Using the following equation:

Pb(SO4)2 + 4 LiNO3 ---> Pb(NO3)4 + 2 Li2SO4

How many grams of lithium nitrate will be needed to make 250 grams of lithium sulfate, assuming that you have an adequate amount of lead (IV) sulfate to do the reaction?

Percent Yield Practice

1) Write the equation for the reaction of iron (III) phosphate with sodium sulfate to make iron (III) sulfate and sodium phosphate.

2) If I perform this reaction with 25 grams of iron (III) phosphate and an excess of sodium sulfate, how many grams of iron (III) sulfate can I make?


3) If 18.5 grams of iron (III) sulfate are actually made when I do this reaction, what is my percent yield?


4) Is the answer from problem #3 reasonable? Explain.


5) If I do this reaction with 15 grams of sodium sulfate and get a 65.0% yield, how many grams of sodium phosphate will I make?

Bellringer 5

Using the equation below, calcuate the following:
Molar ratio
Molar mass
Mole to mole (given 2 moles of HBr)
Mole to gram (given 43 grams of HBr)
Gram to Percent Yield

HBr + ___ KHCO3 --->___ H2O + ___ KBr + ___ CO2

Stoichiometry Tutoring Link

go to this link to assist you with solving stoichiometry problems.

http://www.chemtutor.com/mols.htm

More Stoichiometry

1a) How many moles of chlorine gas (Cl2) would react with 5 moles of sodium (Na) according
to the following chemical equation? (Balance equation.)

Na + Cl2 --> NaCl

1b) Using the equation (after it is balanced) above, determine the amount of product that can be
produced from 24.7 g Na.

1c) How many molecules of product would be produced from 24.7g Na?

__________________________________________________________________________________

2a) In the reaction 2C8H18 + 25O2 --> 16CO2 + 18 H2O, the ratio of volumes of O2 to CO2
is _________________.

2b) If 27.3g of C8H18 are combusted, what mass of water will be produced?

2c) How many molecules of CO2 will be produced?

2d) How many atoms of H are in 2 mol of C8H18?

2e) What is the percentage, by mass, of the H in 2 mol of C8H18?

Homework 1

Balance the following equations:

1) ___ N2 + ___ F2  ___ NF3

2) ___ C6H10 + ___ O2  ___ CO2 + ___ H2O

3) ___ HBr + ___ KHCO3  ___ H2O + ___ KBr + ___ CO2

4) ___ GaBr3 + ___ Na2SO3  ___ Ga2(SO3)3 + ___ NaBr

5) ___ SnO + ___ NF3  ___ SnF2 + ___ N2O3

Using the equation from problem 2 above, answer the following questions:

6) If I do this reaction with 35 grams of C6H10 and 45 grams of oxygen, how many grams of carbon dioxide will be formed?


7) What is the limiting reagent for problem 6? ___________

8) How much of the excess reagent is left over after the reaction from problem 6 is finished?


9) If 35 grams of carbon dioxide are actually formed from the reaction in problem 6, what is the percent yield of this reaction?

Bellringer 4

Calculate the number of grams of NaCl produced as a result of 62.4 gram of sodium reacting with chlorine. The balance equation is :

Na + Cl ---> NaCl

Bellringer 3

When 84.8 grams of iron (III) oxide reacts with an excess of carbon monoxide, 54.3 grams of iron is produced.

a. Write the equation for the reaction.
b. Balance the equation.
c. Calculate the mass of reactants and products.
d. Calculate the percent yield of this reaction.

Dimensional Analysis and Stoichiometry

http://www.slideshare.net/neubla/atoms-molecules-stoichometry-i

Slide presentation of the atomic mole, molar mass and atomic units.

Chemical Equations Review

Review includes notes, bellringers, labs and vocabulary.

More review can be found at

http://misterguch.brinkster.net/worksheets.html

Bellringer 3

Balance the following equation:

HCl + NaOH => NaCl + H2O

Bellringer 2

Calcium carbonate, when heated, form calcium oxide and carbon dioxide gas.

Answer: CaCO3(s) → CaO(s) + CO2(g)

Sulfuric acid, when heated, decomposes to water and sulfur trioxide.

Answer: H2SO4 → H2O(l) + SO3(g)

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

Answer: sodium combined with water yields sodium hydroxide and hydrogen gas.

Writing Chemical Equations

http://www.chymist.com/Equations.pdf

Bellringer 1

Which of the following statements is NOT true about chemical reactions?
a. The atoms rearrange
b. Loss of mass
c. Change in energy
d. New product is formed

More Balancing Equations Practice

1
H2 + O2 => H2O
2
H3PO4 + KOH => K3PO4 + H2O
3
K + B2O3 => K2O + B
4
HCl + NaOH => NaCl + H2O
5
Na + NaNO3 => Na2O + N2
6
C + S8 => CS2
7
Na + O2 => Na2O2
8
N2 + O2 => N2O5
9
H3PO4 + Mg(OH)2 => Mg3(PO4)2 + H2O
10
NaOH + H2CO3 => Na2CO3 + H2O
11
KOH + HBr => KBr + H2O
12
H2 + O2 => H2O2
13
Na + O2 => Na2O
14
Al(OH)3 + H2CO3 => Al2(CO3)3 + H2O
15
Al + S8 => Al2S3
16
Cs + N2 => Cs3N
17
Mg + Cl2 => MgCl2
18
Rb + RbNO3 => Rb2O + N2
19
C6H6 + O2 => CO2 + H2O
20
N2 + H2 => NH3
21
C10H22 + O2 => CO2 + H2O
22
Al(OH)3 + HBr => AlBr3 + H2O
23
CH3CH2CH2CH3 + O2 => CO2 + H2O
24
C + O2 => CO2
25
C3H8 + O2 => CO2 + H2O
26
Li + AlCl3 => LiCl + Al
27
C2H6 + O2 => CO2 + H2O
28
NH4OH + H3PO4 => (NH4)3PO4 + H2O
29
Rb + P => Rb3P
30
CH4 + O2 => CO2 + H2O
31
Al(OH)3 + H2SO4 => Al2(SO4)3 + H2O
32
Na + Cl2 => NaCl
33
Rb + S8 => Rb2S
34
H3PO4 + Ca(OH)2 => Ca3(PO4)2 + H2O
35
NH3 + HCl => NH4Cl
36
Li + H2O => LiOH + H2
37
Ca3(PO4)2 + SiO2 + C => CaSiO3 + CO + P
38
NH3 + O2 => N2 + H2O
39
FeS2 + O2 => Fe2O3 + SO2
40
C + SO2 => CS2 + CO

Balancing Chemical Equations Practice

Balancing Chemical Equations Worksheet
1. _____ H2 + _____ O2 �� _____ H2O
2. _____ N2 +_____ H2 ��_____ NH3
3. _____ S8 + _____ O2 �� _____ SO3
4. _____ N2 + _____ O2 �� _____ N2O
5. _____ HgO �� _____ Hg + _____ O2
6. _____ CO2 + _____ H2O �� _____ C6H12O6 + _____ O2
7. _____ Zn + _____ HCl �� _____ ZnCl2 + _____ H2
8. _____ SiCl4 + _____ H2O �� _____ H4SiO4 + _____ HCl
9. _____ Na + _____ H2O �� _____ NaOH + _____ H2
10. _____ H3PO4 �� _____ H4P2O7 + _____ H2O
11. _____ C10H16 + _____ Cl2 �� _____ C + _____ HCl
12. _____ CO2 + _____ NH3 �� _____ OC(NH2)2 + _____ H2O
13. _____ Si2H3 + _____ O2 �� _____ SiO2 + _____ H2O3
14. _____ Al(OH)3 + _____ H2SO4 �� _____ Al2(SO4)3 + _____ H2O
15. _____ Fe + _____ O2 �� _____ Fe2O3
16. _____ Fe2(SO4)3 + _____ KOH �� _____ K2SO4 + _____ Fe(OH)3
17. _____ C7H6O2 + _____ O2 �� _____ CO2 + _____ H2O
18. _____ H2SO4 + _____ HI �� _____ H2S + _____ I2 + _____ H2O
19. _____ FeS2 + _____ O2 �� _____ Fe2O3 + _____ SO2
20. _____ Al + _____ FeO �� _____ Al2O3 + _____ Fe
21. _____ Fe2O3 + _____ H2 �� _____ Fe + _____ H2O
22. _____ Na2CO3 + _____ HCl �� _____ NaCl + _____ H2O + _____ CO2
23. _____ K + _____ Br2 �� _____ KBr
24. _____ C7H16 + _____ O2 �� _____ CO2 + _____ H2O
25. _____ P4 + _____ O2 �� _____ P2O5

Balancing Chemical Equations

http://richardbowles.tripod.com/chemistry/balance.htm

This link will give you the essentials of balancing a chemical equation.

http://www.fordhamprep.org/gcurran/sho/sho/lessons/lesson81.htm

Creating an atom inventory.

Chemical Equation Notes

Reaction Equations

Key terms

Energy, exothermic reaction, endothermic reactionPhysical reactions, chemical reactions, phase transitionsReactants, productsReaction stoichiometry

Skills to develop
To distinguish chemical changes from physical changes.
To write chemical equations to describe a chemical reaction.
To balance chemical equations.
To calculate the quantities of reactants required or the quantities produced in a chemical reaction.

Chemical Reaction Equations

Changes in a material or system are called reactions, and they are divided into chemical and physical reactions.
Energy is the driving force of all changes, both physical and chemical reactions. Energy is always involved in these reactions. If a system is more stable by losing some energy, a reaction takes place, releasing energy. Such a reaction is said to be exothermic. Supplying energy to a system also causes a reaction. Energy absorbing reactions are called endothermic reactions. Sometimes, the amount of energy involved in a reaction may be so small that the change in energy is not readily noticeable.

An equation can be used to describe a physical reaction, which involves a change of states. For example, melting, sublimation, evaporation, and condensation can be represented as follow.

In these equations, (s) stands for solid, (l) for liquid (l), and (g) for gas,

H2O(s) ® H2O(l) . . . melting
H2O(s) ® H2O(g) . . . sublimation
C2H5OH(l) ® C2H5OH(g) . . . evaporation
NH3(g) ® NH3(l) . . . condensationIn these changes, no chemical bonds are broken or formed, and the molecular identities of the substances have not changed.

Is the phase transition between graphite and diamond is a chemical or physical reaction?
C(graphite) ® C(diamond).

The crystal structures of diamond and graphite are very different, and bonding between the carbon atoms are also different in the two solid states. Because chemical bonds are broken and new bonds are formed, the phase transition of diamond and graphite is a chemical reaction.
Chemicals or substances change converting to one or more other substances, and these changes are called chemical reactions. At the molecular level, atoms or groups of atoms rearrange resulting in breaking and forming some chemical bonds in a chemical reaction. The substances undergoing changes are called reactants, whereas substances newly formed are called products. Physical appearances of products are often different from reactants. Chemical reactions are often accompanied by the appearance of gas, fire, precipitate, color, light, sound, or odor.

These phenomena are related to energy and properties of the reactants and products. For example, the oxidation of propane releases heat and light, and a rapid reaction is an explosion,
C3H8 + 5 O2 ® 3 CO2 + 4 H2O A balanced equation also shows a macroscopic quantitative relationship. This balanced reaction equation shows that five moles of oxygen reacts with one mole of propane generating three moles of carbon dioxide and four moles of water, a total of 7 moles of products in the combustion reaction.

At the molecular level, this equation shows that for each propane molecule, 5 oxygen molecules are required. The three carbon atoms are converted to three molecules of carbon dioxide, whereas the 8 hydrogen atoms in propane are oxidized to 4 water molecules. The numbers of H, C, and O atoms are the same on both sides of the equation.

We study properties of substances so that we know how to make use of them. Tendencies of a substance to react, either by itself or with others, are important chemical properties. Via properties, we understand chemical reactions, which are best studied by experimentation and observation. After you have performed many experiments, you may generalize certain rules and facts. Knowing these rules and facts enable you to solve problems that you have not yet encountered.

The most important aspect of a chemical reaction is to know what are the reactants and what are the products. For this, the best description of a reaction is to write an equation for the reaction. A chemical reaction equation gives the reactants and products, and a balanced chemical reaction equation shows the mole relationships of reactants and products. Often, the amount of energy involved in the reaction is given. Dealing with the quantitative aspect of chemical reactions is called reaction stoichiometry.

For example, when clamshells, CaCO3, are heated, a gas CO2 will be released, leaving a white powder (solid CaO) behind. This reaction is represented by the reaction as depicted in the picture, and the equation of the reaction is written as:
CaCO3 ® CaO + CO2The equation indicates that one mole of CaCO3 gives one mole each of CaO and CO2. Amounts of substances represented by chemical formulas have been introduced on the two previous pages, and these concepts should help to figure out the stoichiometry of reactions when a reaction equation is given.


Example 1
When 10.0 g pure calcium carbonate is heated and converted to solid calcium oxide CaO, how much calcium oxide should be obtained? If the only 5.0 grams CaO is obtained, what is the actual yield? Hint:Under ideal condition, amounts of substance in the reaction equation is as indicated below:
CaCO3 ® CaO + CO2100.0 . . . . . 56 . . . 44 g/mol (formula weights) 1 mol CaCO3 1 mol CaO 56 g CaO
10.0 g CaCO3 ------------ ----------- --------- = 5.6 g CaO
100 g CaCO3 1 mol CaCO3 1 mol CaO
DiscussionAn inefficient conversion is given here, but the method shows the details of consideration. If the amount of CaO obtained is not 5.6 g, one can conclude that the sample may not be pure.

Example 2
When 10.0 g pure calcium carbonate is heated and converted to solid calcium oxide CaO, how much CO2 at standard condition is released? Hint:
CaCO3 ® CaO + CO2 1 mol CO2 22.4 L CO2
10.0 g CaCO3 ----------- ----------- = 2.24 L CO2
100 g CaCO3 1 mol CO2

Discussion

We have taken a short cut in this formulation compared to Example 1. Example 1 and 2 illustrate the evaluation of quantities in g and in L.

Writing Equations for Chemical Reactions
Chemical reaction equations truly represent changes of materials. For many reactions, we may only be able to write equations for the overall reactions. For example, common sense tells us that when sugar is fully oxidized, carbon dioxide and water are the final products. The oxidation reaction is the same as the combustion reaction. Thus we write
C12H22O11 + 12 O2 ® 12 CO2 + 11 H2O This illustrates the methods used for writing balanced reaction equations:

Determine the reactants and productsIn this case, the products are CO2 and H2O, determined by common sense. We know that.

Apply the fundamental principle of conservation of atomsNumbers of atoms of each kind must be the same before and after the reactions.
Balance one type of atoms at a timeBR> We may use H or C to begin. Since there are 12 C atoms on the left, the coefficient is 12 for CO2. Similarly, 22 H atoms produce 11 H2O molecules.
Balance the oxygen atoms on both sides. There are a total of 35 O atoms on the right hand, and the coefficient for O2 should be 11.

Example 3
The compound N2O5 is unstable at room temperature. It decomposes yielding a brown gas NO2 and oxygen. Write a balanced chemical reaction equation for its decomposition. Hint:The first step is to write an unbalanced equation indicating only the reactant and products:
N2O5 ® NO2 + O2 A N2O5 molecule decomposes into two NO2 molecule, and half of O2.
N2O5 ® 2 NO2 + 1/2O2 In order to give whole number stoichiometric coefficients to the equation, we multiply all the stoichiometric coefficients by 2.
2 N2O5 ® 4 NO2 + O2 DiscussionThis example illustrate the steps used in writing a balance equation for a chemical reaction. This balanced equation does not tell us how a N2O5 molecule decompose, it only illustrate the overall reaction.

Example 4
When solutions of CaCl2 and AgNO3 are mixed, a white precipitate is formed. The same precipitate is also observed when NaCl solution is mixed with AgCH3CO2 solution. Write a balanced equation for this the reaction between CaCl2 and AgNO3. Hint:The common ions between NaCl and CaCl2 are Cl- ions, and Ag+ ions are common between the two silver containing compounds. The question illustrates a scientific deduction used in the determination of products. The product is AgCl, and the balanced reaction is
CaCl2 + 2 AgNO3 ® 2 AgCl + Ca(NO3)2 DiscussionIn reality, solutions of salts contain ions. In this case, the solutions contain Ca2+, Cl-, Ag+, and NO3- ions. The Cl- and Ag+ ions form an insoluble solid, and a precipitate is formed,
Cl- + Ag+ ® AgCl(s) Ca2+ and NO3- are by-stander ions.

Chemical Reactions
One of the most important topics in chemistry is chemical reaction. In this page, we only concentrate on the stoichiometry conveyed by reaction equations.

Other topics related to chemical reactions are:
Excess and Limiting Reagents or reactants left over or used upFeatures of chemical reactions or classification of reactionsChemical kinetics or reaction ratesReaction mechanism or how actually reaction proceedThe first two topics are included in this group, but the later topics will be discussed in another course (CHEM123).
Balancing Redox Reactions Balancing oxidation and reduction reaction equations is a little more complicated than what we discussed here. You have to have the skills to assign oxidation states, explain oxidation and reduction in terms of oxidation-state change, and write half reaction euqations. Then you will be able to balance redox reactions. All these are given in the next module on Chemical Reactions.

http://www.science.uwaterloo.ca/~cchieh/cact/c120/reaction.html

Chemical Equation Vocabulary

activity series of metals
complete ionic equation
balanced equation
decomposition reaction
catalyst
double-replacement reaction
chemical equation
net ionic equation
coefficient
single-replacement reaction
combination reaction
skeleton equation
combustion equation
spectator ion

Electron Configuration with Aufbau Diagram

Electron Configuration


http://www.teachersdomain.org/assets/wgbh/phy03/phy03_doc_qmechatom/phy03_doc_qmechatom.pdf

The electron configuration of an atom denotes the distribution of electrons among available shells. The standard notation lists the subshell symbols, one after another. The number of electrons contained in each subshell is stated explicitly. For example, the electron configuration of beryllium, with an atomic (and electron) number of 4, is 1s22s2 or [He]2s2.

Electron Configuration Notes

Electron Dot Configurations

http://www.uoregon.edu/~ch111/L12.htm

Contructing Lewis Dot StructuresStarting with a structure indicating only atom connections (single bonds), you can practice constructing a Lewis dot structure. Just click on the atom or bond you wish to modify. Nonzero formal charges are indicated for each atom in the structure once the total number of electrons is correct.
A recommended procedure might be:

Count the total number of valence electrons (N) needed to account for the atoms (based on the column of the atom in the periodic table) and charge (add one electrons for each negative charge; subtract one electron for each positive charge).
Draw the framework with single bonds. Some knowledge of the way the atoms are connected may be required.
Using lone pairs, complete octets around the noncentral atoms.
Count the number of electrons depicted (two for each bond and two for each lone pair). If this number is less than N, then add electrons to the central atom until the total number of electrons depicted is N.
If the octet rule is not satisfied for the central atom and lone-pair electrons are nearby, use those electrons to make double or triple bonds to the central atom.
Check each atom to see if it has a formal charge. (Singly bonded oxygen will require a negative charge, for example.)

BOHR MODEL
http://csep10.phys.utk.edu/astr162/lect/light/bohr.html

VESPR THEORYhttp://cost.georgiasouthern.edu/chemistry/general/molecule/vsepr.htm

Five Major Chemical Reactions Video

Law of Conservation Video

http://www.youtube.com/watch?v=dExpJAECSL8

Naming Compounds Practice

Ionic Compound Names and Formulas

For the list on the left, name the compound. For the list on the right, give the chemical formula that corresponds to the name

Name Formula
1) NaF 13) potassium fluoride
2) K2CO3 14) ammonium sulfate
3) MgCl2 15) magnesium iodide
4) Be(OH)2 16) copper (II) sulfite
5) SrS 17) aluminum phosphate
6) Cu2S 18) lead (II) nitrite
7) ZnI2 19) cobalt (II) selenide
8) Ca3(PO4)2 20) silver cyanide
9) NH4I 21) copper (II) bicarbonate
10) Mn(NO3)3 22) iron (II) oxide
11) FePO4 23) lithium cyanide
12) CoCO3 24) lead (IV) sulfite


Naming Covalent Compounds Worksheet

Write the formulas for the following covalent compounds:

1) antimony tribromide __________________________________

2) hexaboron silicide __________________________________

3) chlorine dioxide __________________________________

4) hydrogen iodide __________________________________

5) iodine pentafluoride __________________________________

6) dinitrogen trioxide __________________________________

7) ammonia __________________________________

8) phosphorus triiodide __________________________________


Write the names for the following covalent compounds:

9) P4S5¬ __________________________________

10) O2 __________________________________

11) SeF6 __________________________________

12) Si2Br¬6 __________________________________

13) SCl4 __________________________________

14) CH4 __________________________________

15) B2Si __________________________________

16) NF3 _________________________

Bellringers 1-4

1. Valence electrons are the electrons that

a. orbit the outside shell.
b. are stable.
c. do not have a charge.
d. have no energy.

2. What is the difference between a monatomic and polyatomic ion?

3. Complete the following table:

Element # Electrons # Valence Electrons Oxidation # (Charge)
Na
S
He
Ar
Si

4. Name the following ionic compounds:

Ca3(PO4)2
Mn(NO3)3
CoCO3

Common Ions

Naming Compounds

This link will provide with you a series of videos on naming chemical compounds:

http://www.onlinemathlearning.com/chemical-names.html

Electron Configuration Notes

Compounds

A compound is a group of atoms with a specific number and type of atoms arranged in a specific way. Exactly the same elements in exactly the same proportions are in every bit of the compound.
Example: Water is a compound composed of one oxygen atom and two hydrogen atoms. Each hydrogen atom is attached to an oxygen atom by a chemical bond. H2O is the formula for the compound, water.

If any other elements are attached, it is not water. For example, H2S is hydrogen sulfide. Hydrogen sulfide does not have the same types of atoms as water, so it is a different compound.
If a different number of atoms of hydrogen or oxygen are attached, it is not water. H2O2 is the formula for hydrogen peroxide. It might have the right elements in it to be water, but it does not have them in the right proportion. A molecule is a single formula of a compound joined by covalent bonds.

The Law of Constant Proportions states that a given compound always contains the same proportion by weight of the same elements.

Electron Configuration and Valence Electrons
In a stable atom, the number of electrons is equal to the number of protons.
Electrons in atoms are present in discrete orbits or "shells" around the nucleus of the atom.
There is a ranking or heirarchy of the shells, with the shells further from the nucleus having a higher energy.
The innermost electron shell holds only two electrons.
The outermost shell contains the valence electrons. The maximum number of electrons that can occupy the outer shell is eight. When there are eight electrons in the outer shell, it is said to have an octet of electrons.
The valence of an atom is the likely charge it will take on as an ion.
A valence is the amount of positive or negative charge on an ion of an element.
Example: Hydrogen only has one electron. It can lose an electron to become H+, a hydrogen ion, or it can gain an electron to become H-, a hydride ion.

The Octet Rule
The octet rule states that atoms are most stable when they have a full shell of 8 electrons in the outside electron shell.
Octet = 8
An atom with eight electrons in the outer shell is more stable than an atom which as fewer electrons in the outer shell.

The exception to this is Helium (atomic number 2) which only has two electrons in its outer shell. It has a full shell, so it is a stable inert element.

Valence electrons are the only electrons involved in chemical bonds.

Atoms will form chemical bonds with other atoms by either sharing electrons, or by transferring electrons in order to complete their octet and get 8 electrons in the outer shell.

Ions
In a stable atom, the number of electrons is equal to the number of protons.
An atom which has a different number of electrons than it does protons is called an ion.

Ions are charged particles. Types of ions:

Cation - A positively charged ion.A cation is an atom or group of atoms with a net positive charge, caused by the loss of one or more electrons. Examples: Na+, NH4+, Mg+2

Anion - a negatively charged ion.An anion is an atom or group of atoms with a net negative charge, caused by the gain of one or more electrons.Examples: F-, S2-, NO3-

Polyatomic ion - a group of atoms which function as a group and which has a net positive or negative charge (cation or anion).Examples: NH4+ or NO3-

The Periodic Chart can show how the octet rule works. All of the Group I elements have one electron in the outside shell and they all have a valence of plus one. Group I elements will lose that one electron in the outside shell, to become a single positive ion with a full electron shell of eight electrons (an octet) in the s and p subshells under it.

Bonding
A bond is an attachment among atoms. Atoms may be held together for any of several reasons, but all bonds have to do with the electrons (particularly the outside electrons) of atoms.
There are several types of bonds:

Ionic bonds occur due to a full electrical charge difference attraction.

Covalent bonds occur due to sharing electrons.
There are bonds that come about from partial charges or the position or shape of electrons about an atom.

Ionic Bonds
The attraction between a positive ion and a negative ion is an ionic bond.

Some atoms (such as metals) tend to lose electrons to make the outside ring of electrons more stable. When an atom loses electrons it becomes a positive ion (or cation) because the number of protons exceeds the number of electrons.

Other atoms tend to gain electrons to complete the outside electron ring. The non-metal ions tend to gain electrons to fill out the outer shell. When the number of electrons exceeds the number of protons, the ion is negative. (Non-metal ions and most of the polyatomic ions have a negative charge.)

Ionic compounds - composed of cations and anions which are ionically bonded to each other due to attractions of opposite charges
1. Cations and anions combine in a ratio that produces a neutral compound; smallest whole number ratio is used for formula of an ionic compound.
e.g., Na+ + Cl- --> NaCl (one of each is needed to balance the charges: +1 and -1)
Mg+2 + Cl- ---> MgCl2(two Cl's are needed to balance the charges since Cl is -1 and Mg is +2 charge)
2. Cation is listed first, then anion in the formula

pH Notes and Calculations

http://www.sciencegeek.net/Chemistry/chempdfs/pHCalculations.pdf

Lab 2: Penny Lab: Cohesion/Adhesion

Penny Lab: Post- Lab Analysis

3. Explain what surface tension is.
4. Why were many trials taken and averaged?
5. In this experiment, what was your control group?
6. Identify the independent variable in the experiment.
7. Identify the dependent variable in the experiment.
8. What if the experimental question was "How does sugar affect the surface tension of water?" Describe how you would answer this question using the scientific method. If you have time, you can test this.

Bellringer 7-8

A substance with a pH of 2 is ____.
a. Acidic
b. Basic
c. Caustic
d. Volatile

The water molecule is ______ .

a. polar
b. non-polar
c. bipolar
d. codependent

Water Molecule Notes

http://www.bethel.k12.or.us/schools/teachers/bcollins/IB%20Biology%20files/chemistry%20of%20water.pdf

The Water Molecule

This week we will discuss the properties of the water molecule. Due to its structure, the ability of the water molecule to bond with other substances makes it unique and the universal solvent for many substances.

The link will give you an overall view of the water molecule:

http://www.chem1.com/acad/sci/aboutwater.html

SUMMARY OF TRENDS IN THE PERIODIC TABLE

SUMMARY OF TRENDS IN THE PERIODIC TABLE can be found at this site
http://www.avon-chemistry.com/p_table_lecture.html

Summary of Trends

Moving Left --> Right

Atomic Radius Decreases
Ionization Energy Increases
Electronegativity Increases

Moving Top --> Bottom
Atomic Radius Increases
Ionization Energy Decreases
Electronegativity Decreases

Lab 2: We are Family

Objective:

Today's lab will focus on the grouping of families of elements. Their characteristics and behavior are the same but their reactivity towards each other is different. You will be forming compounds with elements from the same family and observe its reaction.

Adopt An Element Project

Project Timeframe
Sept 18 Assignment- Begin Reading
Sept 30/Oct 1 Library Research
Due Oct 7


Requirements:
1) Complete an Adopt An Element information sheet. (60% of grade)
You may use a variety of reference sources. Possible ideas are encyclopedias (book
or CD Rom), science encyclopedias, science catalogs, magazines, and/or Internet sites*.
Information sheets must be neat, written in black ink, and contain all the information
requested. You also need to provide a list of your sources on the back of your
information sheet. A minimum of three sources are required.


2) Create an advertisement for your element. (40% of grade)
The advertisement must include the element’s name, symbol, atomic number,
atomic mass, cost, and an advertising slogan that describes one or more of its important
uses. Advertisements must be neat, colorful, and contain all the information listed
above. You may add pictures that relate to your advertisement theme.


Be sure to include:
Ö Element’s symbol
Ö Element’s name
Ö Atomic number
Ö Atomic mass
Ö Ad slogan
Ö Cost
Ö Your name


You may add pictures or
drawings that illustrate
the various uses for your
element.


Your ad must follow
the same format
as this example!
Atomic Number
33 74.9
As
Arsenic
Arsenic’s a sure fire way
to deal with a nasty rat,
It works better than
a mean old cat!
Cost = $3.20 for 1 gram
John Smith
Atomic Mass
Symbol & Name
Slogan
Cost
Name

Student Conferences This Week

Lab 1: We Are Family

Objective:

Today's lab will focus on the grouping of families of elements. Their characteristics and behavior are the same but their reactivity towards each other is different. You will be forming compounds with elements from the same family and observe its reaction.

Bellringer 6

Match the word to its definition.

Atomic Radius __
Electronegativity (Affinity) __
Ionization Energy __
Reactivity __

A. the amount of energy required to remove the outmost electron. It is closely related to electronegativity

B. an atom's 'desire' to grab another atom's electrons

C. refers to how likely or vigorously an atom is to react with other substances. This is usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom's electrons (electronegativity) because it is the transfer/interaction of electrons that is the basis of chemical reactions

D. an indication of the atom's volume

Bellringer 5

Which element is in Group 3, Period 4?

Which of the following element is NOT a metal?
a. Li
b. Mg
c. Al
d. Br

Which of the following element is NOT found in air?
a. C
b. O
c. Ne
d. H

Period Table Summary of Trends

SUMMARY OF TRENDS IN THE PERIODIC TABLE can be found at this site
http://www.avon-chemistry.com/p_table_lecture.html

Summary of Trends

Moving Left --> Right

Atomic Radius Decreases
Ionization Energy Increases
Electronegativity Increases

Moving Top --> Bottom
Atomic Radius Increases
Ionization Energy Decreases
Electronegativity Decreases

Period Table...More Trends

Review
Period - a row of elements on the periodic table. Remember that sentences are written in rows and end with a period.

Group - a column of elements on the periodic table. Remember that group is spelled group and groups go up and down.

Atomic Radius - Atomic radius is simply the radius of the atom, an indication of the atom's volume.

Period - atomic radius decreases as you go from left to right across a period.
Why? Stronger attractive forces in atoms (as you go from left to right) between the opposite charges in the nucleus and electron cloud cause the atom to be 'sucked' together a little tighter.

Group - atomic radius increases as you go down a group.
Why? There is a significant jump in the size of the nucleus (protons + neutrons) each time you move from period to period down a group. Additionally, new energy levels of elections clouds are added to the atom as you move from period to period down a group, making the each atom significantly more massive, both is mass and volume.

Electronegativity - Electronegativity is an atom's 'desire' to grab another atom's electrons.

Period - electronegativity increases as you go from left to right across a period.
Why? Elements on the left of the period table have 1 -2 valence electrons and would rather give those few valence electrons away (to achieve the octet in a lower energy level) than grab another atom's electrons. As a result, they have low electronegativity. Elements on the right side of the period table only need a few electrons to complete the octet, so they have strong desire to grab another atom's electrons.

Group - electronegativity decreases as you go down a group.
Why? Elements near the top of the period table have few electrons to begin with; every electron is a big deal. They have a stronger desire to acquire more electrons. Elements near the bottom of the chart have so many electrons that loosing or acquiring an electron is not as big a deal. This is due to the shielding affect where electrons in lower energy levels shield the positive charge of the nucleus from outer electrons resulting in those outer electrons not being as tightly bound to the atom.

Ionization Energy - Ionization energy is the amount of energy required to remove the outmost electron. It is closely related to electronegativity.

Period - ionization energy increases as you go from left to right across a period.
Why? Elements on the right of the chart want to take others atom's electron (not given them up) because they are close to achieving the octet. The means it will require more energy to remove the outer most electron. Elements on the left of the chart would prefer to give up their electrons so it is easy to remove them, requiring less energy (low ionization energy).

Group - ionization energy decreases as you go down a group.
Why? The shielding affect makes it easier to remove the outer most electrons from those atoms that have many electrons (those near the bottom of the chart).

Reactivity - Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom's electrons (electronegativity) because it is the transfer/interaction of electrons that is the basis of chemical reactions.

Metals

Period - reactivity decreases as you go from left to right across a period.

Group - reactivity increases as you go down a group
Why? The farther to the left and down the periodic chart you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity.Non-metals

Period - reactivity increases as you go from the left to the right across a period. Group - reactivity decreases as you go down the group.
Why? The farther right and up you go on the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electron.

Ionic Radius vs. Atomic Radius

Metals - the atomic radius of a metal is generally larger than the ionic radius of the same element.
Why? Generally, metals loose electrons to achieve the octet. This creates a larger positive charge in the nucleus than the negative charge in the electron cloud, causing the electron cloud to be drawn a little closer to the nucleus as an ion.

Non-metals - the atomic radius of a non-metal is generally smaller than the ionic radius of the same element.
Why? Generally, non-metals loose electrons to achieve the octet. This creates a larger negative charge in the electron cloud than positive charge in the nucleus, causing the electron cloud to 'puff out' a little bit as an ion.

Melting Point

Metals - the melting point for metals generally decreases as you go down a group.

Non-metals - the melting point for non-metals generally increases as you go down a group.

http://www.geocities.com/CapeCanaveral/Lab/4097/chem/chap4/periodictrends.html
from

Bellringer 4

The three major groups of elements on the periodic table are:

a. Metals, non-metals, gases
b. Metals, mettaloids, gases
c. Metals, alkalines, lactinides
d. Metals, alkaline earth metals, gases

Bellringer 3

Which statement is true?

a. There are 18 groups and 7 periods on the periodic table.
b. There are 18 periods and 7 groups on the periodic table.
c. There are 18 classes and 7 kingdoms on the periodic table.
d. There are 18 families and 7 groups on the periodic table.

Bellringer 2

Determine the number of protons, neutrons and electrons for the following elements:

S
Cu
Fe

Bellringer 1

Atoms contain

a. protons
b. electrons
c. neutrons
d. all of the above

Color Code the Periodic Table

Color Coding the Periodic Table
Student Information Sheet

The Periodic Table is a list of all the known elements. It is organized by increasing atomic number. There are two main groups on the periodic table: metals and nonmetals. The left side of the table contains elements with the greatest metallic properties. As you move from the left to the right, the elements become less metallic with the far right side of the table consisting of nonmetals. The elements in the middle of the table are called “transition” elements because they are changed from metallic properties to nonmetallic properties. A small group whose members touch the zigzag line are called metalloids because they have both metallic and nonmetallic properties.
The table is also arranged in vertical columns called “groups” or “families” and horizontal rows called “periods.” Each arrangement is significant. The elements in each vertical column or group have similar properties. Group 1 elements all have the electron in their outer shells. This gives them similar properties. Group 2 elements all have 2 electrons in their outer shells. This also gives them similar properties. Not all of the groups, however, hold true for this pattern. The elements in the first period or row all have one shell. The elements in period 2 all have 2 shells. The elements in period 3 have 3 shells and so on.
There are a number of major groups with similar properties. They are as follows:
Hydrogen: This element does not match the properties of any other group so it stands alone. It is placed above group 1 but it is not part of that group. It is a very reactive, colorless, odorless gas at room temperature. (1 outer level electron)
Group 1: Alkali Metals – These metals are extremely reactive and are never found in nature in their pure form. They are silver colored and shiny. Their density is extremely low so that they are soft enough to be cut with a knife. (1 outer level electron)
Group 2: Alkaline-earth Metals – Slightly less reactive than alkali metals. They are silver colored and more dense than alkali metals. (2 outer level electrons)
Groups 3 – 12: Transition Metals – These metals have a moderate range of reactivity and a wide range of properties. In general, they are shiny and good conductors of heat and electricity. They also have higher densities and melting points than groups 1 & 2. (1 or 2 outer level electrons)
Lanthanides and Actinides: These are also transition metals that were taken out and placed at the bottom of the table so the table wouldn’t be so wide. The elements in each of these two periods share many properties. The lanthanides are shiny and reactive. The actinides are all radioactive and are therefore unstable. Elements 95 through 103 do not exist in nature but have been manufactured in the lab.
Group 13: Boron Group – Contains one metalloid and 4 metals. Reactive. Aluminum is in this group. It is also the most abundant metal in the earth’s crust. (3 outer level electrons)
Group 14: Carbon Group – Contains on nonmetal, two metalloids, and two metals. Varied reactivity. (4 outer level electrons)
Group 15: Nitrogen Group – Contains two nonmetals, two metalloids, and one metal. Varied reactivity. (5 outer level electrons)
Group 16: Oxygen Group – Contains three nonmetals, one metalloid, and one metal. Reactive group. (6 outer level electrons)
Groups 17: Halogens – All nonmetals. Very reactive. Poor conductors of heat and electricity. Tend to form salts with metals. Ex. NaCl: sodium chloride also known as “table salt”. (7 outer level electrons)
Groups 18: Noble Gases – Unreactive nonmetals. All are colorless, odorless gases at room temperature. All found in earth’s atmosphere in small amounts. (8 outer level electrons)
Color Coding the Periodic Table
Student Worksheet

This worksheet will help you understand how the periodic table is arranged. Your teacher will give you a copy of the periodic table to color. Using map pencils, color each group on the table as follows:

1. Color the square for Hydrogen pink.
2. Lightly color all metals yellow.
3. Place black dots in the squares of all alkali metals.
4. Draw a horizontal line across each box in the group of alkaline earth metals.
5. Draw a diagonal line across each box of all transition metals.
6. Color the metalloids purple.
7. Color the nonmetals orange.
8. Draw small brown circles in each box of the halogens.
9. Draw checkerboard lines through all the boxes of the noble gases.
10. Using a black color, trace the zigzag line that separates the metals from the nonmetals.
11. Color all the lanthanides red.
12. Color all the actinides green.

When you are finished, make a key that indicates which color identifies which group.

Dalton's Atomic Theory

Dalton's Atomic Theory

1) All matter is made of atoms. Atoms are indivisible and indestructible.
2) All atoms of a given element are identical in mass and properties
3) Compounds are formed by a combination of two or more different kinds of atoms.
4) A chemical reaction is a rearrangement of atoms

Periodic Table Vocabulary

alkali metals
group
nonmetal
alkaline earth metal
halogen
nucleus
atom
inner transition metal
period
atomic mass
isotope
periodic law
atomic mass unit (amu)
mass number
periodic table
atomic number
metal
proton
cathode ray
metalloid
representative element
Dalton’s atomic theory
neutron
transition metal
electron
noble gas

Introduction to the Periodic Table Vocabulary

History
Mendeleev
Organization
Periods, Groups, Families
Trends
Atomic Radii
Ionization
Electricity

September 27/28

First Six Weeks Exam

Lab Journal Check

You must have the following in your journal as September 25, 2010

Lab 1 Advertisement
Lab 2 Milk
Lab 3 Sugar Dissolving Rates
Lab 4 Separation of Mixtures

Bellringer 1-9

TAKS April 2009 Benchmark

http://ritter.tea.state.tx.us/student.assessment/resources/release/tests2009/taks_g10_science.pdf

Atomic Math Games, Periodic Table and Videos

Link to the following for a copy of the periodic table:
http://www.nysedregents.org/testing/reftable/archreftable/chempertable.pdf

For Atomic Math Calculations Game, go to

http://education.jlab.org/elementmath/index.html

For a video to explain atomic number and mass, go to

http://www.nysedregents.org/testing/reftable/archreftable/chempertable.pdf

Atomic Theory Notes

education.jlab.org/jsat/powerpoint/atomos.ppt

Bellringer Week 2

1. The smallest particle of matter is a ____ .
a. atom
b. electron
c. proton
d. quark

2. Physical changes are indicated by
a. change of substance
b. rearranging the atoms
c. new products formed
d. new atoms formed

Separation Techniques

Distillation
METHODS of SEPARATING MIXTURES and purifying substances
Simple Distillation

Distillation involves 2 stages and both are physical state changes.

(1) The liquid or solution mixture is boiled to vaporise the most volatile component in the mixture (liquid ==> gas). The ant-bumping granules give a smoother boiling action.

(2) The vapour is cooled by cold water in the condenser to condense (gas ==> liquid) it back to a liquid (the distillate) which is collected.

This can be used to purify water because the dissolved solids have a much higher boiling point and will not evaporate with the steam, BUT it is too simple a method to separate a mixture of liquids especially if the boiling points are relatively close.


Fractional Distillation
Fractional Distillation
Fractional distillation involves 2 main stages and both are physical state changes. It can only work with liquids with different boiling points. However, this method only works if all the liquids in the mixture are miscible (e.g. alcohol/water, crude oil etc.) and do NOT separate out into layers like oil/water.

(1) The liquid or solution mixture is boiled to vaporise the most volatile component in the mixture (liquid ==>gas). The ant-bumping granules give a smoother boiling action.

(2) The vapour passes up through a fractionating column, where the separation takes place (theory at the end). This column is not used in the simple distillation described above.

(3) The vapour is cooled by cold water in the condenser to condense (gas ==> liquid) it back to a liquid (the distillate) which is collected.

This can be used to separate alcohol from a fermented sugar solution.

It is used on a large scale to separate the components of crude oil, because the different hydrocarbons have different boiling and condensation points (see oil).

FRACTIONAL DISTILLATION THEORY:

Imagine green liquid is a mixture of a blue liquid (boiling point 80oC) and a yellow liquid (boiling point 100oC), so we have a coloured diagram simulation of a colourless alcohol and water mixture! As the vapour from the boiling mixture enters the fractionating column it begins to cool and condense. The highest boiling or least volatile liquid tends to condense more i.e. the yellow liquid (water). The lower boiling more volatile blue liquid gets further up the column. Gradually up the column the blue and yellow separate from each other so that yellow condenses back into the flask and pure blue distils over to be collected. The 1st liquid, the lowest boiling point, is called the 1st fraction and each liquid distils over when the top of the column reaches its particular boiling point to give the 2nd, 3rd fraction etc.

To increase the separation efficiency of the tall fractionating column, it is usually packed with glass beads, short glass tubes or glass rings etc. which greatly increase the surface area for evaporation and condensation.

In the distillation of crude oil the different fractions are condensed out at different points in a huge fractionating column. At the top are the very low boiling fuel gases like butane and at the bottom are the high boiling big molecules of waxes and tar.


Chromatography
Paper or Thin Layer Chromatography
This method of separation is used to see what coloured materials make up e.g. a food dye analysis.

The material to be separated e.g. a food dye (6) is dissolved in a solvent and carefully spotted onto chromatography paper or a thin layer of a white mineral material on a glass sheet. Alongside it are spotted known colours on a 'start line' (1-5).

The paper is carefully dipped into a solvent, which is absorbed into the paper and rises up it. The solvent may be water or an organic liquid like an alcohol (e.g. ethanol) or a hydrocarbon, so-called non-aqueous solvents. For accurate work the distance moved by the solent is marked on carefully with a pencil and the distances moved by each 'centre' of the coloured spots is also measured. These can be compared with known substances BUT if so, the identical paper and solvent must be used (See Rf values below).

Due to different solubilities and different molecular 'adhesion' some colours move more than others up the paper, so effecting the separation of the different coloured molecules.

Any colour which horizontally matches another is likely to be the same molecule i.e. red (1 and 6), brown (3 and 6) and blue (4 and 6) match, showing these three are all in the food dye (6).

The distance a substance moves, compared to the distance the solvent front moves (top of grey area on 2nd diagram) is called the reference or Rf valueand has a value of 0.0 (not moved - no good), to 1.0 (too soluble - no good either), but Rf ratio values between 0.1 and 0.9 can be useful for analysis and identification.

Rf = distance moved by dissolved substance (solute) / distance moved by solvent.

Some technical terms: The substances (solutes) to be analysed must dissolve in the solvent, which is called the mobile phase because it moves. The paper or thin layer of material on which the separation takes place is called the stationary or immobile phase because it doesn't move.

It is possible to analyse colourless mixture if the components can be made coloured e.g. protein can be broken down into amino acids and coloured purple by a chemical reagent called Ninhydrin and many colourless organic molecules fluoresce when ultra-violet light is shone on them. These are called locating agents.

Thin layer chromatograpy (t.l.c) is where a layer of paste is thinly and evenly spread on e.g. a glass plate. The paste consists of the solid immobile phase like aluminium oxide dispersesd in a liquid such as water. The plate is allowed to dry and then used in the same way as paper chromatography.


Crystallization


Filtration

Go to this link to view seperation diagrams:
http://www.gcsescience.com/e4-mixture-separation.htm
Posted by Hao Tran at 8:09 PM 0 comments

Compounds, Mixtures, Elements and Separation Methods

Here's a good resource for definitions of
compounds, mixtures, elements and seperation methods

http://www.docbrown.info/page01/ElCpdMix/EleCmdMix.htm#Introduction

Seperation of Mixtures Quiz

Go to this link, answer the quiz and print your results.

http://www.gcsescience.com/q/qelcomsep.html

Lab 4: Separation of Mixtures

Sand, salt, iron filings, seeds

Mixtures Vocabulary

Chromatography
Concentrated
Magnet
Evaporation
Solvent
Solute
Dissolve
Solution
Crystalline
Amorphous
Suspension
Insoluble
Dilute
Saturated
Mixture
Filtration

Seperation of Mixtures

Evaporation
Filtration
Paper Chromatography
Distillation
Centrifuge
Seperating funnel
Magnet
Decanting

Lab 3: Case of the Disappearing Sugar Cube

STANDARD Students will observe and describe chemical and physical change.

OBJECTIVE
Differentiate between common chemical and physical changes.
Analyze factors that influence chemical and physical change.

INTENDED LEARNING OUTCOMES

1a. Make observations and measurements
2d. Collect and record data using procedures designed to minimize error.
2e. Analyze data and draw warranted inferences.

Introduction

To help teach students how stirring, temperature, concentration, surface area and crushing affect reaction rates. This is an open ended activity which can also be used to assess students understanding of the affects of variables on reaction rates. Students should work in groups of at least two.

Materials

1 sugar cube per group
stirring rods
mortar and pestle or something to crush the cubes
50 or 100 ml beakers
Hot plate and pot to boil water with a ladle
graduated cylinders
stop watches

Procedure

1. Place 1 sugar cube in a 50 or 100 ml beaker with 40 ml of cold water and time how long it takes to dissolve and record the time. Have students use the classroom clock to time this one.
2. Place 1 sugar cube crushed in a 50 or 100 ml beaker with 40 ml of cold water, time how long it takes to dissolve and record the time.
3. Place 1 sugar cube crushed in a 50 or 100 ml beaker with 40 ml of cold water and stir it until it dissolves. Time how long it takes to dissolve and record the time.
4. Place 1 sugar cube crushed in a 50 or 100 ml beaker with 40 ml of hot water and stir it until it dissolves. Time how long it takes to dissolve and record the time.

Analysis

Begin discussion by asking students what things were the same in the four procedures? What things did we change? What affect did these changes have on the time it took the sugar to dissolve? What is a variable? Define for students what a variable is and then have them come up with a procedure that will dissolve the sugar cube the fastest and then have them test their hypothesis by racing against the rest of the students in the class.

Variation

If you have the time you can have the students do first solid cube cold water and then solid cube hot water. Then have the students do a crushed cube in cold and hot water. Finally have them do a crushed cube with stirring in cold and hot water.

Physical and Chemical Properties Powerpoint

http://www.dentonisd.org/52720812134413/lib/52720812134413/_files/Physical_Chemical_Properties.pdf

Physical and Chemical Properties/States of Matter

The properties of a substance are those characteristics that are used to identify or describe it. When we say that water is "wet", or that silver is "shiny", we are describing materials in terms of their properties. Properties can be divided into the categories of physical properties and chemical properties. Physical properties are readily observable, like; color, size, luster, or smell. Chemical properties are only observable during a chemical reaction. For example, you might not know if sulfur is combustible unless you tried to burn it.

Another way of separating kinds of properties is to think about whether or not the size of a sample would affect a particular property. No matter how much pure copper you have, it always has the same distinctive color. No matter how much water you have, it always freezes at zero degrees Celsius under standard atmospheric conditions. Methane gas is combustible, no matter the size of the sample. Properties, which do not depend on the size of the sample involved, like those described above, are called intensive properties. Some of the most common intensive properties are; density, freezing point, color, melting point, reactivity, luster, malleability, and conductivity.

Extensive properties are those that do depend on the size of the sample involved. A large sample of carbon would take up a bigger area than a small sample of carbon, so volume is an extensive property. Some of the most common types of extensive properties are; length, volume, mass and weight.

Pieces of matter undergo various changes all of the time. Some changes, like an increase in temperature, are relatively minor. Other changes, like the combustion of a piece of wood, are fairly drastic. These changes are divided into the categories of Physical and Chemical change. The main factor that distinguishes one category form the other is whether or not a particular change results in the production of a new substance.

Physical changes are those changes that do not result in the production of a new substance. If you melt a block of ice, you still have H2O at the end of the change. If you break a bottle, you still have glass. Painting a piece of wood will not make it stop being wood. Some common examples of physical changes are; melting, freezing, condensing, breaking, crushing, cutting, and bending. Special types of physical changes where any object changes state, such as when water freezes or evaporates, are sometimes called change of state operations.

Chemical changes, or chemical reactions, are changes that result in the production of another substance. When you burn a log in a fireplace, you are carrying out a chemical reaction that releases carbon. When you light your Bunsen burner in lab, you are carrying out a chemical reaction that produces water and carbon dioxide. Common examples of chemical changes that you may be somewhat familiar with are; digestion, respiration, photosynthesis, burning, and decomposition.




http://www.fordhamprep.org/gcurran/sho/sho/lessons/lesson15.htm

More Vocabulary

Chapter 1
analytical chemistry
hypothesis
physical chemistry
biochemistry
inorganic chemistry
scientific law
chemistry
observation
scientific method
experiment
organic chemistry
theory

Chapter 2
chemical property
homogeneous mixture
physical change
chemical reaction
law of conservation of mass
physical property
chemical symbol
liquid
product
compound
mass
reactant
distillation
matter
solid
element
mixture
solution
gas
phase
substance
heterogeneous mixture
vapor

Lab 2: Milk Lab Video and Write Up

http://schoolwaxtv.com/op_video/1434/embed" width="450" height="337" frameborder="0" scrolling="no">">

Lab Instructions and Video
http://hhs.tsc.k12.in.us/webpages/teacherpages/teachers/bcreech/Lab1-1.pdf

Metrics Conversion Practice

Write the correct abbreviation for each metric unit.
1) Kilogram _____ 4) Milliliter _____ 7) Kilometer _____
2) Meter _____ 5) Millimeter _____ 8) Centimeter _____
3) Gram _____ 6) Liter _____ 9) Milligram _____

Try these conversions, using the ladder method.
1) 2000 mg = _______ g 6) 5 L = _______ mL 11) 16 cm = _______ mm
2) 104 km = _______ m 7) 198 g = _______ kg 12) 2500 m = _______ km
3) 480 cm = _____ m 8) 75 mL = _____ L 13) 65 g = _____ mg
4) 5.6 kg = _____ g 9) 50 cm = _____ m 14) 6.3 cm = _____ mm
5) 8 mm = _____ cm 10) 5.6 m = _____ cm 15) 120 mg = _____ g

Compare using <, >, or =.
16) 63 cm 6 m 17) 5 g 508 mg 18) 1,500 mL 1.5 L
19) 536 cm 53.6 dm 20) 43 mg 5 g 21


http://sciencespot.net/Media/metriccnvsn2.pdf

Significant Figures

The rules for identifying significant digits when writing or interpreting numbers are as follows:
All non-zero digits are considered significant. For example, 91 has two significant digits (9 and 1), while 123.45 has five significant digits (1, 2, 3, 4 and 5).
Zeros appearing anywhere between two non-zero digits are significant. Example: 101.12 has five significant digits: 1, 0, 1, 1 and 2.
Leading zeros are not significant. For example, 0.00052 has two significant digits: 5 and 2.
Trailing zeros in a number containing a decimal point are significant. For example, 12.2300 has six significant digits: 1, 2, 2, 3, 0 and 0. The number 0.000122300 still has only six significant digits (the zeros before the 1 are not significant). In addition, 120.00 has five significant digits. This convention clarifies the precision of such numbers; for example, if a result accurate to four decimal places is given as 12.23 then it might be understood that only two decimal places of accuracy are available. Stating the result as 12.2300 makes clear that it is accurate to four decimal places.
The significance of trailing zeros in a number not containing a decimal point can be ambiguous. For example, it may not always be clear if a number like 1300 is accurate to the nearest unit (and just happens coincidentally to be an exact multiple of a hundred) or if it is only shown to the nearest hundred due to rounding or uncertainty. Various conventions exist to address this issue:
A bar may be placed over the last significant digit; any trailing zeros following this are insignificant. For example, has three significant digits (and hence indicates that the number is accurate to the nearest ten).
The last significant digit of a number may be underlined; for example, "20000" has two significant digits.
A decimal point may be placed after the number; for example "100." indicates specifically that three significant digits are meant.[1]
However, these conventions are not universally used, and it is often necessary to determine from context whether such trailing zeros are intended to be significant. If all else fails, the level of rounding can be specified explicitly. The abbreviation s.f. is sometimes used, for example "20 000 to 2 s.f." or "20 000 (2 sf)". Alternatively, the uncertainty can be stated separately and explicitly, as in 20 000 ± 1%, so that significant-figures rules do not apply.

Metrics practice

1. Rounded correctly, 2.000 cm × 10.0 cm =
20.000 cm2 20.00 cm2 20 cm2 20.0 cm2
2. The number of significant figures in 0.00230300 m is
9 6 4 3 8
3. 5.5234 mL of mercury is transfered to a graduated cylinder with scale marks 0.1 mL apart. Which of the following will be the correct reading taken from the graduated cylinder?
5.5234 mL 5.52 mL 5.523 mL 5 mL 5.5 mL
4. Correctly rounded, 20.0030 - 0.491 g =
19.5120 g 19.512 g 19.5 g 20 g 19.51 g
5. Correctly rounded, the quotient 2.000 g / 20.0 mL is
0.100 g/mL 0.1000 g/mL 0.1 g/mL 0.10 g/mL

Metrics

Learning objectives

Use the SI system.
Know the SI base units.
State rough equivalents for the SI base units in the English system.
Read and write the symbols for SI units.
Recognize unit prefixes and their abbreviations.
Build derived units from the basic units for mass, length, temperature, and time.
Convert measurements from SI units to English, and from one prefixed unit to another.
Use derived units like density and speed as conversion factors.
Use percentages, parts per thousand, and parts per million as conversion factors.
Use and report measurements carefully.
Consider the reliability of a measurement in decisions based on measurements.
Clearly distinguish between
precision and accuracy
exact numbers and measurements
systematic error and random error
Count the number of significant figures in a recorded measurement. Record measurements to the correct number of digits.
Estimate the number of significant digits in a calculated result.
Estimate the precision of a measurement by computing a standard deviation.
Lecture outline

Measurement is the collection of quantitative data. The proper handling and interpretation of measurements are essential in chemistry - and in any scientific endeavour. To use measurements correctly, you must recognize that measurements are not numbers. They always contain a unit and some inherent error. The second lecture focuses on an international system of units (the SI system) and introduces unit conversion. In the third lecture, we'll discuss ways to recognize, estimate and report the errors that are always present in measurements.

Measurement

quantitative observations
include 3 pieces of information
magnitude
unit
uncertainty
measurements are not numbers
numbers are obtained by counting or by definition; measurements are obtained by comparing an object with a standard "unit"
numbers are exact; measurements are inexact
mathematics is based on numbers; science is based on measurement
The National Institute of Standards and Technology (NIST) has published several online guides for users of the SI system.
The SI System

Le Systéme Internationale (SI) is a set of units and notations that are standard in science.
Four important SI base units (there are others)
Quantity SI
Base Unit English
Equivalent
length meter (m) 1 m = 39.36 in
mass kilogram (kg) 1 kg = 2.2 lbs
time second (s)
temperature kelvin (K) °F = 1.8(oC)+32
K = °C + 273.15
derived units are built from base units
Some SI derived units
Quantity Dimensions SI units Common name
area length × length m2 square meter
velocity length/time m/s
density mass/volume kg/m3
frequency cycles/time s-1 hertz (Hz)
acceleration velocity/time m/s2
force mass × acceleration kg m/s2 Newton (N)
work, energy, heat force × distance kg m2/s2 Joule (J)
Prefixes are used to adjust the size of base units
Commonly used SI prefixes (there are others).
Prefix Meaning Abbreviation Exponential
Notation
Giga- billion G 109
Mega- million M 106
kilo- thousand k 103
centi- hundredths of c 10-2
milli- thousandths of m 10-3
micro- millionths of µ 10-6
nano- billionths of n 10-9
pico- trillionths of p 10-12
several non-SI units are encountered in chemistry
Non SI unit Unit type SI conversion Notes
liter (L) volume 1 L = 1000 cm3 1 quart = 0.946 L
Angstrom (Å) length 1 Å = 10-10 m typical radius of an atom
atomic mass unit (u) mass 1 u = 1.66054×10-27 kg about the mass of a proton or neutron; also known as a 'dalton' or 'amu'


Arithmetic with units

addition and subtraction: units don't change
2 kg + 3 kg = 5 kg
412 m - 12 m = 400 m
consequence: units must be the same before adding or subtracting!
3.001 kg + 112 g = 3.001 kg + 0.112 kg = 3.113 kg
4.314 Gm - 2 Mm = 4.314 Gm - 0.002 Gm = 4.312 Gm
multiplication and division: units multiply & divide too
3 m × 3 m = 9 m2
10 kg × 9.8 m/s2 = 98 kg m/s2
consequence: units may cancel
5 g / 10 g = 0.5 (no units!)
10.00 m/s × 39.37 in/m = 393.7 in/s


Converting Units

5 step plan for converting units
identify the unknown, including units
choose a starting point
list the connecting conversion factors
multiply starting measurement by conversion factors
check the result: does the answer make sense?
Common variations
series of conversions
example: Americium (Am) is extremely toxic; 0.02 micrograms is the allowable body burden in bone. How many ounces of Am is this?
converting powers of units
converting compound units
starting point must be constructed
using derived units as conversion factors
mass fractions (percent, ppt, ppm) convert mass of sample into mass of component
density converts mass of a substance to volume
velocity converts distance traveled to time required
concentration converts volume of solution to mass of solute
Uncertainty in Measurements

making a measurement usually involves comparison with a unit or a scale of units
always read between the lines!
the digit read between the lines is always uncertain
convention: read to 1/10 of the distance between the smallest scale divisions
significant digits
definition: all digits up to and including the first uncertain digit.
the more significant digits, the more reproducible the measurement is.
counts and defined numbers are exact- they have no uncertain digits!
Tutorial: Uncertainty in Measurement
counting significant digits in a series of measurements
compute the average
identify the first uncertain digit
round the average so the last digit is the first uncertain digit
counting significant digits in a single measurement
convert to exponential notation
disappearing zeros just hold the decimal point- they aren't significant.
exception: zeros at the end of a whole number might be significant
Precision of Calculated Results
calculated results are never more reliable than the measurements they are built from
multistep calculations: never round intermediate results!
sums and differences: round result to the same number of fraction digits as the poorest measurement
products and quotients: round result to the same number of significant digits as the poorest measurement.
Quiz
Using Significant Figures
Precision vs. Accuracy
good precision & good accuracy
poor accuracy but good precision

good accuracy but poor precision
poor precision & poor accuracy


Precision Accuracy
reproducibility correctness
check by repeating measurements check by using a different method
poor precision results from poor technique poor accuracy results from procedural or equipment flaws
poor precision is associated with 'random errors' - error has random sign and varying magnitude. Small errors more likely than large errors. poor accuracy is associated with 'systematic errors' - error has a reproducible sign and magnitude.
Estimating Precision
Consider these two methods for computing scores in archery competitions. Which is fairer?
Score by distance from bullseye
Score by area or target
The standard deviation, s, is a precision estimate based on the area score: where
xi is the i-th measurement
is the average measurement
N is the number of measurements.
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Details

http://antoine.frostburg.edu/chem/senese/101/measurement/index.shtml

Lab 1: Advertisement

Lab Safety ppt
Scientific Method ppt
Observations
Hypothesis
Experiment
Data
Conclusion

Students choose an advertisement. Using scientific method to "prove" the ad. Hypothesis must be testable.
Data can be qualitative or quantitative or both.

First Day of School

A day, 3 pre-AP Chem classes, discussed syllabus and evaluated ourselves with Gardner's multiple intelligence. Assignment was to log into the class blog.

Welcome Video

Create your own video slideshow at animoto.com.

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Students Tips on How to Reach the Top

Tarrant County high schools are loaded with smart students, with infinitely small percentage points separating the best from the rest.
But the students who make it to the top of their high school graduating class don't get there by accident. Well before they enter their freshman year, these students have figured out what it takes to be valedictorian or salutatorian.
"Ever since I was in elementary school, I have been very focused on keeping high grades," said Katie Skinner, valedictorian at Calvary Christian Academy in Fort Worth.
Before Skinner and several other top 2010 graduates packed up and headed off to college, the Star-Telegram asked them to outline eight strategies and traits to earn a top ranking.
Some of the qualities are common sense.
When teachers say it's smart to get plenty of rest before a big test, believe them. Pulling an all-nighter to cram does more harm than good.
Participate in class.
Don't cheat.
Here are other ways to think like a valedictorian .
Go for it
Choose your courses carefully and don't settle for the "recommended" path to graduation.
When Dat Nguyen's family moved from Vietnam in 2004, he spent three months at the Fort Worth school district's International Newcomer Academy. After less than a year at Meadowbrook Middle School's Language Center, he finished eighth grade among the top students in his class and in 2010 was Dunbar High School's valedictorian.
"It's about making a goal and sticking with it. It's all about the mindset at the beginning," he said. "I just kind of looked around to see who is No. 2 and No. 3, and I always wanted to get a better grade."
Nguyen, 18, will study biomedical engineering at the University of Texas at Arlington.
Load up on advanced classes
Understand how grade-point average is calculated, because fractions of a point can be the difference. It's not enough to be a straight-A student. Many schools use a weighted GPA system to calculate class rank, with higher points for more-challenging courses.
Several valedictorians said they took as many Advanced Placement and honors courses as they could handle.
And they stayed away from unweighted courses that bring down GPAs.
Students in the Arlington school district who meet certain criteria can exclude certain courses from GPA calculations.
Bowie High School valedictorian Kosisio Mora and her twin sister, Ifunanya Mora, the Arlington school's salutatorian, both used that option for a nonhonors anatomy course.
"It wasn't that it was too hard; it just wouldn't help me," said Ifunanya Mora, 18.
The Mora twins, of Grand Prairie, left Thursday for University of the Incarnate Word in San Antonio. Both plan to major in biology/pre-med.
Challenge yourself
Don't get sidetracked by hanging out with friends, then try to tackle a semester project in one weekend.
Working toward valedictorian helped motivate the students to keep academics their top priority.
"I decided I wanted to be valedictorian as a freshman. I thought if I set a high goal for myself, it would help me stay focused and keep me from slacking off," said Skinner, 18, who will major in telecommunications and media studies at Texas A&M University in College Station. "I prayed about it daily. I would ask God to help me keep the right mindset and keep the right goals."

Don't settle for less
Be proactive. If you get stuck with a bad teacher, transfer to another class. If your guidance counselor is not effective, ask for a different one.
Keller Central High School valedictorian Forrest Ripley said he researched which teachers were best at their subject before signing up for classes. He asked upperclassmen and his two older brothers for advice on what teachers to avoid.
But he was assigned to six different counselors in four years of high school, leaving him largely on his own in selecting classes, Ripley said.
"My counselors were not very helpful, so I didn't rely on them," said Ripley, 19, who will study business at the University of Texas at Austin. "So much is getting the schedule that you want. Be prepared to be involved."
Homework and extra credit
Students say it's important to go beyond assigned class work and homework.
Always do extra credit, and research subjects that pique your interest.
Asked whether he had studied a lot, Nguyen replied, "Not really."
We disagree.
Nguyen said that after he finished his assigned reading, math problems and other homework, he would study two more hours each day to prepare for classes and tests.
Students say it's critical to study every day and to plan ahead for big tests.
And never turn in work late.
Pay attention to details
Read the requirements of each assignment. Review main course points with the teacher and ask what will be on tests. Use a planner to keep track of assignment due dates, upcoming tests and long-term projects. Double-check your work.
Kosisio Mora had a PowerPoint presentation graded down because she did not follow instructions to put a photograph in every slide.
"I did everything right," she said. "But since I didn't pay attention to that detail, it cost me."
Work with others
Take responsibility for your schoolwork, but it's smart to cultivate teachers and upperclassmen as allies, get tutoring in weaker subjects and study with other high-achieving classmates.
"You can't understand everything. You can't be a genius in everything, so if you help someone else out, then they're more apt to help you," said Brooke Awtry, 18, salutatorian of the first graduating class at Westlake Academy charter school.
"Have some friends over, study for a couple hours and then watch a movie or have dinner," said Awtry, a Keller resident who will study English and international affairs at Southern Methodist University in Dallas. "That was a way to easily put together studying and having a life."
Get parents involved
Parents who are plugged into their children's school help them succeed.
They meet and communicate regularly with teachers and administrators and are often involved in booster clubs and parent-teacher organizations. That's where they find out about scholarships, tutoring and other opportunities.
"The parents are aware of what is going on. If there is an opportunity for a kid to take, it's the parents that are around there talking to one another and talking to teachers," said Jennifer Latu, lead counselor at Fossil Ridge High School in the Keller district.


Read more: http://www.star-telegram.com/2010/08/21/2417308_p2/how-to-become-a-valedictorian.html#ixzz0xLAXWk5s


Read more: http://www.star-telegram.com/2010/08/21/2417308/how-to-become-a-valedictorian.html#my-headlines-default#ixzz0xLAQkOEj
http://www.star-telegram.com/2010/08/21/2417308/how-to-become-a-valedictorian.html#my-headlines-default

Week 1 Bellringers

1. The statement, "A chemical reaction never creates products that weigh more or less than the reactants", is based on three centuries of experimental observation. The statement is an example of:
a. a hypothesis b. a theory c. a datum d. a law


2. A hypothesis is
a. obeyed under any circumstances.
b. a theory that has been proved
c. a tentative explanation for a natural phenomenon
d. a description of a pattern or relationship in experimental data

3. A number of people become ill after eating dinner in a restaurant. Which of the following statements is a hypothesis?
a. The cooks felt really bad about it.
b. Everyone who ate oysters got sick.
c. Bacteria in the oysters may have caused the illness.
d. Symptoms include nausea and dizziness
e. People got sick whether the oysters were raw or cooked.

4.A natural law is
a. a description of a pattern or relationship in experimental data
b. an explanation that has been proved
c. a tentative explanation for a natural phenomenon
d. obeyed under any circumstances.

5. Which of the following is least important to know about a liquid solution you are using during a laboratory investigation?
a. price of the solution per mL
b. flammability of the solution
c. first aid procedures to follow for skin contact
d. recommended procedures for appropriate Disposal

Chapter 1 Vocabulary

biochemistry

The chemistry of living things, including the structure and function of biological molecules and the mechanism and products of their reactions.

chemistry
The study of matter and its transformations. See What is chemistry? for other definitions.

computer-assisted drug design
Using computational chemistry to discover, enhance, or study drugs and related biologically active molecules.

computational chemistry
A branch of chemistry concerned with the prediction or simulation of chemical properties, structures, or processes using numerical techniques.

dependent variable

Compare with independent variable.
A dependent variable changes in response to changes in independent variables. For example, in an experiment where the vapor pressure of a liquid is measured at several different temperatures, temperature is the independent variable and vapor pressure is the dependent variable.

environmental chemistry

chemical ecology
The study of natural and man-made substances in the environment, including the detection, monitoring, transport, and chemical transformation of chemical substances in air, water, and soil.

experiment
An experiment is direct observation under controlled conditions. Most experiments involve carefully changing one variable and observing the effect on another variable (for example, changing temperature of a water sample and recording the change volume that results).

geochemistry
geological chemistry
The study of materials and chemical reactions in rocks, minerals, magma, seawater, and soil.

hypothesis
hypotheses Compare with theory
A hypothesis is a conjecture designed to guide experimentation. Hypotheses are extremely useful in problem solving, and are essential in developing new theories.

independent variable

Compare with dependent variable.
An independent variable that can be set to a known value in an experiment. Several independent variables may be controlled in an experiment. For example, in an experiment where the vapor pressure of a liquid is measured at several different temperatures, temperature is the independent variable and vapor pressure is the dependent variable.

inorganic chemistry.
The study of inorganic compounds, specifically their structure, reactions, catalysis, and mechanism of action.

law
natural law; scientific law
Natural laws summarize patterns that recur in a large amount of data. Unlike human laws, natural laws don't forbid or permit; they describe.

matter

Matter is anything that has mass. Air, water, coffee, fire, human beings, and stars are matter. Light, X-rays, photons, gravitons, information, and love aren't matter.

medicinal chemistry

A branch of chemistry concerned with the discovery, design, synthesis, and investigation of biologically active compounds and reactions that these compounds undergo in living things.

organic chemistry

The study of compounds that contain carbon chemically bound to hydrogen, including synthesis, identification, modelling, and reactions of those compounds.

pharmacology

The study of drugs, which includes determination of biological activity, biological effects, breakdown and synthesis, and delivery.
pharmacognosy. Identification, isolation, and characterization of biologically active substances in living things.

physical chemistry
chemical physics
A branch of chemistry that studies chemical phenomena from a physical and mathematical perspective. Physical chemistry includes chemical thermodynamics, kinetics, spectroscopy, quantum chemistry, and statistical mechanics.

scientific notation
exponential notation.

A system for reporting very small or very large numbers by writing the number as a decimal number between 1 and 10, multiplied by a power of 10. For example, 602000000000000000000000 is written in scientific notation as 6.02 x 1023. 0.000323 is written in scientific notation as 3.23 x 10-4.

theory. theories
Compare with hypothesis.
Theories are well-established explanations for experimental data. To become established, the theory must experimentally tested by many different investigators. Theories usually can not be proven; a single contrary experiment can disprove a theory.

toxicology

The study of poisons, including identification, isolation, biological effects, mechanism of action, and development of antidotes.

variable
Compare with independent variable and dependent variable.
A quantity that can have many possible values. In designing experiments, variables that affect measurements must be identified and controlled. For example, an experiment that measures reaction rates must control temperature, because temperature is a variable that can change the rate of reaction.

Definition of Chemistry

chem·is·try n., pl. -tries. 1. the science that systematically studies the composition, properties, and activity of organic and inorganic substances and various elementary forms of matter. 2. chemical properties, reactions, phenomena, etc.: the chemistry of carbon. 3. a. sympathetic understanding; rapport. b. sexual attraction. 4. the constituent elements of something; the chemistry of love. [1560-1600; earlier chymistry].
The first definition captures many of the essential ingredients of chemistry (although definitions 3 and 4 might make a more entertaining paper):

Chemistry is a science. There is only one sanctioned procedure for determining whether a statement about matter is really chemistry: the exhaustive, inefficient, but highly successful scientific method. Chemists often arrive at new results by nonscientific means (like luck or sheer creativity), but their work isn't chemistry unless it can be reproduced and verified scientifically.
Chemistry is a systematic study. Chemists have devised several good methods for solving problems and making observations. For example, analytical chemists often use protocols (thoroughly tested recipes) for determining the concentrations of substances in a sample. Chemists use well-defined techniques like spectroscopy and chromatography to study new or unknown substances.
Chemistry is the study of the composition and properties of matter. Chemistry answers questions like, "What kind of stuff is this sample made of? What does the sample look like on a molecular scale? How does the structure of the material determine its properties? How do the properties of the material change when I increase temperature, or pressure, or some other environmental variable?"
Chemistry is the study of the reactivity of substances. One material can be changed into another by a chemical reaction. A complex substance can by made from simpler ones. Chemical compounds can break down into simpler substances. Fuels burn, food cooks, leaves turn in the fall, cells grow, medicines cure. Chemistry is concerned with the essential processes that make these changes happen.
Chemistry is the study of organic and inorganic substances. Organic substances contain hydrogen combined with carbon; inorganic substances don't. It was once believed that organic compounds were exclusively produced by living things, but today chemists can synthesize many organic materials from inorganic ones. Carbon can link with itself and other atoms in many diverse ways, and its chemistry is far more complex than that of other elements. So while the organic/inorganic distinction is artificial, it's still useful.
Chemistry is the study of connections between the everyday world and the molecular world. Chemists use atoms and molecules to explain properties and behaviors of matter. For example, you can find molecular explanations for flavor and color changes elsewhere on this site.
If you'd like some historical perspective, a good reference is The Enlighment of Matter: The Definition of Chemistry from Agricola to Lavoisier, by Marco Baretta.

Author: Fred Senese senese@antoine.frostburg.edu

Introduction

Learning objectives
• State the central objectives of chemistry (and this course).
• Outline the scientific method.
o Classify statements and explanations as observations, experimental data, laws , hypotheses , or theories . Quiz
o Understand the importance of making controlled comparisons and obtaining reproducible data.
Lecture outline
The introductory lecture discusses the scope, objectives, and methods of chemistry.
What is Chemistry?
• the study of matter and its transformations
• the study of connections between molecular and macroscopic events
Why Study Chemistry?
• learn fundamental physical models
• gain technical perspective on current events
• develop problem solving skills
• appreciate life's little mysteries
The Scientific Method
• a systematic procedure for solving problems and exploring natural phenomena
• Observations (data)
o are the foundation of the scientific method
o data can be qualitative or quantitative.
o data is most useful when collected under controlled conditions (experiments )
o experiments must be repeatable and reproducible
• Natural laws
o compactly summarize patterns in a large amount of data
o often apply only under special conditions
o are descriptions of nature, not facts or explanations
• Hypotheses
o tentative explanations designed to guide experimentation
o a useful hypothesis must be testable
o must be rejected or corrected when they conflict with experiment
• Theories
o a well-tested explanation for experimental data based on a set of hypotheses.
o must be discarded or refined when they can't explain new experimental results
o scientific theories have three aspects: philosophical, mathematical, and empirical.
 Understand all three, or risk misusing the theory!
o a good theory...
 explains currently available data
 is as simple as possible (but no simpler!)
 accurately predicts results of future experiments
 suggests new lines of work and new ways to think
 clearly shows underlying connections
• Serendipity and intuition
o Many important scientific discoveries were not arrived at using the scientific method
 Charles Goodyear - vulcanization of rubber
 Teflon
 Plastics
 Saccharin


http://antoine.frostburg.edu/chem/senese/101/intro/index.shtml

First Week of School

I. Introductions, Room Orientation, Technology

II. Icebreaker

III. Syllabus and class rules

IV. Why study chemistry? What is chemistry?

V. Matter Concept Map

VI. Language of science

Lab Journals

VII. Lab safety

Student Expectations

Expectations of Me

1. I will always do my best.
2. I will show respect for myself, my teacher, my peers, and materials.
3. I have read, understood, and agreed to the terms of the safety contract.
4. I will take responsibility for my grade.
5. I will seek help if I need it. I understand everyone gets confused, but Mrs. Tran doesn’t know when to help you if you don’t ask.
6. I will take pride in being a South Hills Scorpion!
7. I choose to make it an awesome year!

What I can expect from Mrs. Tran

1. She will treat me fairly.
2. She is concerned about me and my education.
3. She will update me on my progress and grades.
4. She will expect a lot from me.
5. She will teach me the things I’m willing to learn.
6. She will be honest with me.

Where to go for help
Please visit the classroom for help between these hours:

email address: myscienceclass@yahoo.com

website: http://scienceeinstein.blogspot.com

Lab Safety Contract

Science is a hands-on laboratory class. Students will be doing many laboratory activities that may require the use of chemicals, laboratory equipment, and other items which, if used incorrectly, can be hazardous. Safety in the science classroom is the number 1 priority for students, teachers, and parents. To ensure a safe science classroom, a list of rules has been developed and provided to you in this student safety contract. These rules must be followed at all times. The student and a parent must sign their copy. Please read the entire contract before you sign. Students will not be allowed to perform experiments until all their contracts are signed and given to the teacher.
GENERAL GUIDELINES
1. Conduct yourself in a responsible manner at all times in the classroom.
2. Follow all written and verbal instructions carefully. If you do not understand a direction or part of a procedure, ASK YOUR TEACHER BEFORE PROCEEDING WITH THE ACTIVITY.

3. When first entering a science room, do not touch any equipment, chemicals, or other materials in the laboratory area until you are instructed to do so.
4. Perform only those experiments authorized by your teacher. Carefully follow all instructions, both written and oral. Unauthorized experiments are not allowed.

5. Be prepared for your work in the laboratory. Read all procedures thoroughly before entering the laboratory. Never fool around in the laboratory. Horseplay, practical jokes, and pranks are dangerous and prohibited.

6. Be alert and proceed with caution at all times in the laboratory. Notify the teacher immediately of any unsafe conditions you observe.
7. Keep hands away from face, eyes, mouth, and body while using chemicals or lab equipment. Wash your hands with soap and water after performing all experiments.
8. Experiments must be personally monitored at all times. Do not wander around the room, distract other students, startle other students or interfere with the laboratory experiments of others.




CLOTHING

9. Any time chemicals, heat, or glassware are used, students will wear safety goggles. NO EXCEPTIONS TO THIS RULE!
10. Dress properly during a laboratory activity. Long hair, dangling jewelry, and loose or baggy clothing are a hazard in the laboratory. Long hair must be tied back, and dangling jewelry and baggy clothing must be secured. Shoes must completely cover the foot. No sandals allowed on chemical lab days.
ACCIDENTS AND INJURIES
11. Report any accident (spill, breakage, etc.) or injury (cut, burn, etc.) to the teacher immediately, no matter how trivial it seems. Do not panic.
HANDLING CHEMICALS

12. Do not taste, or smell any chemicals.
13. Do not return unused chemicals to their original container unless specifically instructed by your teacher.
14. Never remove chemicals or other materials from the laboratory area.
QUESTIONS (answers are confidential)
15. Do you wear contact lenses? Yes_______ No______
16. Are you color blind? Yes _______ No______
17. Do you have allergies? Yes _______ No ______
If so, please list specific allergies _____________________________________________________________________________

_____________________________________________________________________________
AGREEMENT
I, __________________________________ (student's name) have read and agree to follow all of the safety rules set forth in this contract. I am aware that any violation of this safety contract that results in unsafe conduct in the laboratory or misbehavior on my part, may result in my being removed from the lab classroom, detention, receiving a failing grade, and/or further disciplinary action.

Student signature Date
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