CHEMICAL EQUATIONS AND REACTIONS
reactant- A chemical substance that is present at the start of a chemical reaction
Decomposition- A reaction where a single compound breaks down into simpler compounds
Products- A substance that is formed as the result of a chemical reaction
Coefficient- A number or symbol multiplied with a variable or an unknown quantity in an algebraic term
Synthesis- Formation of a compound from simpler compounds or elements.
Single Replacement- A chemical reaction in which an element replaces one element in a compound
Doule Replacement- A chemical reaction between compounds in which the elements in the reactants recombine to form two different compounds, each of the products having one element from each of the reactants.
Law of Conservation of matter- The total amount of matter and energy available in the universe is fixed
Exothermic- Reacting that gives off heat to the environment.
Endothermic- Reaction that absorbs heat from its surroundings as the reaction proceeds.
Stoichiometry-The study of the relationships or ratios between two or more substances undergoing a physical or chemical change



http://www.youtube.com/watch?v=RnGu3xO2h74






















Chemical Reaction: A change in which a substance is changed into one or more new substances.

chemical%20reaction.gif

Endothermic and Exothermic

Exothermic- A chemical reaction in which more energy is released than required to break bonds in the initial reaction.

-an example of exothermic is table salt.

Na(s) + 0.5Cl2(s) = NaCl(s)

Endothermic- A chemical reaction in which a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form in the product molecules.

-an example of endothermic is photosynthesis.
sunlight + 6CO2(g) + H2O(l) = C6H12O6(aq) + 6O2(g)





LAW OF CONSERVATION OF MATTER / LAW OF CONSERVATION OF MASS

states that mass is neither created nor destroyed during a chemical reaction but is conserved.

404239464-B41.jpg

Balance of a chemical equation given the

reactants and products


2H2 + O2 =2 H2O


arrow
arrow
http://www.youtube.com/watch?v=4B8PFqbMNIw

2H2O

hydrogen-small
hydrogen-small
hydrogen-small
hydrogen-small

hydrogen-small
hydrogen-small
hydrogen-small
hydrogen-small
+

oxygen-small
oxygen-small
oxygen-small
oxygen-small
arrow
arrow
water molecule-small
water molecule-small
water molecule-small
water molecule-small
2 * 2.02g

+

32.00g

=

2 * 18.02g

Chemical

2H2
+
O2
arrow
arrow

2H2O

hydrogen-small
hydrogen-small
hydrogen-small
hydrogen-small

hydrogen-small
hydrogen-small
hydrogen-small
hydrogen-small
+

oxygen-small
oxygen-small
oxygen-small
oxygen-small
arrow
arrow
water molecule-small
water molecule-small
water molecule-small
water molecule-small
2 * 2.02g

+

32.00g

=

2 * 18.02g


The total mass of the reactants, 36.04g, is exactly equal to the total mass of the products, 36.04g.





external image images?q=tbn:ANd9GcT1osYtiX2zDrss0K-nMX-dkzQClflluYosk_wR1_hsU1ndIMA&t=1&h=178&w=209&usg=__KKB8B1frLYnDOUstj3rEmsDOq9k=
chemwiki.ucdavis.edu



The Four Types of Chemical Reaction
The Four Types of Chemical Reaction

CLASSIFICATION OF CHEMICAL REACTIONS

  1. Decomposition reactions

  2. Combination reactions (Synthesis reactions)

  3. Single-replacement reactions (Displacement reactions)

  4. Double-replacement reactions (Metathesis reactions)

  5. Combustion reactions

  6. Oxidation-reduction reactions (Redox reactions)


· Synthesis: A synthesis reaction is when two or more simple compounds combine to form a more complicated one. These reactions come in the general form of:
A + B ---> AB
One example of a synthesis reaction is the combination of iron and sulfur to form iron (II) sulfide:
8 Fe + S8 ---> 8 FeS








































· Decomposition: A decomposition reaction is the opposite of a synthesis reaction - a complex molecule breaks down to make simpler ones. These reactions come in the general form:
AB ---> A + B
One example of a decomposition reaction is the electrolysis of water to make oxygen and hydrogen gas:
2 H2O ---> 2 H2 + O2








































· Single displacement: This is when one element trades places with another element in a compound. These reactions come in the general form of:
A + BC ---> AC + B
One example of a single displacement reaction is when magnesium replaces hydrogen in water to make magnesium hydroxide and hydrogen gas:
Mg + 2 H2O ---> Mg(OH)2 + H2








































· Double displacement: This is when the anions and cations of two different molecules switch places, forming two entirely different compounds. These reactions are in the general form:
AB + CD ---> AD + CB
One example of a double displacement reaction is the reaction of lead (II) nitrate with potassium iodide to form lead (II) iodide and potassium nitrate:
Pb(NO3)2 + 2 KI ---> PbI2 + 2 KNO3





5 basic chemical reactions
1. Synthesis reactions: a reaction where two or more substances combine to form a new compound.
For example:
A + X à AX
Where A and X can be elements or compounds and AX is a compound.

2. Decomposition Reaction: the reverse of a synthesis reaction and is a reaction in which a single compound undergoes a reaction that produces two or more simpler substances.
For example:
AX à A + X
Where AX is the compound and A and X can be elements or compounds.
Single replacement reactions: also known as displacement reaction; is a reaction in which one element in a compound.
For example:
A + BX à AX + B
Or
Y + BX à BY + X

Where A, B, X, and Y are elements and AX, BX, and BY are compounds.

4. Double-replacement reactions: a reaction in which the ions of two compounds exchange places in an aqueous solution to form two new compounds.

For example:
AX + BY à AY + BX
Where A, X, B, and Y in the reactants represent ions and products AY and BX represent ionic or molecular compounds.

5. Combustion Reactions: a reaction in which a substance combines with oxygen and releases a large amount of energy in the form of heat and light.
There is no example for combustion reactions.
http://www.ric.edu/faculty/ptiskus/chem_review/Index.htm





EQUATIONS SHOWS;
· the reactants which enter into a reaction.
· the products which are formed by the reaction.

· the amounts of each substance used and each substance produced.
IMPORTANT PRINCIPLES TO REMEMBER;
· Every chemical compound has a formula which cannot be altered.
A chemical reaction must account for every atom that is used. This is an application of the Law of Conservation of Matter which states that in a chemical reaction atoms are neither created nor destroyed.
Four basic types of chemical reactions:
Synthesis (composition);
· two or more elements or compounds may combine to form a more complex compound.
· Basic form: A + X AX

Examples of synthesis reactions:
1. Metal + oxygen metal oxide
2. EX. 2Mg(s) + O2(g)
2MgO(s
3. Nonmetal + oxygen
nonmetallic oxide
4. EX. C(s) + O2(g)
CO2(g)
5. Metal oxide + water
metallic hydroxide
6. EX. MgO(s) + H2O(l)
Mg(OH)2(s)
7. Nonmetallic oxide + water
acid
8. EX. CO2(g) + H2O(l)
; H2CO3(aq)
9. Metal + nonmetal
salt
10. EX. 2 Na(s) + Cl2(g)
2NaCl(s)

11. A few nonmetals combine with each other.
12. EX. 2P(s) + 3Cl2(g) 2PCl3(g)
These two reactions must be remembered:
1. N2(g) + 3H2(g) 2NH3(g

STOICHIOMETRY

COMBINATIONS OF ELEMENTS AND THEIR REACTIONS: Study chemical reactions by reading
This is an example of a chemical reaction and a reversal.
This is an example of a chemical reaction and a reversal.
sections on stoichiometry in chemistry text books and demonstrating them with laboratory experiments. (See laboratory cautions below). Combining elements to form new combinations is a very important part of chemistry. (If you don't know the meaning of the words here look them up in the "TERMS" links provided below. Also look up the words in your dictionary for additional understanding.) Learn about the concept of a MOLE.


external image moz-screenshot-6.png-study of quantitive relationship between amounts of reactants used and products formed by chemical reaction
- also based on law of conservation of mass.

As tiny particles of iron react with oxygen in the air, iron (III) oxide (Fe2O3) is produced
4Fe(s)+ 3O2(g) --> 2Fe2O3(S)



4mol Fe x 55.85g Fe= 2234g Fe
1molFe

3molO2 x 32.00g O2 = 96.00g O2
1mol O2

STEPS ON SOLVING STOICHIOMETRY PROBLEMS

http://www.youtube.com/watch?v=qEibF3UOO1U
BASIC STOICHIOMETRY
Pronounce stoichiometry as “stoy-kee-ah-met-tree,” if you want to sound like you know what you are talking about, or “stoyk:,” if you want to sound like a real geek. Stoichiometry is just a five dollar idea dressed up in a fifty dollar name. You can compare the amounts of any materials in the same chemical equation using the formula weights and the coefficients of the materials in the equation. Let’s consider the equation for the Haber reaction, the combination of nitrogen gas and hydrogen gas to make ammonia.
N2 + 3 H2 external image arrow.gif 2 NH3
The formula for nitrogen is N2 and the formula for hydrogen is H2. They are both diatomic gases. The formula for ammonia is NH3. The balanced equation requires one nitrogen molecule and three hydrogen molecules to make two ammonia molecules, meaning that one nitrogen molecule reacts with three hydrogen molecules to make two ammonia molecules or one MOL of nitrogen and three MOLS of hydrogen make two MOLS of ammonia. Now we are getting somewhere. The real way we measure amounts is by weight (actually, mass), so 28 grams (14 g/mol times two atoms of nitrogen per molecule) of nitrogen and 6 grams of hydrogen (1 g/mol times two atoms of hydrogen per molecule times three mols) make 34 grams of ammonia. Notice that no mass is lost or gained, since the formula weight for ammonia is 17 (one nitrogen at 14 and three hydrogens at one g/mol) and there are two mols of ammonia made. Once you have the mass proportions, any mass-mass stoichiometry can be done by good old proportionation. What is the likelihood you will get just a simple mass-mass stoich problem on your test? You should live so long. Well, you should get ONE.
Rather than thinking in terms of proportions, think in mols and mol ratios, a much more general and therefore more useful type of thinking. A mol ratio is just the ratio of one material in a chemical equation to another material in the same equation. The mol ratio uses the coefficients of the materials as they appear in the balanced chemical equation. What is the mol ratio of hydrogen to ammonia in the Haber equation? 3 mols of hydrogen to 2 mols of ammonia. Easy. In the standard stoichiometry calculations you should know, ALL ROADS LEAD TO MOLS. You can change any amount of any measurement of any material in the same equation with any other material in any measurement in the same equation. That is powerful. The setup is similar to Dimensional Analysis.
1. Start with what you know (GIVEN), expressing it as a fraction.
2. Use definitions or other information to change what you know to mols of that material.
3. Use the mol ratio to exchange mols of the material given to the mols of material you want to find.
4. Change the mols of material you are finding to whatever other measurement you need.
How many grams of ammonia can you make with 25 grams of hydrogen? (Practice your mol math rather than doing this by proportion. Check it by proportion in problems that permit it.)
You are given the mass of 25 grams of hydrogen. Start there.
Visual for solvingstoichiometry problems in a Dimensional Analysis style.
Visual for solvingstoichiometry problems in a Dimensional Analysis style.
25 g H2/1 Change to mols of hydrogen by the formula weight of hydrogen 1 mol of H2 = 2.0 g. (The 2.0 g goes in the denominator to cancel with the gram units in the material given.) Change mols of hydrogen to mols of ammonia by the mol ratio. 3 mols of hydrogen = 2 mols of ammonia. (The mols of hydrogen go in the denominator to cancel with the mols of hydrogen. You are now in units of mols of ammonia.) Convert the mols of ammonia to grams of ammonia by the formula weight of ammonia, 1 mol of ammonia = 17 g. (Now the mols go in the denominator to cancel with the mols of ammonia.) Cancel the units as you go.
The math on the calculator should be the last thing you do. 2 5 ÷ 2 . 0 x 2 ÷ 3 x 1 7 = and the number you get (141.66667) will be a number of grams of ammonia as the units in your calculations show. Round it to the number of significant digits your instructor requires (often three sig. figs.) and put into scientific notation if required. Most professors suggest that scientific notation be used if the answer is over one thousand or less than a thousandth. The answer is 142 grams of ammonia.
The calculator technique in the preceding paragraph illustrates a straightforward way to do the math. If you include all the numbers in order as they appear, you will have less chance of making an error. Many times students have been observed gathering all the numbers in the numerator, gathering all the numbers in the denominator, presenting a new fraction of the collected numbers, and then doing the division to find an answer. While this method is not wrong, the extra handling of the numbers has seen to produce many more errors.
See the Stoichiometry Roadmap for a way to consider this idea graphically. This example starts at "mass given" and goes throught the mol ratio to "mass find."
mass1--Fw1--mol1--mol ratio2/1--mol2--Fw2--mass2
mass1--Fw1--mol1--mol ratio2/1--mol2--Fw2--mass2
Notice by the chart above we may get the number of mols of material given if we change the mass by the formula weight, but in our continuous running math problem, we don’t have to stop and calculate a number of mols. Students who insist on doing so tend to get more calculator errors.
The more traditional formula for converting mols to mass would be, where Fw is the formula weight, m is the mass, and n is the number of mols: n x Fw = m. You should be able to "see" these formula relationships on the roadmap.

EXAMPLE:
The question: If you react 18.7 grams of Iron(III) Oxide with Carbon(II) Oxide (1) how many
grams of Carbon (II) Oxide would you need? (2) How many grams of Iron would you produce?
And (3) How many ml of Carbon Oxide would you produce?
Molecular weights:
Fe2O3 Fe = 2 X 56 = 112 CO C = 1 X 12 = 12 CO2 C 1 X 12 = 12
O = 3 X 16 = 48 O = 1 X 16 = 16 O 2 X 16 = 32
-------- -------- -------
160 28 44
1. Write and balance the equations.
Iron (III) Oxide + Carbon (II) Oxide Iron + Carbon Oxide
Write the equation
Fe2O3 + CO 2 Fe + CO2
Balance the Equation
1 Fe2O3 + 3 CO 2 Fe + 3 CO2
2. Write underneath the equation what is GIVEN and what is ASKED FOR.
1 Fe2O3 + 3 CO 2 Fe + 3 CO2
18.7 g _?_ g
3. Find moles of what is given
Moles of Fe2O3 = Mass Fe2O3 / Molecular weight Fe2O3
= 18.7 / 160
= 0.116875
4. Write the mole to mole relationship between what you looking for and what you have.
(MULTIPLY BY WHERE YOU ARE GOING AND DIVIDE BY WHERE YOU
CAME FROM. 3 moles CO and 1 Mole Fe2O3
Mole CO = 3/1 Moles Fe2O3
= 3 X (0.116875)
= 0.350625
5. Change moles to units you need, in this case GRAMS of CO was asked for.
Mass CO = (Mole CO) X (Molecular weight CO)
= (0.350625) X (28)
= 9.82 grams of CO needed (Final is THREE SIGNIFICANT FIGURES)
1 Fe2O3 + 3 CO 2 Fe + 3 CO2
18.7 g _?_ g
Now finish the problems and find mass Fe using the steps above.
1. Find moles of what is given
Moles of Fe2O3 = Mass Fe2O3 / Molecular weight Fe2O3
= 18.7 / 160
= 0.116875
2. Find the mole to mole relationship.
Moles Fe = 2/1 Moles Fe2O3
= 2 X (0.116875)
= 0.23375
3. Find what is being asked for.
Mass Fe = (Moles Fe) X (Molecular Weight Fe)
= (0.23375) X (56)
= 13.1 grams of Fe Produced (Final is THREE SIGNIFICANT FIGURES)
1 Fe2O3 + 3 CO 2 Fe + 3 CO2
18.7 g _?_ ml
Now finish the problems and find volume CO2 using the steps above.
1. Find moles of what is given
Moles of Fe2O3 = Mass Fe2O3 / Molecular weight Fe2O3
= 18.7 / 160
= 0.116875
2. Find the mole to mole relationship.
Moles CO2 = 3/1 Moles Fe2O3
= 3 X (0.116875)
= 0.350625
3. Find what is being asked for.
Volume CO2 = (Moles CO2) X (22.4)
= (0.350625) X (22.4)
= 7.854 liters or 7854 ml of CO2 produced (Final is AT LEAST THREE
SIGNIFICANT FIGURES
So we now have the following information:
Fe2O3 + 3 CO 2 Fe + 3 CO2
18.7 g 9.82 g 13.1 g 7854 ml
http:/<object width="480" height="385"><param name="movie" value="http://www.youtube.com/v/b4F9zjWgrDE?fs=1&amp;hl=en_US"></param><param name="allowFullScreen" value="true"></param><param name="allowscriptaccess" value="always"></param><embed src="http://www.youtube.com/v/b4F9zjWgrDE?fs=1&amp;hl=en_US" type="application/x-shockwave-flash" allowscriptaccess="always" allowfullscreen="true" width="480" height="385"></embed></object>www.youtube.com/watch?v=b4F9zjWgrDE mass= http://www.youtube.com/watch?v=b4F9zjWgrDE
http://www.youtube.com/watch?v=b4F9zjWgrDE
Similarly mas of product= 2mol Fe 2O3 x 159.7 Fe2O3 = 319.4g
1mol Fe2O3

FOR RULES, CHECK OUT THIS VIDEO:
http://www.youtube.com/watch?v=rWCv-ymhpfY


The rules followed in the determination of stoichiometric relationships are based on
the laws of conservation of mass and energy and the law of combining weights or volumes

STOICHIOMETRY GRAMS TO GRAMS VIDEO:


http://teachertube.com/viewVideo.php?title=Stoichiometry_grams_to_grams&video_id=185136

Sample Problems

Problem: Given the following equation at STP:

|| N2(g) + H2(g)→NH3(g) || ||

Determine what volume of H2(g) is needed to produce 224 L of NH3(g).
Solution:
br> Step 1: Balance the equation.

|| N2(g) + 3H2(g)→2NH3(g) || ||

Step 2: Convert the given quantity to moles. Note in this step, 22.4 L is on the denominator of the conversion factors since we want to convert from liters to moles. Remember your conversion factors must always be arranged so that the units cancel.

|| external image latex_img28.gif = 10 moles of NH3(g) || ||

Step 3: mole ratio.

|| external image latex_img26.gif = 15 moles H2(g) || ||

Step 4, convert to desired units:

|| external image latex_img29.gif = 336 L H2(g) || ||

Now for a more challenging problem:

Given the following reaction:

|| 2H2S(g) + O2(g)→SO2(g) + 2H2O(s) || ||

How many atoms of oxygen do I need in order to get 18 g of ice?
Solution

Step 1. The equation is partially balanced already, but let's finish the job.

|| 2H2S(g) +3O2(g)→2SO2(g) + 2H2O(s) || ||

Step 2, convert to moles:

1 formula unit of H2O has 2 atoms of H and 1 atom of O
The atomic mass of H is 1 gram/mole
Atomic mass of O = 16 grams/mole


|| GFM of H2O(s) = external image latex_img11.gif + external image latex_img12.gif = 18 grams / mole || ||


external image latex_img7.gif×1 mole = 1 mole of H2O(s)



Step 3, mole ratio:


external image latex_img8.gif×3 moles O2(g) = 1.5 moles O2(g)



Step 4, convert to desired units:


Stochiometry sample quiz:

http://chemistry.about.com/library/weekly/blstoichiometryquiz.htm

Reaction stoichiometry

Reaction stoichiometry allows us to determine the amount of substance that is
consumed or produced by a reaction. The following video considers the first part
of this: how much of a reactant is consumed in a chemical reaction.
Product formation is discussed elsewhere.

CHECK OUT THIS VIDEO!

http://www.chemcollective.org/stoich/reaction_stoi.php



external image latex_img22.gif = 9.03×1023 molecules O2(g)



Is this the answer? No. The question asks for ATOMS of oxygen. There are two atoms of oxygen in each molecule of O2(g).
external image latex_img27.gif×2 atoms O = 1.806×1024 atoms O



INTERPRET A BALANCED EQUATIONIN TERMS OF MOLES AND REPRESENTATIVE PARTICLES.

http://docs.google.com/viewer?a=v&q=cache:WFlWoud8DiUJ:howethnotes.wikispaces.com/file/view/Chemistry%2BCh%2B12%2Bnotes.ppt+interpret+a+balanced+equation+in+terms+of+molecules+and+representative+particles&hl=en&gl=us&pid=bl&srcid=ADGEESgr4f-JLQhSdm23-Xmfqnd-FzFEovGi_b8RLZwLzMqddPvprfPYoqRI90qbKeBicFEG32b1V9X-cni37So65xck-0QDkwhKdtJuOmYVl2bCzOILvbnbUH4AeM6r_kv_FCjBu2pV&sig=AHIEtbTJhbLiLtzeXr4rBHyRde17ktci5w

Perform stoichiometric calculations for a given chemical reaction including mass-mass calculations.



FourTypesofChemicalReactions_ppt_05.jpg



how to balance chemical equations


=


=
STOICHIOMETRY RAP SONG!!
http://www.teachertube.com/viewVideo.php?video_id=168372&title=Awesome_Stoichiometry_Rap_percent_yield


EXAMPLE OF CHEMICAL EQUATION




What is a chemical equation?

When a chemical reaction occurs, it can be described by an equation. This shows the chemicals that react (called the reactants) on the left-hand side, and the chemicals that they produce (called the products) on the right-hand side. The chemicals can be represented by their names or by their chemical symbols.
Unlike mathematical equations, the two sides are separated by an arrow, that indicates that the reactants form the products and not the other way round.

A large number of chemical equations are more complicated than the simple ones you will see in this section. They are reversible, which means that the reactants react together to form the products, but as soon as the products are formed, they start to react together to reform the reactants!
Reversible equations proceed in both directions at once, with reactants forming products and products forming reactants simultaneously. Eventually, the system settles down and a balance (an equilibrium) is reached, with the reactants and products present in stable concentrations. This does not mean that the reaction stops, merely that it proceeds in both directions at the same rate, so that the concentrations do not change.
Reversible reactions are indicated with a double arrow as shown in the example below:
Ethanoic acid + ethanol ethyl ethanoate + water

external image prof_l.gif
In this case, ethanol (which is alcohol, basically) reacts with ethanoic acid (the main constituent of vinegar) to form ethyl ethanoate and water. However, the ethyl ethanoate produced reacts with the water produced to recreate the ethanol and ethanoic acid again. In practice, the chemicals reach a balance point, called equilibrium where all four chemicals are present.


http://richardbowles.tripod.com/chemistry/balance.htm
external image snick1.gif












www.youtube.com


chemical_reactions.jpg





CHEMICAL REACTIONS:

In the late 1890s, a chemist Sir William Ramsay, discovered the
**elements**
helium, neon, argon, krypton, and xenon.
These elements, along with radon, were placed in group VIIIA of the periodic table and nicknamed inert (or noble) gases because of their tendency not to react with other elements. The tendency of the noble gases to not react with other elements has to do with their electron configurations. All of the noble gases have full
**valence shells**; this configuration is a stable configuration and one that other elements try to achieve by reacting together. In other words, the reason **atoms**
react with each other is to reach a state in which their valence shell is filled.

http://www.visionlearning.com/library/module_viewer.php?mid=54&l=

Law of conservation of Matter:


external image image013.jpg
www.http://martine.people.cofc.edu/111LectWeek1_files/image013.jpg





I. Formulas show chemistry at a standstill. Equations show chemistry in action.
    1. the reactants which enter into a reaction.
    2. the products which are formed by the reaction.
    3. the amounts of each substance used and each substance produced.
  • B. Two important principles to remember:

    1. Every chemical compound has a formula which cannot be altered.
    2. A chemical reaction must account for every atom that is used. This is an application of the Law of Conservation of Matter which states that in a chemical reaction atoms are neither created nor destroyed.
  • C. Some things to remember about writing equations:
    1. The diatomic elements when they stand alone are always written H2, N2, O2, F2, Cl2, Br2, I2
    2. The sign, → , means "yields" and shows the direction of the action.
    3. A small delta, (D), above the arrow shows that heat has been added.
    4. A double arrow, ↔ , shows that the reaction is reversible and can go in both directions.
    5. Before beginning to balance an equation, check each formula to see that it is correct. NEVER change a formula during the balancing of an equation.
    6. Balancing is done by placing coefficients in front of the formulas to insure the same number of atoms of each element on both sides of the arrow.**Practice Balancing Equations**
    7. Always consult the **Activity Series** of metals and nonmetals before attempting to write equations for replacement reactions.
    8. If a reactant or product is a solid, (s) is placed after the formula.
    9. If a reactant or product is a gas, (g) is placed after it.
    10. If a reactant or product is in water solution, (aq) is placed after it.
    11. Some products are unstable and break down (decompose) as they are produced during the reaction. You need to be able to recognize these products when they occur and write the decomposition products in their places.
  • Examples:

    • H2CO3(aq) → H2O(l) + CO2(g)Carbonic acid, as in soft drinks, decomposes when it is formed.

    • H2SO3(aq) → H2O(l) + SO2(g)Sulfurous acid also decomposes as it is formed.

NH4OH(aq) → NH3(g) + H2O(l)You can definitely smell the odor of ammonia gas because
whenever "ammonium hydroxide" is formed it decomposes into ammonia and water
For more information :
http://www.files.chem.vt.edu/RVGS/ACT/notes/Types_of_Equations.html






  • Stoichiometry (is a branch of chemistry that deals with the quantitative relationships that exist between the
reactants and produ**toichiometr**cts in chemical reactions. In a balanced chemical reaction, the relations among quantities
of reactants and products typically form a ratio of whole numbers. For example, in a reaction that forms ammonia
(NH3), exactly one molecule of nitrogen (N2) reacts with three molecules of hydrogen (H2) to produce two molecules
of NH3:

N2 + 3H2 → 2NH3




  • MOLES AND REPRESENTATIVE PARTICLES





  • chem.jpg


  • Calculate the mass of a product

    • equation.gif
      Stoichiometry


























aa.jpg

SYNTHESIS REACTION
In a synthesis reaction two or more simple substances combine to form a more complex substance. Two or more reactants yielding one product is another way to identify a synthesis reaction.
For example, simple hydrogen gas combined with simple oxygen gas can produce a more complex substance-----water!
The chemical equation for this synthesis reaction looks like:
synthesisimage
synthesisimage

reactant + reactant -------> product
To visualize a synthesis reaction look at the following cartoon:
birdandwormimage
birdandwormimage


In the cartoon, the skinny bird (reactant) and the worm (reactant) combine to make oneproduct, a fat bird.


DECOMPOSITION REACTION
In a decomposition reaction a more complex substance breaks down into its more simple parts. One reactant yields 2 or more products. Basically, synthesis and decomposition reactions are opposites.
For example, water can be broken down into hydrogen gas and oxygen gas. The chemical equation for this decomposition reaction looks like:
decomposeimage
decomposeimage

reactant -------> product + product
To visualize a decomposition reaction look at the following cartoon:
eggandturtleimage
eggandturtleimage

In this cartoon the egg (the reactant), which contained the turtle at one time, now has opened and the turtle (product) and egg shell (product) are now two separate substances.
SINGLE REPLACEMENT REACTION
In a single replacement reaction a single uncombined element replaces another in a compound. Two reactants yield two products. For example when zinc combines with hydrochloric acid, the zinc replaces hydrogen. The chemical equation for this single replacement reaction looks like:
singlereplaceimage
singlereplaceimage

reactant + reactant ---------> product + product
To visualize a single replacement reaction look at the following cartoon:
dancerimage
dancerimage



Notice, the guy in the orange shirt steals the date of the other guy. So, a part of one of the reactants trades places and is in a different place among the products.
DOUBLE REPLACEMENT REACTION
In a double replacement reaction parts of two compounds switch places to form two new compounds. Two reactants yield two products. For example when silver nitrate combines with sodium chloride, two new compounds--silver chloride and sodium nitrate are formed because the sodium and silver switched places. The chemical equation for this double replacement reaction looks like:
doublereplaceimage
doublereplaceimage

reactant + reactant ---------> product + product
To visualize a double replacement reaction look at the following cartoon:
tradinghatsimage
tradinghatsimage

ENERGY OF CHEMICAL REACTIONS
Chemical reactions always involve a change in energy. Energy is neither created or destroyed. Energy is absorbed or released in chemical reactions. Chemical reactions can be described as endothermic or exothermic reactions.
Endothermic Reactions
Chemical reactions in which energy is absorbed are endothermic. Energy is required for the reaction to occur. The energy absorbed is often heat energy or electrical energy. Adding electrical energy to metal oxides can separate them into the pure metal and oxygen. Adding electrical energy to sodium chloride can cause the table salt to break into its original sodium and chlorine parts.
Exothermic Reactions
Chemical reactions in which energy is released are exothermic. The energy that is released was originally stored in the chemical bonds of the reactants. Often the heat given off causes the product(s) to feel hot. Any reaction that involves combustion (burning) is an exothermic chemical reaction




Stoichiometric Calculations


Stoichiometric Calculations


Now we want to ask (and be able to answer) some very important questions. For example, suppose that I am burning a specific amount of Mg in the presence of oxygen and I want to know how much product I will have. The chemical equation is
Mg(s) + O2(g)
arrow right
arrow right
MgO(s)
Before I can answer the question, I must balance the equation. Just as numbers without units are meaningless, an unbalanced equation is insufficient. So we balance the equation:
2Mg(s) + O2(g)
arrow right
arrow right
2MgO(s)
Now if I have 0.145 g of Mg, how much MgO will this produce?
The key to answering this question is to understand that reactions occur on the same mole basis as the balanced stoichiometric coefficients. This equation says that 2 moles of Mg produce w moles of MgO. So how many moles of Mg did I start with?
.145 g Mg/(24.31 g/mole) = .00596 moles of Mg
So how many moles of MgO will this reaction produce?
.00596 moles Mg * (2 moles MgO/2moles Mg) = .00596 moles of MgO
which is .00596 moles of MgO *(40.31 g MgO/mole MgO) = .240 gms of MgO
Similarly we can calculate the number of grams of oxygen needed to complete the reaction since 2 moles of Mg react with 1 mole of oxygen:
.00596 moles of Mg * (1 mole of O2/2 mole Mg) = .00298 moles O2
and the grams of oxygen would be
.00298 moles O2 * (32.00 gms O2/mole O2 = .0954 g O2
Note that we had to use 32.00 for the molar mass of oxygen as a diatomic



Stoichiometric Calculations

The ability to balance and interpret equations enable us to make calculations involving masses
of substances involved in the reaction. This involves understanding mass, mole, and volume.
Stoichiometry is the study of quantitative relationships involved in chemical reactions.
There are two categories of stoichiometric problems you will be expected to be able to solve.
1. One deals with moles only. For example, a problem might say: Given this reaction:
2 H2 + O2 ---> 2 H2O and excess hydrogen, 1.50 moles of oxygen gas will produce how many
moles of water? Answer = 3.00 moles
2. The second deals with grams. For example, a problem might say: Given the reaction:
2 H2S + 3 O2 ---> 2 H2O + 2 SO2; 24.0 grams of oxygen will react completely with how many
grams of hydrogen sulfide?
Before We Start
Look at the equation in No. 2 just above. It is important to realize that this equation can be
understood two different ways. First, it can be understood to mean that two molecules of
hydrogen sulfide react with three molecules of oxygen gas to produce two molecules of water
and two molecules of sulfur dioxide. Second, IT CAN ALSO be understood in terms of moles.
The same 2:3:2:2 sequence applies, only now it is moles and not molecules.



aaa.JPG
Image.jpg



http://www.youtube.com/watch?v=8KH3laR2iR4&feature=related




For the electrolysis of water, there are 2 hydrogen and 1 oxygen on the reactant side and 2 hydrogen and 2 oxygen on the product side of the equation.
H2O → H2 + O2

#H
2
2
#O
1
2
A 2 can be added in front of the water to balance the oxygen:
2H2O → H2 + O2

#H
4
2
#O
2
2
However, this change unbalances the hydrogen, so a 2 can be added before the hydrogen gas in the product:
2H2O → 2H2 + O2

#H
4
4
#O
2
2
Because the electrolysis of water contains oxygen gas and fractions can be used, the equation could also be balanced:
H2O → H2 + external image Stoichiometry_II_Chemical_Equations_02.gifO2

for more information:
http://www.education.com/reference/article/stoichiometry-ii-chemical-equations/



H + O
arrow
arrow
H2O

The plus sign on the left side of the equation means that hydrogen (H) andoxygen (O) are reacting. Unfortunately, there are two problems with this chemical equation. First, because atoms like to have full valence shells, single H or O atoms are rare. In nature, both hydrogen and oxygen are found as diatomic molecules, H2 and O2, respectively (in forming diatomic molecules the atoms share electrons and complete their valence shells). Hydrogen gas, therefore, consists of H2 molecules; oxygen gas consists of O2. Correcting our equation we get:H2 + O2
arrow
arrow
H2O

for more information:http://www.visionlearning.com/library/module_viewer.php?mid=56



=

=



Balancing Chemical Equations===


Example 1

C5H12 + O2 ---> CO2 + H2O
Answer »
There are five carbons on the left but only one on the right, and on each side the carbon is in a single chemical species. Put a 5 in front of the CO2 on the right hand side.

C5H12 + O2 ---> 5CO2 + H2O

There are twelve hydrogens on the left but only two on the right hand side, and hydrogen is in a single species on each side. Put a 6 in front of the H2O on the right hand side.

C5H12 + O2 ---> 5CO2 + 6H2O

Finally, there are only two oxygens on the left hand side but 16 of them on the right hand side. So put a 8 in front of the O2 on the left hand side.

C5H12 + 8O2 ---> 5CO2 + 6H2O
It's now a balanced chemical equation.




Example 2

Zn + HCl ---> ZnCl2 + H2
Answer »
There are two chlorines on the right but only one on the left, and the chlorine is in a single chemical species on each side . Put a 2 in front of the HCl on the left hand side.

Zn + 2HCl ---> ZnCl2 + H2

And if you look carefully, you will see that the equation is now balanced, with one Zn on each side, two hydrogens on each side and two chlorines on each side. Some examples can be rather easy!



Example 3

Ca(OH)2 + H3PO4 ---> Ca3(PO4)2 + H2O
Answer »
There are three calciums on the right but only one on the left, and the calcium is in a single chemical species on each side . Put a 3 in front of the Ca(OH)2 on the left hand side.

3Ca(OH)2 + H3PO4 ---> Ca3(PO4)2 + H2O

There are two PO4 ions on the right but only one on the left side, and the P doesn't appear anywhere else (so the group remains intact). Put a 2 in front of the H3PO4 on the left side.

3Ca(OH)2 + 2H3PO4 ---> Ca3(PO4)2 + H2O

Finally, there are six oxygens on the left hand side not present as PO4 but only one on the right hand side not in the PO4. So put a 6 in front of the H2O on the right hand side.

3Ca(OH)2 + 2H3PO4 ---> Ca3(PO4)2 + 6H2O
It's now a balanced equation. Note how we treated the PO4 ion as a single species to be balanced.



Example 4

FeCl3 + NH4OH ---> Fe(OH)3 + NH4Cl
Answer »
The most obvious error is that there are three chlorines on the left but only one on the right, and the chlorine is in a single chemical species on each side . Put a 3 in front of the NH4Cl on the right hand side.

FeCl3 + NH4OH ---> Fe(OH)3 + 3NH4Cl

The next most obvious unbalanced part is that there are now three NH4 groups on the right but only one on the left hand side. So put a 3 in front of the NH4OH on the left.

FeCl3 + 3NH4OH ---> Fe(OH)3 + 3NH4Cl

And if you count up the atoms on each side, you will see that this is now a balanced chemical equation.

http://www.sky-web.net/science/balancing_chemical_equations_examples.htm


Balancing chemical equations
http://www.youtube.com/watch?v=lwP_L1R79Xg
http://www.youtube.com/watch?v=VUnYowO5_78__http://www.youtube.com/watch?v=q_rbjDGiyWM&feature=fvst__


TYPES OF CHEMICAL REACTIONS EXAMPLES:



image005.jpg










































Classification of Chemical Reactions

There are many kinds of chemical reactions and also many ways of classifying them. We will look at a few the ways to classify chemical reactions. Five traditional types of chemical reactions are

  1. Decomposition reactions
  2. Combination reactions (Synthesis reactions)
  3. Single-replacement reactions (Displacement reactions)
  4. Double-replacement reactions (Metathesis reactions)
  5. Combustion reactions
  6. Oxidation-reduction reactions (Redox reactions)
Decomposition Reactions
A decomposition reaction is a reaction in which a single compound decomposes to two or more other substances. A general equation that describes a decomposition reaction is AB --> A + B where A and B can be elements or compounds. Most compounds can be broken down into simpler substances or decomposed. Often this can be done by heating the compound. For example the industrial preparation of lime (calcium oxide) involves the decomposition of calcium carbonate by heating it.
CaCO3(s) ---> CaO(s) + CO2(g)
Combination Reactions

A combination reaction is a reaction in which two substances combine to form a third. A general equation that describes a combination reaction is A + B --> AB. Again A and B can be either elements or compounds. Decomposition and combination reactions can be considered to be the reverse of each other. Under some conditions it is possible to change conditions and cause a decomposition reaction to become a combination reaction or vice versa. The reaction of calcium oxide with sulfur dioxide to form calcium sulfite is an example of a combination reaction.
CaO(s) + SO2(g) ---> CaSO3(s)
Single-Replacement Reactions

A single-replacement reaction is a reaction in which an element reacts with a compound and replaces another element in the compound. A general equation that describes a single-replacement reaction is A + BC --> AB + C, where A and C are elements and BC and AB are compounds. The reaction in which copper displaces silver from an aqueous solution of silver nitrate is an example of a single-replacement reaction.
Cu(s) + 2 AgNO3(aq) ---> Cu(NO3)2(aq) + 2 Ag(s)
Double-Replacement Reactions

A double-replacement reaction (also called a metathesis reaction) is a reaction in which there is an exchange of positive ions between two compounds. These reactions generally take place between two ionic compounds in aqueous solution. A general equation that describes a double-replacement reaction is AB + CD --> AD + CB, where A and C are cations and B and D are anions. For a double-replacement reaction to occur, at least one of the products must be a gas or water, or a precipitate.
  • Precipitation reactions are one type of double-replacement reaction. An example is
AgNO3(aq) + NaCl(aq) ---> AgCl(s) + NaNO3(aq)
This is what is called a molecular equation, which is a chemical equation in which the compounds are written as if they were molecular substances,even if they exist in solution as ions. Another type chemical equation is an ionic equation. The ionic equation shows soluble ionic compounds as individual ions in solution. Lets rewrite the equation above.
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) ---> AgCl(s) + Na+(aq) + NO3-(aq)
Next we cancel any ions that appear on both sides of the equation.
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) ---> AgCl(s) + Na+(aq) + NO3-(aq)

Remove the cancelled ions (called spectator ions) from the equation and we have
Ag+(aq) + Cl-(aq) ---> AgCl(s)
This is called a net ionic equation.

  • A neutralization reaction is a reaction that occurs between an acid and a base with the formation of an ionic compound and water, if the reaction is in an aqueous solution. This is another type of double-replacement reaction. An example is
HCl(aq) + NaOH(aq) ---> H2O(l) + NaCl(aq)
The net ionic equation would be
H+(aq) + OH-(aq) ---> H2O(l)
Combustion Reactions
A combustion reaction is a reaction in which a substance reacts with oxygen, usually with the rapid release of heat and the production of a flame. Organic compounds usually burn in the oxygen in air to produce carbon dioxide and if the compound contains hydrogen, another product will be water. For example butane burns in air as follows.
2 C4H10(g) + 13 O2(g) ---> 8 CO2(g) + 10 H2O(l)
Summary of the Types of Chemical Reactions
Many chemical reactions can be classified into one of five general types. In a combination reaction, two or more elements or compounds combine chemically to form a new compound. By contrast, a single reactant is the identifying characteristic of adecomposition reaction. Usually in a decomposition reaction, a compound is decomposed into two or more elements or compounds.
When oxygen reacts with a compound composed of the elements carbon and hydrogen, a combustion reaction takes place. If all the carbon in the product is present as carbon dioxide, the reaction is complete combustion. If carbon monoxide is formed, the reaction is incomplete combustion.
The other two types of chemical reactions are single-replacement and double-replacement reactions. In a single-replacement reaction an element replaces another element in a compound. A new compound and a new element are formed. The element that is being displaced must be less active than the element replacing it. This can be determined from the activity series of metals. In a double-replacement reaction two ionic compounds react by exchanging cations to form two new compounds. A double-replacement reaction usually takes place in aqueous solution. Double-replacement reactions are driven by the formation of a precipitate, a gaseous product, or water.


Balancing Chemical Equations








reactant- A chemical substance that is present at the start of a chemical reaction

Decomposition- A reaction where a single compound breaks down into simpler compounds

Products- A substance that is formed as the result of a chemical reaction

Coefficient- A number or symbol multiplied with a variable or an unknown quantity in an algebraic term

Synthesis- Formation of a compound from simpler compounds or elements.

Single Replacement- A chemical reaction in which an element replaces one element in a compound

Doule Replacement- A chemical reaction between compounds in which the elements in the reactants recombine to form two different compounds, each of the products having one element from each of the reactants.

Law of Conservation of matter- The total amount of matter and energy available in the universe is fixed

Exothermic- Reacting that gives off heat to the environment.

Endothermic- Reaction that absorbs heat from its surroundings as the reaction proceeds.

Stoichiometry-The study of the relationships or ratios between two or more substances undergoing a physical or chemical change


Endothermic Reaction Examples :

-Photosynthesis
Photosynthesis is one of the best illustrations of endothermic reactions, occurring in nature. It is a process in which plants use chlorophyll, in the presence of sunlight to convert carbon dioxide and water to glucose and oxygen.

-Melting Ice
When ice melts, it draws heat from the surroundings, rendering the solid form unstable.


-Water Evaporation
Another natural example of an endothermic reaction. Water in the liquid form, uses heat, to convert into the gaseous form of vapor.

-Electrolysis
Electrolysis, involves the separation or decomposition of the original compounds which occurs as a result of the application of an electric current. This is another example of an endothermic reaction.



Exothermic Reaction Examples:

-Combustion/Burning
When you light a match, or a fire, the compounds that burn, release heat into the surroundings, causing an increase in temperature. This is the most basic illustration of an exothermic reaction. Combustion of fuel is another example of an exothermic reaction.

-Neutralization
Many (but not all) neutralization reactions, are exothermic in nature. A neutralization reaction is one that takes place between a acid and a base, to produce salt and water. For example mixing sodium hydroxide (a base) with hydrochloric acid (the acid) will give you a solution of sodium chloride and water, accompanied by an increase in temperature.

-Rusting
The rusting of iron is an example of a spontaneous exothermic reaction

http://www.schooltube.com/video/db41eba5cdbd45fcbe75/Balancing-Chemical-Equations
A video on how to balance chemical equations ^







The concept of balancing equations

Take a look at this chemical word equation:
Aluminium + Oxygen external image rarrow.gif Aluminium Oxide
This is the equation for the burning of aluminium in oxygen. If we convert each of the chemical names into the appropriate symbols, we get the following:
Al + O2 external image rarrow.gif Al2O3
Note that oxygen gas is diatomic, which means that the oxygen atoms, like policemen, go around in pairs. A molecule of aluminium oxide consists of two aluminium atoms combined with three oxygen atoms. Actually, technically the word "molecule" is inappropriate in that previous sentence. The formula simply tells us the ratio of aluminium atoms to oxygen atoms in the compound. In the solid state, the atoms form a giant structure called a crystal lattice rather than individual discrete molecules. When balancing chemical equations, people often refer to the number of species on each side to avoid this problem.
You can see by looking at it that there is something wrong with this equation. If you count the number of atoms of each type on each side, you will see that there is only one aluminium atom on the left side whereas there are two on the right. There are two oxygen atoms on the left side, as compared to three on the right side. This clearly doesn't match.
Left side:
Aluminium atom
Aluminium atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom


Right side:
Aluminium atom
Aluminium atom
Aluminium atom
Aluminium atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom

We can balance the equation by mutiplying the different atoms and molecules on each side by different amounts. Firstly, multiply the aluminium atoms on the left side by 2:
2
2
Al + O2 external image rarrow.gif Al2O3

Left side:
Aluminium atom
Aluminium atom
Aluminium atom
Aluminium atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom


Right side:
Aluminium atom
Aluminium atom
Aluminium atom
Aluminium atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom

Now there are the same number of aluminium atoms on each side of the equation. We could also multiply the number of oxygen molecules on each side by one and a half (1.5), which would give three oxygen atoms on the left side (1.5 x 2 = 3) to match the three oxygen atoms on the right side:
2 Al + 1.5 O2 external image rarrow.gif Al2O3
Left side:
Aluminium atom
Aluminium atom
Aluminium atom
Aluminium atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom


Right side:
Aluminium atom
Aluminium atom
Aluminium atom
Aluminium atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom

This is now balanced, but that 1.5 is a horrible thing to have in an equation - how can you have one and a half molecules? We can solve this problem by multiplying everything throughout by 2:
4
4
Al +
3
3
O2 external image rarrow.gif
2
2
Al2O3

Left side:
Aluminium atom
Aluminium atom
Aluminium atom
Aluminium atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom


Right side:
Aluminium atom
Aluminium atom
Aluminium atom
Aluminium atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom


Aluminium atom
Aluminium atom
Aluminium atom
Aluminium atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom


Aluminium atom
Aluminium atom
Aluminium atom
Aluminium atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom
Oxygen atom

If you count the number of atoms on each side, you will find that there are four aluminium atoms on each side and six oxygen atoms. Sorted!

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


http://www.nclark.net/ChemicalReactions
This website gives you some labs and activities you can do based on chemical reactions




http://www.youtube.com/watch?v=VbIaK6PLrRM&feature=related


Balancing Chemical Equations

This website has a practice test to help with balancing equations:
http://education.jlab.org/elementbalancing/question.php?55994230

II. Four basic types of chemical reactions:
  • two or more elements or compounds may combine to form a more complex compound.
    • Basic form: A + X → AX





    1. Metal + oxygen → metal oxideEX. 2Mg(s) + O2(g) → 2MgO(s)

    2. Nonmetal + oxygen → nonmetallic oxideEX. C(s) + O2(g) → CO2(g)

    3. Metal oxide + water → metallic hydroxideEX. MgO(s) + H2O(l) → Mg(OH)2(s)

    4. Nonmetallic oxide + water → acidEX. CO2(g) + H2O(l) → ; H2CO3(aq)

    5. Metal + nonmetal → saltEX. 2 Na(s) + Cl2(g) → 2NaCl(s)

    6. A few nonmetals combine with each other.EX. 2P(s) + 3Cl2(g) → 2PCl3(g)

**Practice Predicting Products of Synthesis Reactions**Back to the Top





  • A single compound breaks down into its component parts or simpler compounds.
    • Basic form: AX → A + X
Examples of decomposition reactions:

  • ## Metallic carbonates, when heated, form metallic oxides and CO2(g).EX. CaCO3(s) → CaO(s) + CO2(g)


    1. Most metallic hydroxides, when heated, decompose into metallic oxides and water.EX. Ca(OH)2(s) → CaO(s) + H2O(g)


    2. Metallic chlorates, when heated, decompose into metallic chlorides and oxygen.EX. 2KClO3(s) → 2KCl(s) + 3O2(g)


    3. Some acids, when heated, decompose into nonmetallic oxides and water.EX. H2SO4 → H2O(l) + SO3(g)


    4. Some oxides, when heated, decompose.EX. 2HgO(s) → 2Hg(l) + O2(g)


    5. Some decomposition reactions are produced by electricity.EX. 2H2O(l) → 2H2(g) + O2(g)
      EX. 2NaCl(l) → 2Na(s) + Cl2(g)



**Practice Predicting Products of Decomposition Reactions**


  • a more active element takes the place of another element in a compound and sets the less active one free.
    • Basic form: A + BX → AX + B or AX + Y → AY + X


Examples of replacement reactions:


  • ## Replacement of a metal in a compound by a more active metal.EX. Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

    1. Replacement of hydrogen in water by an active metal.EX. 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
      EX. Mg(s) + H2O(g) → MgO(s) + H2(g)

    2. Replacement of hydrogen in acids by active metals.EX. Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

    3. Replacement of nonmetals by more active nonmetals.EX. Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(l)

NOTE: Refer to the **activity series for metals** and nonmetals to predict products of replacement reactions. If the free element is above the element to be replaced in the compound, then the reaction will occurr. If it is below, then no reaction occurs.





    • occurrs between ions in aqueous solution. A reaction will occurr when a pair of ions come together to produce at least one of the following:
      1. a precipitate
      2. a gas
      3. water or some other non-ionized substance.
    • Basic form: AX + BY → AY + BX
  • Examples of ionic reactions:
  • ## Formation of precipitate.EX. NaCl (aq) + AgNO3(aq) → NaNO3(aq) + AgCl(s)
    • EX. BaCl2(aq) + Na2 SO4(aq) → 2NaCl(aq) + BaSO4(s)

    1. Formation of a gas.EX. HCl(aq) + FeS(s) → FeCl2(aq) + H2S(g)

    2. Formation of water. (If the reaction is between an acid and a base it is called a neutralization reaction.)EX. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

    3. Formation of a product which decomposes.EX. CaCO3(s) + HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)

NOTE: Use the **solubility rules** to decide whether a product of an ionic reaction is insoluble in water and will thus form a precipitate. If a compound is soluble in water then it should be shown as being in aqueous solution, or left as separate ions. It is, in fact, often more desirable to show only those ions that are actually taking part in the actual reaction. Equations of this type are called **net ionic equations**.



    • Hydrocarbon (CxHy) + O2(g) → CO2(g) + H2O(g)
    • EX. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
    • EX. 2C4H10(g) + 13O2(g) → 8CO2(g) + 10H2O(g)





Stoichiometry: The relative proportions in which elements form compounds or in which substances react. (From the Greek word for element, stoikheion.)
Stoichiometric coefficient: The number immediately in front of the formula of the chemical substances in the balanced equation. The coefficients may be interpreted as the relative number of particles of reactants and products.
Example: For the above equation two molecules of dihydrogen reacts with one molecule of dioxygen to form two molecules of water.



external image balancesodium.gif
Some Chemical Reactions in Everyday Life
Science being a subject of common interest, it is very intriguing to analyze visual experiments happening in day-to-day life. There are a plethora of products that you use everyday, which are formulated with application of chemical reaction. Say for example; toothpaste, soap, shampoo, cleaning agent, etc. are all results of chemical reactions. Following are some of the most profound chemical reactions, which we encounter in everyday life :

Aerobic Respiration
Do you know indulging in physical movements is associated with a chemical reaction? The process requires energy, which is yielded by aerobic respiration. Over here, respiration helps breaks down glucose (an energy source) into water, carbon dioxide and energy in form of ATP (adenosine triphosphate). The balanced
cellular respiration equation is represented as:

C6H12O6 + 6O2 → 6CO2+ 6H2O + Energy (36 ATPs) C6H12O6 → 2C6H5OH + 2CO2 + Energy 6 CO2+ 6 H2O + Light energy → C6H12O6 + 6 O2 Rusting of Iron Very often, you notice a coating of rust over unpainted iron surfaces, which gradually leads to disintegration of iron. This is nothing, but a chemical phenomenon called rusting. In this case, iron (a very reactive metal) combines with oxygen in presence of water (more precisely, atmospheric moisture), resulting in formation of iron oxides. The chemical reaction behind rusting can be simply represented as:

Fe + O2 + H2O → Fe2O3. XH2O

In addition to the above mentioned examples, you can find a list of chemical reactions in everyday life. Whether you consider cooking, souring, fermenting or burning, there is a chemical reaction accompanying these everyday processes. Thus, it won't be wrong to say learning chemistry and chemical reactions start at home.

external image science-projects-for-kids-chemical-reaction-12.jpg






All chemical reactions can be placed into one of six categories. Here they are, in no particular order:
1)
Combustion: A combustion reaction is when oxygen combines with another compound to form water and carbon dioxide. These reactions are exothermic, meaning they produce heat. An example of this kind of reaction is the burning of napthalene:

C10H8 + 12 O2 ---> 10 CO2 + 4 H2O



2)
Synthesis: A synthesis reaction is when two or more simple compounds combine to form a more complicated one. These reactions come in the general form of:
A + B ---> AB
One example of a synthesis reaction is the combination of iron and sulfur to form iron (II) sulfide:
8 Fe + S8 ---> 8 FeS



3)
Decomposition: A decomposition reaction is the opposite of a synthesis reaction - a complex molecule breaks down to make simpler ones. These reactions come in the general form:
AB ---> A + B
One example of a decomposition reaction is the electrolysis of water to make oxygen and hydrogen gas:
2 H2O ---> 2 H2 + O2



4)
Single displacement: This is when one element trades places with another element in a compound. These reactions come in the general form of:
A + BC ---> AC + B
One example of a single displacement reaction is when magnesium replaces hydrogen in water to make magnesium hydroxide and hydrogen gas:
Mg + 2 H2O ---> Mg(OH)2 + H2



5)
Double displacement: This is when the anions and cations of two different molecules switch places, forming two entirely different compounds. These reactions are in the general form:
AB + CD ---> AD + CB
One example of a double displacement reaction is the reaction of lead (II) nitrate with potassium iodide to form lead (II) iodide and potassium nitrate:
Pb(NO3)2 + 2 KI ---> PbI2 + 2 KNO3



6)
Acid-base: This is a special kind of double displacement reaction that takes place when an acid and base react with each other. The H+ ion in the acid reacts with the OH- ion in the base, causing the formation of water. Generally, the product of this reaction is some ionic salt and water:
HA + BOH ---> H2O + BA
One example of an acid-base reaction is the reaction of hydrobromic acid (HBr) with sodium hydroxide:
HBr + NaOH ---> NaBr + H2O**
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