## How to convert between pH and concentrations.

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Aim: How to convert between pH and concentrations.

Activity: Draw the pH scale.

Notes: Sörenson defined pH as the negative logarithm of the hydrogen ion concentration.

pH = – log [H+]

Remember that sometimes H3O+ is written, so

pH = – log [H3O+]

means the same thing.

So let’s try a simple problem: The [H+] in a solution is measured to be 0.010 M. What is the pH?

The solution is pretty straightforward. Plug the [H+] into the pH definition:

pH = – log 0.010

An alternate way to write this is:

pH = – log 10¯2

Since the log of 10¯2 is -2, we have:

pH = – (- 2)

Which, of course, is 2.

## Angela Shella Photo and notes on How to identify oxidizing and reducing agents

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Aim: How to identify oxidizing and reducing agents

Notes:

oxidizing agent: the substance that gets reduced

reducing agent: the substance that gets oxidized

An atom gets oxidized when the oxidation number increases.

An atom gets reduced when its oxidation number decreases

## How to write half reactions

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AIM: How to write half reaction

NOTES:(from chemteam)A half-reaction is simply one which shows either reduction OR oxidation, but not both. Here is the example redox reaction used before:

Ag+ + Cu —> Ag + Cu2+

It has BOTH a reduction and an oxidation in it. That is why we call it a redox reaction, from REDuction and OXidation.

What you must be able to do is look at a redox reaction and separate out the two half-reactions in it. To do that, identify the atoms which get reduced and get oxidized. Here are the two half-reactions from the above example:

Ag+ —> Ag

Cu —> Cu2+

The silver is being reduced, its oxidation number going from +1 to zero. The copper’s oxidation number went from zero to +2, so it was oxidized in the reaction. In order to figure out the half-reactions, you MUST be able to calculate the oxidation number of an atom.

Keep in mind that a half-reaction shows only one of the two behaviors we are studying. A single half-reaction will show ONLY reduction or ONLY oxidation, never both in the same equation.

Also, notice that the reaction is read from left to right to determine if it is reduction or oxidation. If you read the reaction in the opposite direction (from right to left) it then becomes the other of our two choices (reduction or oxidation). For example, the silver half-reaction above is a reduction, but in the reverse direction it is an oxidation, going from zero on the right to +1 on the left.

There will be times when you want to switch a half-reaction from one of the two types to the other. In that case, rewrite the entire equation and swap sides for everything involved. If I needed the silver half-reaction to be oxidation, I’d write Ag —> Ag+ rather tha just doing it mentally.

OK, back to the next step, which is both half-reactions must be balanced. However, there is a twist. When you learned about balancing equation, you made equal the number of atoms of each element on each side of the arrow. That still applies, but there is one more thing: the total amount of charge on each side of the half-reaction MUST be the same.

When you look at the two half-reactions above, you will see they are already balanced for atoms with one Ag on each side and one Cu on each side. So, all we need to do is balance the charge. To do this you add electrons to the more positive side. You add enough to make the total charge on each side become EQUAL.

To the silver half-reaction, we add one electron:

Ag+ + e¯ —> Ag

To the copper half-reaction, we add two electrons:

Cu —> Cu2+ + 2e¯

One point of concern: notice that each half-reaction wound up with a total charge of zero on each side. This is not always the case. You need to strive to get the total charge on each side EQUAL, not zero.

One more point to make before wrapping this up. A half-reaction is a “fake” chemical reaction. It’s just a bookkeeping exercise. Half-reactions NEVER occur alone. If a reduction half-reaction is actually happening (say in a beaker in front of you), then an oxidation reaction is also occuring. The two half-reactions can be in separate containers, but they do have to have some type of “chemical connection” between them. The nature of this connection is the subject of another tutorial.

ASSESSMENT: Given reactions, write the half reactions.

Sn + NO3¯ —> SnO2 + NO2

HClO + Co —> Cl2 + Co2+

NO2 —> NO3¯ + NO

## How to use the mole concept?

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The mole is the standard method in chemistry for communicating how much of a substance is present.

Here is how the International Union of Pure and Applied Chemistry (IUPAC) defines “mole:”

The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12. When the mole is used, the elementary entities must be specified and may be atoms, molecules, ions, electrons, other particles, or specified groups of such particles.

This is the fundamental definition of what one mole is. One mole contains as many entities as there are in 12 grams of carbon-12 (or 0.012 kilogram).

In one mole, there are $6.022 x 10^23$ atoms. Here’s another way: there are 6.022 x 1023 atoms of carbon in 12 grams of carbon-12.

Let’s say that real clearly: one mole of ANYTHING contains 6.022 x 1023 entities.

The word “entities” is simply a generic word. For example, if we were discussing atoms, then we would use “atoms” and if molecules were the subject of discussion, the word entities would be replaced in actual use by “molecules.”

Avogadro’s Number has been very carefully measured in a number of ways over many decades. The symbol for mole is “mol.” Why does a four-letter word have a three-letter symbol? That’s really the wrong question. Here’s why.

Here it is again: one mole of ANY specified entity contains 6.022 x 1023 of that entity. For example:

• One mole of donuts contains 6.022 x 1023 donuts
• One mole of H2O contains 6.022 x 1023 molecules
• One mole of nails contains 6.022 x 1023 nails
• One mole of Fe contains 6.022 x 1023 atoms
• One mole of dogs contains 6.022 x 1023 dogs
• One mole of electrons contains 6.022 x 1023 electrons

Get the idea?

6.022 x 1023 is so important in chemistry that it has a name. It is called Avogadro’s Number and has the symbol N. It is so named in honor of Amedeo Avogadro, an Italian chemist, who, in 1811, made a critical contribution (recognized only in 1860 after his death) which helped greatly with the measurement of atomic weights.

Avogadro’s Number has a unit associated with it. It is mol¯1, as in 6.022 x 1023 mol¯1. The superscripted minus one means the unit mol is in the denominator. There is an understood numerator of one, as in 1/mol.

Why is there no unit in the numerator? There could be, but it would vary based on the entity involved. If we were discussing an element, we might write atoms/mol. If we were discussing a compound, we would say “molecules per mol.” What is in the numerator depends on what “entity” (atom, molecule, ion, electron, etc.) is being used in the problem.

Consequently, units names in the numerator are not used and a one is used instead.

Getting back to Avogadro’s Number role in chemistry; please note that counting atoms or molecules is very difficult since they are so small. However, we can “count” atoms or molecules by weighing large amounts of them on a balance.

When we weigh one mole of a substance on a balance, this is called a “molar mass” and has the units g/mol (grams per mole). This idea is very critical because it is used all the time.

• A molar mass is the weight in grams of one mole.
• One mole contains 6.022 x 1023 entities.

Therefore, a molar mass is the mass in grams of 6.022 x 1023 entities.

OK. How does one calculate a molar mass? Get ready, because you already know how to calculate a molar mass.

The molar mass of a substance is the molecular weight in grams.

All you need to do is calculate the molecular weight and stick the unit “g/mol” after the number and that is the molar mass for the substance in question.

Calculate the molar mass of Al(NO3)3

(1 x 26.98) + (3 x 14.007) + (9 x 16.00) = 213.00 g/mol

213.00 grams is the mass of one mole of aluminum nitrate.

213.00 grams of aluminum nitrate contains 6.022 x 1023 entities of Al(NO3)3

## How to perform and predict precipitation reactions.(LAB)

Double Replacement Reactions (Precipitation Reactions)

Explain that no matter how substances within a closed system interact with one another, or how they combine or break apart, the total mass of the system remains the same.  Understand that the atomic theory explains the conservation of matter: if the number of atoms stays the same no matter how they are rearranged, then their total mass stays the same.

Objectives

1 Recognize when a chemical reaction has taken place (as opposed to a physical change)
2 Observe and recognize when a precipitate has formed
3 Recognize a double replacement reaction
4 Write the equation for a chemical reaction

Required Materials

For 100 students 100 mL each of :        ( in plastic dropper bottles)

0.1M NaCl (twice as many bottles)
0.1MKNO3
0.1M AgNO3 (twice as many bottles)
0.1 M NaI
0.1 M Cu(NO3)2
3  M NaOH (twice as many bottles)
0.1 M Fe(NO3)3
distilled water
gloves
5 test tubes / pair of students

Lesson Introduction

Begin by cautioning students about safety.  Have the students wear gloves and goggles.  There should be no eating or drinking in lab.  Make sure no students aim test tubes at other students.  Caution students to keep chemicals away from face.  Clean up spills immediately.  The students should wash their hands after the experiment.  No horseplay is allowed.

On day 1:  (This is a 2-day exercise.) Explain to students that a chemical change has taken place when substances are used up and others are formed to take their place.  Substances have changed identity.  In a physical change like melting or boiling, substances don’t change their identity.  Their chemical formula doesn’t change.  In a chemical reaction no substance is lost, but the substance changes by rearrangement of atoms so that the total mass stays the same.

All the compounds in this exercise are inorganic and ionic, meaning that the substance can dissociate in water to form a cation and an anion (species with a  charge).  Table salt can be written NaCl.  In water it is actually Na+(aq) and Cl-(aq).  These ions can recombine with other substances to form a different salt- some soluble and some insoluble.  Insoluble substances that form in solution are called precipitates.  For example NaCl + AgNO3 can dissolve and recombine in water to form NaNO3 + AgCl.  NaNO3 is soluble and will remain dissolved in the water, but AgCl is mostly insoluble and will precipitate out- or form a white solid.  It might be useful also to show the students the list of polyatomic anions on the worksheet so that they can see that NO3- remains in solution in a whole unit as does SO42- and OH-.  Also show the students the list of substances that form precipitates in water.  Explain that these substances have a limited solubility in water so that they fall (precipitate) out of the water as particles.  Work through several examples (several are suggested below) with explanations. Then allow students to do the worksheet.

On day 2 during this exercise students will try in lab five combinations of ionic solutions, decide whether or not a precipitate has formed and write the chemical equation for the reaction.  They will need to practice writing equations on Day 1.  Show students that the prototype double displacement reaction is AB + CD        AD + CB.  Show them examples and discuss:
Example 1:
NaCl(aq) + AgNO3(aq)     NaNO3(aq) + AgCl (s)

Or

Na+(aq) + Cl-(aq) + Ag+(aq) + NO3-(aq)        Na+(aq) + NO3-(aq) + AgCl (s)

Example 2:
K2SO4 + NaOH     no precipitation reaction  (Note that double replacement doesn’t result in an insoluble salt.)
or
2K+(aq) + SO42-(aq) + Na+(aq) + OH-(aq)    2K+(aq) + SO42-(aq) + Na+(aq) + OH- (aq)

Example 3:
Al2(SO4)3 (aq)  +6 KOH (aq)        2 Al(OH)3 (s) + 2K+(aq) + SO42-(aq)

Example 4:
LiCl (aq) + AgNO3 (aq)        AgCl (s) + Li+(aq) +NO3-(aq)

Example 5:
Pb(NO3)2 (aq) + 2 NaI (aq)        PbI2(s) + Na+ (aq) + NO3- (aq)

Example 6:
CaCl2(aq) + Na2CO3    (aq)        CaCO3 (s) + 2 NaCl (aq)

Students should then be allowed to try out what they have learned from the discussion and examples on Worksheet I.  Provide students with the list of polyatomic ions and insoluble salts.  A list follows worksheet 2.

On Day 2:
Procedure

Review lab safety.  Students should mix 10 drops of each of the two solutions that are paired on the worksheet in a test tube.  They should mix well and record observations.  If a reaction occurred they should write a complete and balanced equation.  If no reaction occurred, they should record “no reaction.”  Discard the mixtures from the test tubes in a waste jar.  Students should not pour their solutions in the sink.  All of the waste may be later diluted by the teacher and discarded, except for AgCl and Cu(OH)2.  Consult local authorities for disposal of these two precipitates.

Closure

Explain again that these are called double replacement reactions.  They can be used to identify the presence of various ions in solution.  Often wastewater is studied this way.

Assessment

The worksheet may be graded along with the students’ observations.  An unknown solution containing the ions the students have used may be prepared.  The students may be asked to identify their unknown solution containing a metal by testing with the other solutions KNO3, AgNO3, H2SO4, K2CrO4 or NaOH.

Worksheet 1
Name______________________________
Class Period________________________

Complete the following equations.  Note that precipitates are insoluble and are followed by (s).  Species in solution are followed by (aq).  Note the list of insoluble salts.  These are precipitates.  Note the list of polyatomic ions.  These atoms always stay together as a unit.

1.  Ba(NO3)2 (aq) + K2SO4 (aq)

2.  AgNO3 (aq) + NaBr (aq)

3.  FeCl3 (aq) + 3 KOH (aq)

4.  Pb(NO3)2 (aq) + K2SO4 (aq)

5.  Cu(NO3)2 (aq)  + 2 NaOH

Polyatomic Ions

NH4+ ammonium
OH-  hydroxide
NO3- nitrate
CO32- carbonate
SO42- sulfate
PO43- phosphate

Worksheet 2
Name______________________________
Class Period________________________

Add 10 drops of each reagent in the equation to a test tube.  Mix well.  Record your observations.

Reactions:

1.  _____NaCl(aq)  + _____ KNO3(aq)
Observations:

2. _____NaCl(aq) + _____AgNO3 (aq)
Observations:

3. _____AgNO3(aq)  +  _____NaI (aq)
Observations:

4.  _____Cu(NO3)2 + ______NaOH
Observations:

5.  _____Fe(NO3)3 (aq)  + _____NaOH(aq)
Observations:

List of Insoluble Salts

AgCl          silver chloride (white)
Ag2CrO4    silver chromate (red)
AgIO3         silver iodate (white)
AgI        silver iodide (yellow)
BaSO4        barium sulfate (white)
Cu(OH)2    copper(II) hydroxide (blue)
Fe(OH)3    iron(III) hydroxide (red)

Polyatomic Ions

NH4+ ammonium
OH-  hydroxide
NO3- nitrate
CO32- carbonate
SO42- sulfate
PO43- phosphate

1.  Ba(NO3)2 (aq) + K2SO4 (aq)         BaSO4 (s) + 2 K+ (aq) + 2 NO3- (aq)

2.  AgNO3 (aq) + NaBr (aq)         AgBr (s) + Na+ (aq) + NO3- (aq)

3. FeCl3 (aq) + 3 KOH (aq)        Fe(OH)3 (s) + 3 K+ (aq) + 3 Cl- (aq)

4.  Pb (NO3)2 (aq) + K2SO4 (aq)         PbSO4 (s) + 2 K+ (aq) + 2 NO3-(aq)

5.  Cu(NO3)2 (aq) + 2 NaOH     Cu(OH)2 (s) + 2 Na+ (aq) + 2 NO3- (aq)

1.  1 NaCl(aq)  + 1 KNO3(aq)         No reaction
Observations:
No reaction

2.  1 NaCl(aq) + 1 AgNO3 (aq)        AgCl (s) + Na+ (aq) + NO3- (aq)
Observations:
White precipitate

3.  1 AgNO3(aq)  +  1 NaI (aq)        AgI (s) + Na+ (aq) + NO3- (aq)
Observations:
Yellow precipitate

4. 1  Cu(NO3)2 + 2 NaOH        Cu (OH)2 (s) + 2 Na+ (aq) + 2 NO3- (aq)
Observations:
Blue precipitate

5.  1 Fe(NO3)3 (aq)  + 3 NaOH(aq)          Fe (OH)3 + 3 Na+ (aq) + 3 NO3- (aq)
Observations:
Red precipitate

## How to Balance Chemical Reactions

Chemical equations do not come already balanced. This must be done before the equation can be used in a chemically meaningful way.

All chemical calculations to come must be done with a balanced equation.

A balanced equation has equal numbers of each type of atom on each side of the equation.

The The Law of Conservation of Mass is the rationale for balancing a chemical equation. The law was discovered by Antoine Laurent Lavoisier (1743-94) and this is his formulation of it, translated into English in 1790 from the Traité élémentaire de Chimie (which was published in 1789):

“We may lay it down as an incontestible axiom, that, in all the operations of art and nature, nothing is created; an equal quantity of matter exists both before and after the experiment; the quality and quantity of the elements remain precisely the same; and nothing takes place beyond changes and modifications in the combination of these elements.”

A less wordy way to say it might be:

“Matter is neither created nor destroyed.”

Therefore, we must finish our chemical reaction with as many atoms of each element as when we started.

Here is the example equation for this lesson:

H2 + O2 —> H2O

It is an unbalanced equation (sometimes also called a skeleton equation). This means that there are UNEQUAL numbers at least one atom on each side of the arrow.

In the example equation, there are two atoms of hydrogen on each side, BUT there are two atoms of oxygen on the left side and only one on the right side.

Remember this: A balanced equation MUST have EQUAL numbers of EACH type of atom on BOTH sides of the arrow.

An equation is balanced by changing coefficients in a somewhat trial-and-error fashion. It is important to note that only the coefficients can be changed, NEVER a subscript.

The coefficient times the subscript gives the total number of atoms.

Three quick examples before balancing the equation.

(a) 2 H2 – there are 2 x 2 atoms of hydrogen (a total of 4).

(b) 2 H2O – there are 2 x 2 atoms of hydrogen (a total of 4) and 2 x 1 atoms of oxygen (a total of 2).

(c) 2 (NH4)2S – there are 2 x 1 x 2 atoms of nitrogen (a total of 4), there are 2 x 4 x 2 atoms of hydrogen (a total of 16), and 2 x 1 atoms of sulfur (a total of 2).

So, now to balancing the example equation:

H2 + O2 —> H2O

The hydrogen are balanced, but the oxygens are not. We have to get both balanced. We put a two in front of the water and this balances the oxygen.

H2 + O2 —> 2 H2

O

However, this causes the hydrogen to become unbalanced. To fix this, we place a two in front of the hydrogen on the left side.

2 H2 + O2 —> 2 H2O

This balances the equation.

Two things you CANNOT do when balancing an equation.

1) You cannot change a subscript.

You cannot change the oxygen’s subscript in water from one to two, as in:

H2 + O2 —> H2O2

True, this balances the equation, but you have changed the substances in it. H2O2 is a completely different substance from H2O.

2) You cannot place a coefficient in the middle of a formula.

The coefficient goes at the beginning of a formula, not in the middle, as in:

H2 + O2 —> H22O

Water only comes as H2O and you can only use whole formula units of it.

There is another thing you should avoid. Make sure that your final set of coefficients are all whole numbers with no common factors other than one. For example, this equation is balanced:

4 H2 + 2 O2 —> 4 H2O

However, all the coefficients have the common factor of two. Divide through to eliminate common factors like this.

The equation just above is correctly balanced, but it is not the BEST answer. The best answer has all common factors greater than one removed.

Balance this equation: H2 + Cl2 —> HCl

Remember that the rule is: A balanced equation MUST have EQUAL numbers of EACH type of atom on BOTH sides of the arrow.

The correctly balanced equation is:

H2 + Cl2 —> 2 HCl

Placement of a two in front of the HCl balances the hydrogen and chlorine at the same time.

Balance this equation: O2 —> O3

Hint: think about what the least common multiple is between 2 and 3. That’s right – six.

The LCM tells you how many of each atom will be needed. Your job is to pick coefficients that get you to the LCM.

The correctly balanced equation is:

3 O2 —> 2 O3

### Practice Problems

How many oxygens are indicated: 3 Ca(NO3)2

Balance these equations:

Zn + HCl —> ZnCl2 + H2

KClO3 —> KCl + O2

S8 + F2 —> SF6

Fe + O2 —> Fe2O3

### Balancing Worksheet #1

Please note that several of these equations are already balanced as written. The answers are in this file and are several lines below the last problem. There are 50 problems in two columns.

` 1.  H2 +  O2 --->  H2O				26.  N2   +   H2  --->   NH3 2.  S8 +  O2 --->  SO3				27.  N2   +   O2  --->   N2O 3.  HgO --->  Hg +  O2				28.  CO2   +   H2O  --->   C6H12O6   +   O2 4.  Zn +  HCl --->  ZnCl2 +  H2			29.  SiCl4   +   H2O  --->   H4SiO4   +   HCl 5.  Na +  H2O --->  NaOH +  H2			30.  H3PO4  --->   H4P2O7   +   H2O 6.  C10H16 +  Cl2 --->  C +  HCl		31.  CO2   +   NH3  --->   OC(NH2)2   +   H2O 7.  Si2H3   +   O2  --->   SiO2   +   H2O	32.  Al(OH)3   +   H2SO4  --->   Al2(SO4)3   +   H2O 8.  Fe +  O2 --->  Fe2O3			33.  Fe2(SO4)3   +   KOH  --->   K2SO4   +   Fe(OH)3 9.  C7H6O2 +  O2 --->  CO2 +  H2O		34.  H2SO4   +   HI   --->   H2S   +   I2   +   H2O10.  FeS2 +  O2 --->  Fe2O3 +  SO2		35.  Al   +   FeO   --->   Al2O3   +   Fe11.  Fe2O3  +  H2  --->   Fe   +   H2O		36.  Na2CO3   +   HCl   --->   NaCl   +   H2O   +   CO212.  K   +   Br2  --->    KBr			37.  P4   +   O2   --->   P2O513.  C2H2   +   O2  --->   CO2   +   H2O		38.  K2O   +   H2O   --->   KOH14.  H2O2   --->   H2O   +   O2			39.  Al   +   O2   --->   Al2O315.  C7H16   +    O2   --->   CO2    +   H2O	40.  Na2O2   +   H2O   --->   NaOH   +   O216.  SiO2   +   HF  --->   SiF4    +    H2O	41.  C   +   H2O   --->   CO   +   H217.  KClO3   --->   KCl    +   O2		42.  H3AsO4   --->   As2O5   +   H2O18.  KClO3   --->   KClO4   +   KCl		43.  Al2(SO4)3 +  Ca(OH)2 --->  Al(OH)3 +  CaSO419.  P4O10    +   H2O   --->   H3PO4		44.  FeCl3   +   NH4OH   --->   Fe(OH)3   +   NH4Cl20.  Sb   +   O2   --->   Sb4O6			45.  Ca3(PO4)2   +  6 SiO2  --->   P4O10  +  CaSiO321.  C3H8   +   O2   --->   CO2    +   H2O	46.  N2O5   +   H2O   --->   HNO322.  Fe2O3   +    CO   --->   Fe   +   CO2	47.  Al   +   HCl   --->   AlCl3   +   H223.  PCl5    +   H2O   --->   HCl   +   H3PO4	48.  H3BO3   --->   H4B6O11    +    H2O24.  H2S   +   Cl2   --->   S8   +   HCl		49.  Mg   +   N2   --->    Mg3N225.  Fe   +   H2O   --->   Fe3O4    +   H2	50.  NaOH   +   Cl2   --->   NaCl   +   NaClO   +   H2O`

## How to Calculate Formula Weight

`1. AlCl3 	14. Ba(SCN)2	27. LiH		40. Ba(BrO3)2	53. AlBr3	66. HCl2. TeF4		15. K2S		28. CO		41. Hg2Cl2	54. P2O5	67. K2SO43. PbS 		16. NH4Cl	29. SnI4	42. Cr2(SO3)3	55. NH4NO3	68. NaCl4. Cu2O		17. KH2PO4	30. KOH		43. Al(MnO4)3	56. Ba(OH)2	69. LiI 5. AgI 		18. C2H5NBr	31. K2O		44. CoSO4	57. PbSO4	70. Hg2O6. N2O		19. Ba(ClO3)2	32. H2SO4	45. Ca(NO3)3	58. Ba3(PO4)2	71. HF7. MoCl5 	20. Fe(OH)3	33. Hg3N2	46. NaH2PO4	59. NaC2H3O2	72. FeCl38. Hg2Br2 	21. (NH4)2S	34. SiF4	47. (NH4)3PO4	60. Ba(OH)2	73. NaHSO49. Ta2O5	22. CoCl2	35. NH4OH	48. KAl(SO4)2	61. NaHCO3	74. Ag2O10. HgF2	23. KMnO4	36. N2O5	49. Hg2SO4	62. Al(OH)3	75. Pb(ClO2)211. KCl		24. CaSO4	37. SnCrO4	50. Al2(SO4)3	63. NH4MnO4	76. CoF312. KF		25. H2CO3	38. Al2O3	51. FePO4	64. Fe2O3 	77. Al(C2H3O2)313. ZnO		26. CO2		39. CuCO3	52. Ca(C2H3O2)2	65. CaCO3	78. Na2Al2(SO4)4Answers (each answer has the units g/mol)1. 133.34 	14. 255.26	27. 7.95	40. 393.1314	53. 266.69	66. 36.4612. 203.59	15. 110.26	28. 28.01	41. 472.09	54. 141.944	67. 174.253. 239.3 	16. 53.49	29. 626.31	42. 344.1666	55. 80.04	68. 58.4434. 143.09	17. 136.08	30. 56.106	43. 383.788	56. 171.34	69. 133.846 5. 234.77 	18. 122.97	31. 94.20	44. 154.99	57. 303.26	70. 417.1796. 44.01	19. 304.23	32. 98.07	45. 226.09	58. 601.93	71. 20.0067. 273.20 	20. 106.87	33. 629.78	46. 119.977	59. 82.03	72. 162.2068. 560.98	21. 68.14	34. 104.08	47. 149.087	60. 171.34	73. 120.0559. 441.89	22. 129.84	35. 35.046	48. 258.195	61. 84.007	74. 231.7410. 238.59	23. 158.03	36. 108.01	49. 497.24	62. 78.00	75. 342.1011. 74.55	24. 136.14	37. 234.68	50. 342.136	63. 136.97	76. 115.92812. 58.10	25. 62.02	38. 101.96	51. 150.82	64. 159.69 	77. 204.1213. 81.38	26. 44.01	39. 123.555	52. 158.169	65. 100.09	78. 484.173`

## How to assign oxidation numbers

Rule Number One: All free, uncombined elements have an oxidation number of zero.This includes diatomic elements such as O2 or others like P4 and S8.

Rule Number two: Hydrogen, in all its compounds except hydrides, has an oxidation number of +1 (positive one)

Rule number three: Oxygen, in all its compounds except peroxides, has an oxidation number of -2 (negative two). another exception is when oxygen is with fluorine in a binary compound.

With only a very few exceptions, oxidation states can be assigned to all atoms in a formula. There are more extensive sets of rules and, for the most part, they derive from the three above rules.

There are some complex examples which are not discussed in the usual set of introductory rules. For example, potassium superoxide is KO2. That means that the oxidation number on the O2 is negative one.

Rule Number 4: all the oxidation numbers in a compound must add up to zero.

Rule Number 5. all the oxidation numbers in a polyatomic ion must add up to the charge of the polyatomic ion.

Rule Number 6. the oxidation number of fluorine in a compound is -1.

Rule Number 7. The oxidation number of the elements of monatomic ions in a compound is the same as the charge of the monatomic ion.

Rule Number 8. In a compound, the oxidation number of group 1 elements is +1.

Rule Number 9. In a compound, the oxidation number of group 2 elements is +2

Practice Examples

## Chemical Bonds and Reactions

Do Now: Students will write the electron configuration of Sodium and chloride. They write a paragraph discussing the octet rule. Why would Na and Cl make a bond?

UnitedStreaming.com  ” elements of Chemistry: Compounds and Reactions”

Aim: How do elements combine in a chemical reaction?

Notes:

Main concepts include changes in matter, ionic bonds, covalent bonds, chemical reactions, and preview of electrochemistry.

There are three different types of changes in matter: nuclear, physical, and chemical changes. When chemical changes occur, there is an actual change in the chemical composition of the substances.  The original atoms are preserved but they combine  in such a way that a new substance is created with a different  chemical composition.

In a chemical compound, the valence electrons are lost, gained, or shared between the different atoms. These unions are called chemical bonds. There are two basic kinds of chemical bonds: ionic bonds and covalent bonds. Electrons have the ability to move between atoms. When a neutral atom loses an electron, it becomes a positive ion and when a neutral atom gains an electron it becomes a negative ion. These ions are then attracted to each other because of their opposite electrical charges.

Covalent bonds occur when electrons are shared between atoms in structures called molecules.  There are a vast number of substances that have molecular structures including plastic, paper, water, and all of the bonds in plants and animals. Molecules can be made up with as few as two atoms or hundreds and even a billion atoms. Molecular formulas are precise descriptions of how many atoms there are in a single molecule.

A chemical reaction is a process in which one or more substances are converted into new substances with different physi-
cal and chemical properties.

Electrochemistry is a special type of chemical reaction called redox reactions. There are a wide variety of redox reactions,
from respiration in animals to the rusting of iron. One of the most interesting is the production of electricity in a battery.

Video Quiz to follow.

Directions: Answer the following either true or false, or fill in the blank with the correct word to
make it true.
1. There are only 92 elements that are found naturally in the universe.
T_______  F_______
2. The process when atoms are broken apart to release energy is called fusion.
T_______  F_______.
3. When there is a change in the chemical composition of a substance, it is called a chemical __________.
4. When an atom loses or gains one or more electron, it becomes an ____________.
5. The octet rule states that atoms tend to gain, lose, or share electrons in order to acquire a full set of
valence electrons.
T_______  F_______
6. When electrons are shared between atoms in covalent bonds, they are called ____________.
7. Molecules never arrange themselves in the same way.
T_______  F_______
8. A chemical reaction is a process in which one or more substances are converted into new substances
with different physical and chemical properties.
T_______  F_______
9. Chemical equations do not need to balance.  There can be different numbers of atoms on each side of
the equation.
T_______  F_______
10. The electric power in a battery is generated by a chemical reaction.
T_______  F_______

## How to Draw Structural Formulas

Aim: How To Draw electron dot notation for bonding in several molecules.
How to Draw Structural Formulas
How to Distinguish between single, double, and triple bonding

Notes: frostburg.edu
1.  Draw a skeleton structure. A skeleton structure is a rough map showing the arrangement of atoms within the molecule. In general, you need to determine the skeleton experimentally, but here are a few guidelines for predicting skeleton structures from molecular formulas.
* Central atoms are usually the atoms with highest valence, or
the largest atoms, or
the least electronegative atom.
* H and the halogens are usually outside atoms.
* Don’t put more than four atoms around a central atom unless the central atom is third period or lower.
2. Count total valence electrons.
* Add the number of electrons in the valence shells of all atoms in the molecule.
* If the molecule is charged, add an electron for each negative charge and subtract an electron for each positive charge.
* Noble gas compounds are very uncommon (except on general chemistry tests!) Should you encounter one, each noble gas atom has 8 valence electrons.
3. Connect the structure.
* Draw a bond between the central atom and each outside atom.
* Each bond uses 2 valence electrons.
4. Place electrons on outside atoms.
* Use remaining electrons to satisfy the octets for each of the outside atoms.
* If you run out of electrons at this point, the skeleton structure was wrong. Go back to step I.
5. Place all remaining electrons on the central atom.
* If there are more than 8 electrons on the central atom, and the central atom is not third period or lower, you counted the number of valence electrons incorrectly. Go back to step II.
* If the octet on the central atom is not complete, try sharing lone pairs of outside atoms to form double or triple bonds. Write one multiply bound structure for each outside atom with a lone pair to share; these are resonance structures.
* If you can’t get an octet on the central atom, at this point, check to see whether the total number of valence electrons for this molecule is odd. It’s impossible to give octets to all atoms in an odd electron molecules. Get as close to an octet as possible by forming multiple bonds.

Assessment:

Draw dot diagrams and structural formulas for HCl,  H2O, CH4, C2H6, CO2, NH3, CS2, H2S, HBr, PH3

Homework or Summary

Vocab definitions

draw lewis structures

## Exploring Different Forms of Energy

Aim:  How to connect real world examples of different Forms of Energy to mathematical formulas.

Unitedstreaming.com   “elements of physics: energy work and power”

This video demonstrates how the formula W = Fx is a definition of energy, explains kinetic energy and givesI a number of examples. This is a great review because it also ties in potential energy and gives a number of examples. It illustrates and  explains some of the types of energy such as electromagnetism, nuclear energy, chemical energy, sound energy, and heat energy. Students will see real life examples of energy exchanges and  see how the three laws of thermodynamics apply to real world scenarios. This video also discusses why Einstein’s formula E = mc 2  explains how matter and energy are different aspects of the same thing.This is a great overall review that ties the energy concept together.

Video Quiz follows.

Review for Test(unit study guide)

Notes

1.       Distinguish between electrical, chemical, radiant, and thermal energy.   What do they all have in common?

a)Electrical – energy associated with the movement of charged particles

b)Chemical – energy stored in chemical bonds

c)Radiant – energy carried by electromagnetic waves

d)Thermal – kinetic & potential energy of the particles in an object

e) All of these forms of energy have the ability to do work.

2.      What is kinetic energy?   What is potential energy?   Describe how a swinging pendulum demonstrates the back-and-forth conversion of potential energy to kinetic energy.

Kinetic energy is energy associated with moving objects, no matter how big or how small.   Potential energy is energy stored in an object due to is position and location.

A swinging pendulum demonstrates both kinetic energy and potential energy.   When a swinging pendulum is at the highest point in its swing, it has maximal potential energy and a minimal kinetic energy.   At the bottom or lowest point in its swing, it has maximal kinetic energy and minimal potential energy.

3.      Describe and give an example of the following types of potential energy…

a)      elastic potential energy – energy stored when an object is compressed or stretched, such as a stretched rubber band

b)     chemical potential energy – energy stored in chemical bonds which can be released by a chemical reaction, such as the energy stored in a typical household battery

c)      electrical potential energy – energy due to the position of an electrical charge relative to other electrical charges, such as static electricity

d)     nuclear potential energy –energy stored in the nuclei of atoms and released during nuclear reactions;   the energy produced at nuclear power plants are an example

4.     Calculate the gravitational potential energy of a 500-kg rock resting on the edge of a 200-meter cliff.

Gravitational potential energy   =   mass   x   gravity   x   height

GPE   =   500 kg   x   9.8 m/s2   x   200 meters

GPE   =   980,000 Joule

5.             Calculate the kinetic energy the rock mentioned in question 4 will have just before it hits the ground if it is moving at 10 meters per second.

Kinetic energy   = ½( m   x   v2)

KE   =   ½   x   [500 kg   x (10 m/s)2]

KE   =   2500 Joules

Homework:

Review for Test

## How to Name Compounds with Polyatomic Ions

Aim: How to Name compounds with polyatomic ions.

Materials: Index Cards, markers

Background information: REVIEW: (These compounds to follow ARE NOT binary compounds. They contain three or more elements, as opposed to only two in a binary compound. Don’t use the Greek method. That naming technique is used only for binary compounds of two nonmetals. That means, if you see a formula like BaSO4, the name is not barium monosulfur tetraoxide. Many unaware students over the years have made this error and suffered for it.)

Goals: students should be able to learn to recognize the presence of a polyatomic ion in a formula.

Activity:

students will make a set of flashcards with the name on one side and the ion and its charge on the other. Then, carry them everywhere and use them. The cations used will be a mix of fixed charges AND variable charges. You must know which are which.

Students will associate the charges with each polyatomic ion. For example, NO3¯ is called nitrate and it has a minus one charge.

Notes: glencoe chemistry,  zumdahl & zumdahl

1. When more than one polyatomic ion is required, parenthesis are used to enclose the ion with the subscript going outside the parenthesis. For example, the very first formula used is Fe(NO3)2. This means that two NO3¯ are involved in the compound. Without the parenthesis, the formula would be FeNO32, a far cry from the correct formula.

2. How to say a formula. When you speak a formula involving parenthesis out loud, you use the word “taken” as in the formula for ammonium sulfide, which is (NH4)2S. Out loud, you say “N H four taken twice S.” OR with the formula for copper(II) chlorate, which is Cu(ClO3)2. You say ” Cu Cl O three taken twice.”

Example #1 – write the name for Fe(NO3)2

Step #1 – decide if the cation is one showing variable charge. If so, a Roman numeral will be needed. In this case, iron does show variable charge.

If a variable charge cation is involved, you must determine the Roman numeral involved. You do this by computing the total charge contributed by the polyatomic ion. In this case, NO3¯ has a minus one charge and there are two of them, making a total of minus 2.

Therefore, the iron must be a positive two, in order to keep the total charge of the formula at zero.

Step #2 – determine the name of the polyatomic ion. Nitrate is the name of NO3¯.

The correct name is iron(II) nitrate. The common name would be ferrous nitrate.

Example #2 – write the name for NaOH

Step #1 – the cation, Na+, does not show a variable charge, so no Roman numeral is needed. The name is sodium.

Step#2 – OH¯ is recognized as the hydroxide ion.

The name of this compound is sodium hydroxide.

There are three things you must memorize: the name (hydroxide), the symbol (OH) and the charge (minus one).

Example #3 – write the name for KMnO4

Step #1 – the cation, K+, does not show a variable charge, so no Roman numeral is needed. The name is potassium.

Step#2 – MnO4¯ is recognized as the permanganate ion.

The name of this compound is potassium permanganate.

Example #4 – write the name for Cu2SO4

Step #1 – decide if the cation is one showing variable charge. If so, a Roman numeral will be needed. In this case, copper does show variable charge.

If a variable charge cation is involved, you must determine the Roman numeral involved. You do this by computing the total charge contributed by the polyatomic ion. In this case, SO42¯ has a minus two charge and there is only one, making a total of minus 2.

Therefore, the copper must be a positive one. Why? Well, there must be a positive two to go with the negative two in order to make zero. Since the formula shows two copper atoms involved, each must be a plus one charge.

Step #2 – determine the name of the polyatomic ion. Sulfate is the name of SO42¯.

The correct name is copper(I) sulfate. The common name would be cuprous sulfate.

Example #5 – write the name for Ca(ClO3)2

The first part of the name comes from the first element’s name: calcium. You also determine that it is not a cation of variable charge.

The second part of the name comes from the name of the polyatomic ion: chlorate.

This compound is named calcium chlorate.

Example #6 – write the name for Fe(OH)3

Iron is an element with two possible oxidation states. The iron is a +3 charge because (1) there are three hydroxides, (2) hydroxide is a minus one charge, (3) this gives a total charge of negative three and (40 there is only one iron, so it must be a +3.

Therefore the first part of the name is iron(III).

The second part of the name is hydroxide, the name of the polyatomic ion.

The name of this compound is iron(III) hydroxide (or ferric hydroxide when using the common method).

Home work:

The cations in this first set are all of fixed oxidation state, so no Roman numerals are needed.

Write the correct name for:

1) AlPO4

2) KNO2

3) NaHCO3

4) CaCO3

5) Mg(OH)2

6) Na2CrO4

7) Ba(CN)2

8) K2SO4

9) NaH2PO4

10) NH4NO3

These formulas involve the use of a polyatomic ion. The cations are all of variable oxidation state, so Roman numerals are needed.

Write the correct name for:

11) Sn(NO3)2

12) FePO4

13) Cu2SO4

14) Ni(C2H3O2)2

15) HgCO3

16) Pb(OH)4

17) Cu2Cr2O7

18) Cu(ClO3)2

19) FeSO4

20) Hg2(ClO4)2

These formulas mix the use of the two types of cations.

Write the correct name for:

21) KClO3

22) SnSO4

23) Al(MnO4)3

24) Pb(NO3)2

25) Mg3(PO4)2

26) CuH2PO4

27) CaHPO4

28) Fe(HCO3)3

29) Na2CO3

30) MnSO4

Homework: handout completion of chemical formulas  or (p224 #19-23) ion pairs

## How to Name Binary Molecular Compounds

Aim: How to name binary molecular compounds, when 2 nonmetals are involved.

Notes: glencoe science assessment

1. A binary compound is one made of two different elements. There can
be one of each element such as in CO or NO. There can also be several
of each element such as BF or OCl2.

2. In this type of compound there are no cations!

3. You dont have to know the charges. Just use the element names and their prefixes. Be aware that
heavy use of Greek number prefixes are used in this lesson.Here are the
first ten:

`		one	mono-			six	hexa-		two	di-			seven	hepta-		three	tri-			eight	octa-		four	tetra-			nine	nona-		five	penta-			ten	deca-`

Example #1 – write the name for N2O.

Example #2 – write the name for NO2.

Step #1 – part of the first name is the unchanged name of the first
element in the formula. In the examples above, it would be nitrogen.

If the subscript of the first element is 2 or more, you add a prefix
to the name. In the first example above, you would write dinitrogen. If
the subscript is one as in the second example above, you DO NOT use a prefix. You simply write the name, in this example it would be nitrogen.

Step #2 – the anion is named in the usual manner of stem plus “ide.”
In addition, a prefix is added. In the first example, the prefix is
“mono-” since there is one oxygen. In the second example, use “di-”
because of two oxygens.

The correct names of the two examples are dinitrogen monoxide and nitrogen dioxide.

Note that “monoxide” is written rather than “monooxide.” It sounds better when spoken out loud.

Example #3 – write the name for IF7.

Step #1 – the first element is iodine and there is only one. This part of the name will be “iodine”, NOT “monoiodine.”

Step #2 – the second element is fluorine, so “fluoride” is used. Since there are seven, the prefix “hepta” is used.

The name of this compound is iodine heptafluoride.

Example #4 – write the name for N2O5.

Step #1 – the first element is nitrogen and there are two. This part of the name will be “dinitrogen.”

Step #2 – the second element is oxygen, so “oxide” is used. Since there are five, the prefix “penta” is used.

The name of this compound is “dinitrogen pentaoxide.” Many write is
as “dinitrogen pentoxide.” The ChemTeam believes that both are
considered correct, but the second is to be prefered.

Example #5 – write the name for XeF2.

The first part of the name comes from the first element’s name: xenon. Since there is only one atom present, no prefix is used.

The second part of the name comes from the root of the second symbol
plus ‘ide’ as well as the prefix “di-,”therefore di + fluor + ide =
difluoride.

This compound is named xenon difluoride.

Example #6 – write the name for N2O4.

The first part of the name comes from the first element’s name:
nitrogen. Since there are two atoms, the prefix “di-” is used giving
dinitrogen.

The second part of the name comes from the root of the second symbol
plus ‘ide’ as well as the prefix “tetra-,”therefore tetr + ox + ide =
tetroxide.

This compound is named dinitrogen tetroxide. Notice the dropping of the “a” in tetra.

Just a reminder: this system of naming does not really have an
offically accepted name, but is often called the Greek system (or
method). It involves use of Greek prefixes when naming binary compounds
of two nonmetals.

Sometimes you will see the Stock system applied to these types of
compounds. Here is what the IUPAC currently says about that practice:
“The Stock notation can be applied to both cations and anions, but
preferably should not be applied to compounds between nonmetals.”

Acitivity
Write the correct name for:

1) As4O10

2) BrO3

3) BN

4) N2O3

5) NI3

6) SF6

7) XeF4

8) PCl3

9) CO

10) PCl5

Write the correct name for:

11) P2O5

12) S2Cl2

13) ICl2

14) SO2

15) P4O10

16) UF6

17) OF2

18) ClO2

19) SiO2

20) BF3

Write the correct name for:

21) N2S5

22) CO2

23) SO3

24) XeF6

25) KrF2

26) BrCl5

27) SCl4

28) PF3

29) XeO3

30) OsO4

## How to draw orbital diagrams?

Aim:     How to draw orbital diagrams?

Step 1. Count the number of electrons in the species.

Step 2.   Use the Aufbau principle to write the electron configuration using the correct order.

Step 3. Use boxes to represent orbitals in the   sublevels.

Step 4. Each box can hold a maximum of 2 electrons. Show electrons as arrows.

Step 5.   Hund’s Rule: Don’t fill up an orbital in a sublevel at a specific energy until each orbital is half filled. (empty bus seat rule.)

Step 6. Electrons in the same orbital must have opposite spins: Pauli exclusion principle.

## Atoms emit light

Do Now: Identify group and period properties of the periodic table.
Write the electron configuration of  O, C, N, P, H, He, Ba, Li, Ne, Na, B, using the periodic table.

Aim: How to Calculate wavelength, frequency, and energy of electron emissions.

materials needed: reference tables – bohr model
Notes:
Using spectroscopy to analyze electron arrangement.

(If  time permits)(

What energy level transition is indicated when the light emitted by a hydrogen atom has a wavelength of 103 nm?
n=3 to n=1
Which color of light would a hydrogen atom emit when an electron changes from the n=5 level to the n=2 level?
blue
Electrons release energy as they return to the ground state( excited to ground: high to low)

What is the formula that relates frequency with wavelength,lamda?
According to the visble light spectrum(reference table), which color light has the highest frequency.
Which has the highest wavelength?

Calculating energy changes within the atom.

Describe the relationship between ground state, excited state, and photons.
Describe Bohr’s model and compare it to quantum theory.
All moving particles, ie. electrons and light have wave-particle duality. Explain.

Activity OR Emission Spectra LAB handout
Define ground state
Define excited state
Draw and Analyze the Bohr model of the atom.
Draw and Describe the main parts of the electromagnetic spectrum(colored pencils) and give an example of each.
Describe the electron location of an electron in terms of probability.

Which configurations represent excited atoms

1s12s1

1s22s22p63s23p64s23d10

Practice Questions:

1) Which electron configuration represents an atom in the excited state?
A) 2-7        B) 2-8-2        C) 2-7-1        D) 2-8-1
2) What is the total number of electrons in the second principal energy level of a calcium
atom in the ground state?
A) 8        B) 6        C) 2        D) 18
3) If the principal quantum number n = 4, what is the maximum number of electrons that
can be found?
A) 18        B) 8        C) 32        D) 6
4) The maximum number of electrons that a single orbital of 3d could hold is
A) 3        B) 2        C) 10        D) 6
5) What is the frequency of an X ray that has a wavelength of 1.15 x 10-1 m?
A) 5.32 x 109  sec-1    B) 2.61 x 109 sec-1    C) 3.44 x 109 sec-1
D) 7.8 x 107 sec-1
6) As an electron moves from its ground state to its excited state, the potential energy of
the electron :
A) remains the same    B) decreases    C) increases
7) The characteristic bright line spectrum is produced when the electron:
A) moves to a higher level            B) is lost by an atom
C) forms an ionic bond            D) moves to a lower level.
8) The greatest energy absorption occurs when the electron moves from:
A) 1s to 3s    B) 3p  to 3s    C) 4d to 4s    D) 4s to 3p
9) The relationship between the frequency and the wavelength is:
A) direct        B) linear        C) equal       D) inverse
10) What is the probability of finding an electron outside of its orbital?
A) zero        B) 5%        C) 10 %        D) 95%

Homework

p70-74 and 116-149, 154-207

## Naming Binary Compounds with Fixed Cations

Aim: How to name binary compounds with fixed cation.

Notes:

1. A binary compound is one made of two different elements. There can
be one of each element such as in NaCl or KF. There can also be several
of each element such as Na2O or AlBr3.

2. Please remember that all elements involved in this lesson have ONLY ONE charge. That includes BOTH the cation AND the anion involved in the formula.

3.  a)The order for names in a binary compound is first the cation, then the anion.
b)Use the name of cation with a fixed oxidation state directly from the periodic table.
c)The name of the anion will be made from the root of the element’s name plus the suffix “-ide.”

Example 1: Write the name of the following formula: H2S

Step #1 – Look at first element and name it. Result of this step = hydrogen.

Step #2 – Look at second element. Use root of its full name ( which
is sulf-) plus the ending “-ide.” Result of this step = sulfide.

These two steps give the full name of H2S. Notice that
the presence of the subscript is ignored. There are other types of
binary compounds where you must pay attention to the subscript. Those
compounds involve cations with variable charges.

Example 2: Write the name of the following formula: NaCl

Step #1 – Look at first element and name it. Result of this step = sodium.

Step #2 – Look at second element. Use root of its full name ( which
is chlor-) plus the ending “-ide.” Result of this step = chloride.

Example 3: Write the name of the following formula: MgBr2

Step #1 – Look at first element and name it. Result of this step = magnesium.

Step #2 – Look at second element. Use root of its full name ( which
is brom-) plus the ending “-ide.” Result of this step = bromide.

Note the presence of the subscript does not play a role in this name.

Example 4: Write the name of the following formula: KCl

The first part of the name comes from the first element symbol: potassium.
The second part of the name comes from the root of the second symbol plus ‘-ide,’ therefore chlor + ide = chloride.

This compound is named potassium chloride

Example 5: Write the name of the following formula: Na2S

First symbol is Na, so the first part of the name is sodium. (Note
the presence of the subscript does not play a role in this name.)
Second element is sulfur (from the symbol S), so the name is sulf + ide
= sulfide.

This compound is named sodium sulfide.

Here are examples of common roots:

 Cl: chlor- F: fluor- Br: brom- O: ox- I: iod- N: nitr-

Activity:

Write the correct name for:

1) MgS

2) KBr

3) Ba3N2

4) Al2O3

5) NaI

6) SrF2

7) Li2S

8) RaCl2

9) CaO

10) AlP

Write the correct name for:

11) K2S

12) LiBr

13) Sr3P2

14) BaCl2

15) NaBr

16) MgF2

17) Na2O

18) SrS

19) BN

20) AlN

Write the correct name for:

21) Cs2O

22) RbI

23) MgO

24) CaBr2

25) LiI

26) BeBr2

27) K2O

28) SrI2

29) BF3

30) Al2S3

## Representative Elements in periodic trends

Aim: How to relate the representative elements to trends of the periodic table
Activity: Students will create a wall of Groups. Each group poster must contain the symbols, the dot diagrams, the trends going down( ionization energy, electronegativity, atomic radius). Each must also contain electronconfigurations.

Summary of Representative Elements & Periodic Trends

1.              The representative elements are in groups 1,2,3, 13, 14, 16, 17, 18.

2.            The representative elements follow a perfect trend where the group number corresponds to the number of valence electrons

3.             The Dot diagrams correspond to the group number.

4.           The atomic radius general decreases as you go across a period. Why? From left to right atoms are in the same principal energy level(similar distance from the nucleus) but   experience very strong increase in positive nuclear charge which pulls the atom inward.

5.             The atomic radius general increases down a group. Why? Because more principal energy levels   occur which shield more electrons from the positive pull of the nucleus.

6.            The Period number corresponds to the outermost (highest) principal energy level.

## How to Count Atoms

Aim: How to count atoms when given a formula.

Notes:

Step One: Determine How many elements are in the formula. (capital letter)

Step Two: Look to the right of each element and determine the subscript. No subscript means, 1 atom. A subscript of 2 means 2 atoms of the element to the left of the subscript.

Step three : If there is a parthesis, count the atoms of each element in the parenthesis then multiply by the subscript outside of the parenthesis.

Activity: Students will create a table, and organize the number of atoms of each element.

1) MgS

2) KBr

3) Ba3N2

4) Al2O3

5) NaI

6) SrF2

7) Li2S

8) Ra(OH)2

9) CaO

10) AlP

11) K2S

12) LiBr

13) Sr3P2

14) BaCl2

15) NaBr

16) MgF2

17) Na2O

18) SrS

19) BN

20) AlN

21) Cs2O

22) (NH4)3PO4

23) MgO

24) CaBr2

25) LiI

26) BeBr2

27) K2O

28) SrI2

29) BF3

30) CuS

31) PbBr4

33) Pb3(PO4)2

34) Fe(OH)3

35) FeI2

36) Sn3P4

37) Cu2S

38) SnCl2

39) HgO

40) Hg2F2

41) AlPO4

42) KNO2

43) NaHCO3

44) CaCO3

45) Mg(OH)2

46) Na2CrO4

47) Ba(CN)2

48) K2SO4

49) NaH2PO4

50) NH4NO3

## How to Classify Matter

Motivation Idea/ Do Now : (Several Objects)What is this? what is matter?  How are all these objects similar? The word “matter” describes everything that has physical
existence. it
occupies space and has mass.

Aim: How to classify matter

Activity: Students will create tables, and organize differences between elements compounds and mixtures. (5 facts about elements. 7 facts about compounds. Flow diagram, etc

Question #1: We can successively separate matter into
categories by asking a sequence of yes or no questions. All matter can
be separated into two categories by first asking the question “Is only
one chemical substance present in the sample being considered?”

 YES - Pure Substance NO - Mixture

We can represent this question graphically:

 Keep in mind that terms like “chemical substance” or “mixture” haven’t really been defined yet. Hopefully, their definitions will be clearer as we go on.

Question #2: All pure substances can be separated into two
categories by asking the question “Can the sample be further decomposed
by chemical means?” (Notice we are ignoring mixtures for the moment.)

 YES - Compound NO - Element

We can represent this question graphically:

Notes: zuhmdal

1. Elements Only

An element= pure substance which cannot be broken down by
further chemical techniques. These include heating, cooling,
electrolysis and reacting with other chemicals. (By the way, it is
correct that an atom can be destroyed, but NOT by chemical means. You
must use a more powerful reaction, called a nuclear reaction, to
destroy or change atoms. That is a topic for another lesson.)

A sample of an element contains only one kind of atom in the
sample. Suppose you had a lump of copper in your hand. The ONLY type of
atom in the lump is copper. In the lump there are trillions and
trillions and trillions of copper atoms. NOTHING else.

If you were to heat the lump of copper, it would melt and
eventually vaporize. The smallest unit of the copper, called the atom,
would remain unaffected by this. The atoms of copper would be in the
solid state, the liquid state or the gaseous state, but they would be
EXACTLY the same in each state.

The atom is the smallest subdivision of an element which still
retains the properties of that element. In fact, a very good definition
of an atom is:

the smallest part of an element that can enter into a chemical combination

There are around 118 elements known to man, of which 20-30 are
really, really important. Almost every element that exists has some
form of use. There are some which are so unstable they only last for
seconds or even tiny fractions of a second and no use has yet been
found for them. However, ya never know!

Elements have names and symbols. For example hydrogen has the
symbol H and iron has the symbol Fe. Please note that Fe is one symbol,
not two. Also, make sure to use lower case for the second letter.
Writing BR for bromine is incorrect, writing it as Br is correct.

2. Just Compounds

A compound is a pure substance composed of two or more different
atoms chemically bonded to one another. A compound can be destroyed by
chemical means. It might be broken down into simpler compounds, into
its elements or a combination of the two. The key distinction is that
compounds break down whereas the SAME techniques do not cause an
element to break down.

The molecule is smallest subdivision of a compound that still
retains the properties of that compound. The parallel definition (to
the element one above) for the molecule is:

the smallest part of a compound that can enter into a chemical combination

Another definition, equally good, is that a molecule is the smallest stable part of a compound.

Water is a typical example of a compound. One molecule of water
is composed of two hydrogen atoms and one oxygen atom, chemically
bonded together. It is identified with its formula: H2O.

If you were to heat water (let’s start with ice), it would eventually melt, then vaporize. Each water molecule (each H2O)
would act as an independent unit and zoom around in the gas sample. The
three atoms making the water molecule would stay attached to each
other. In addition, water would enter into a chemical reaction acting
like a water molecule, NOT little separate atoms of hydrogen and
oxygen.

The compound is going to
have distinctly different properties that its elements. Hydrogen has a
set of properties, as does oxygen. However, the set of properties that
water has in no way like the two elements. For example, at room
temperature (about 20-25 °C) water is a liquid while hydrogen and
oxygen are gases.

Compounds have names and formulas. The formula is made from the
symbols of the elements in the molecule and how many of each element
there are. For example, glucose’s formula is C6H12O6.

3 Compounds and Molecules

The difference between compound and molecule causes distress among students. I will try and explain it more.

The word compound is meant when you are making general reference
to a chemical substance, as in “Go get a bottle of glucose from the
storeroom.” or “Glucose is the one of the end products of
photosynthesis.”

Inside the bottle is 500 grams of the compound glucose. Making
up the 500 grams of glucose are trillions and trillions of individual
glucose molecules, the formula of which is C6H12O6.

The plant makes the chemical compound called glucose. A sample
of the chemical compound glucose is made up of many glucose molecules,
all having the formula C6H12O6.

4. Elements and Molecules

At room temperature and room pressure, almost all elements are
considered to exist as single atoms. However, there are some which
exist as molecules.(1) There are seven diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2
(2) P4 and S8 also exist.At elevated temperatures, these molecules will break down
into single atoms, but in the above state listed, they are considered
to be elements.

The atoms are chemically bonded to each other, so they
are considered to be molecules, but they are not considered to be
compounds. They are elements.
If I am discussing the element oxygen, I mean O2. If I am discussing the element sulfur, I mean S8.
If I wished to discuss oxygen atoms or sulfur atoms, I would have to
say it explicitly or the context would clearly have to demand that
individual atoms of the element are being discussed.

Homework

Science Writing journal : Pick and element compound and a mixture and discuss how each relates to the definitions discussed in class.

## Electron configurations

Do Now: How many electrons are in each of the following neutral atoms? Na, C, N, S, P, K, Ca, H, He, Ne, Cl, O

Aim: How are electrons organized in the atom?

(demo:Use models. Tape 1s orbital on board , add 2 electrons,  Tape the 2s, 2p’s etc.What does it mean?)

Notes: Electron cloud model, quantum theory

1s, 2s 2p, 3s 3p 3d, 4s 4p 4d 4f. Each principle energy level has a corresponding number of sub levels.

Aufbau principle: building up
Students must remember the correct order. Use triangle mneumonic device
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f 5g
6s 6p 6d 6f 6g 6h
7s 7p 7d etc

ground state, excited state, Heisenberg Uncertainty principle

Activity: Have students complete an organization chart. Columns: Princ Energy level, sub levels, orbitals, # electrons

Activity: Have students use the orbital models to build their own atoms.
Activity: Write electron configurations for the first 35 atoms

Homework:
p99 #12,13
p113#60
p139 #18-20

## History of Atomic Theory

Do Now:
Define isotope
C-12   C-14    Na-23    H-3   U-236  U-235   H-2   C-13   Si-29   Na-22  K-40  Cl-35  K-39   Li-7  Fe-56

Write another symbol for each of these? How many neutrons in each of these.

Aim: To Explore the history of Atomic Structure

Materials needed:  Chalk/white board space, video, poster paper, models of s, p, and d orbitals (if models are not available bring in balloons or not.) Index cards with isotopes

Questions. What was the Plum pudding model( thomson’s model), Who was Dalton and what were his contributions? What did Rutherford’s experiment prove? What is the Bohr Model and draw an example using potassium. What is the quantum mechanical theory? What is the heisenberg uncertainty principle?

Notes: The modern model of the atom has evolved over a long period of time through the work of many scientists. Each atom has a nucleus, with an overall positive charge, surrounded by negatively charged electrons.
Subatomic particles contained in the nucleus include protons and neutrons.

The proton is positively charged, and the neutron has no charge.
The electron is negatively charged.
Protons and electrons have equal but opposite charges.
The number of protons is equal to the number of electrons in an atom. (Optional for review: Teacher will pass out handout for practice with protons, electrons and neutrons.)
The mass of each proton and each neutron is approximately equal to one atomic mass unit. An electron is much less massive than a proton or neutron.

In  the wave-mechanical model(electron cloud), the electrons are in orbitals, which are defined as regions of most probable electron location(ground state). Each electron in an atom has its own distinct amount of energy. When an electron in an atom gains a specific amount of energy, the electron is at a higher energy state (excited state.) When an electron returns from a higher energy state to a lower energy state, a specific amount of energy is emitted. The emitted energy can be used to identify an element.

Acitivity:

Students will break out into groups. Each will create a poster project and present to class. The topics: 1. rutherford experiment , 2. daltons atomic theory,  3. bohr model  4. electron cloud model (quantum theory)  5. Thomson model, plum pudding model p87-90

Each poster must include pictures, or drawings and pertinent notes about the topic.

Summary

history of atomic theory presentations

Homework

p113 #59-62
p112 #29-33

## physical and chemical change

Do Now: convert 56 mL to L. What are SI Units? What are SI prefixes?
Convert 6000 to scientific notation. Write the formula of percent
error.

materials needed: Huge Periodic table chart in the classroom. Index cards. Data projector or overhead projector. Exothermic and Endothermic salt. 2 beakers. wood block. water. paper towels. (possible chemicals) Zn, HCl, Cu, water, vinegar, baking soda, calcium chloride, sodium hydroxide, copper sulfate

Activity:
1.Devise an experiment that uses physical properties to identify
substances. List all materials needed.  Use EOC reference tables. Pick
various physical properties from the table and list them on an index
card. List density, solubility, boiling points, and melting points.
Then give the list to another classmate. The classmate must figure out
what substance it is.
Activity 2. Science Writing Journal:
Write a paragraph that discusses indicators of a chemical change.
List specific examples. Devise an experiment to show 3 chemical changes. (see chemicals)
Activity 3. List properties of metals.p155. List properties of nonmetals. p158

Notes:
A physical change results in the rearrangement of existing particles in
a substance. A chemical change results in the formation of different
substances with changed properties.

Chemical and physical changes can be exothermic or endothermic,(enrichment: potential diagrams)
(DEMO! add calcium chloride and salt b thermo. Show students immediate freeze of water. Also show immediate heat.)

The
structure and arrangement of particles and their interactions determine
the physical state of a substance at a given temperature and pressure.
(Have students draw a simple particle model to differentiate properties
of solids, liquids, and gases.
Physical properties of
substances can be explained in terms of chemical bonds and
intermolecular forces. These properties include conductivity,
malleability, solubility, hardness, melting point, and boiling point.

Summary: Have students come up to the board and make 2 columns, physical change, chemical change. They will list  examples of each.

homework:

p13 #11
p22 #31,32
p29 #1
p35 #25
p65 #10, 11

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