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