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Chemical Reactions

Read Section 8.1

Information available from a chemical equation:

                           Reactants --------> Products


__________________ - starting substances in a chemical reaction

__________________ - substances formed in a chemical reaction

----------> means “yields”


Example:

Word equation: hydrogen plus oxygen yields water


Chemical equation:        H2       +       O2        -------->      H2O


The Law of Conservation of Mass tells us that we can neither create
nor destroy matter in a chemical reaction, ...


         H2        +        O2        --------->        H2O


But you can change their partners.

 

             2 H2    +            O2            ---------->        2 H2O

2 molecules H2 + 1 molecules O2 -----------> 2 molecules H2O

2 moles H2       +        1 mole O2  -----------> 2 moles H2O

4.04 g H2           +         32.00 g O2 -----------> 36.04 g H2O



Symbols used in writing chemical equations 

+                used to separate two reactants or two products
-------->   “yields”, separates reactants from products

-------->     heat applied to the reaction

catalyst
-------->      a catalyst is required (gets reaction started faster)

(s) _________________

(l) _________________

(g) __________

(aq) _______________; substance is dissolved in water


                 gas being produced by the reaction



                 solid being produced by the reaction



Writing and Balancing Chemical Equations

1. Make sure you have the correct formulas for all the reactants and products in the reaction. Are elements present in the reaction that should be written as diatomic molecules? 


2. Write the reactants on the left and the products on the right with a yield sign in between. If there are two or more reactants or products involved, use plus signs to separate their formulas.


3. Compare the number of atoms of each element in the reactants with the numbers of atoms of each element in the products. If the numbers on both sides are equal you are finished; the equation is balanced.


4. Balance the elements on each side of the equation by placing a coefficient, a whole number, in front of the chemical formulas as appropriate. Start with elements that appear only once on each side of the equation. Never change subscripts to balance the equation. That would alter the chemical formula of a substance, making it another substance altogether. 


5. Double-check each kind of atom to be sure that the equation is balanced.


6. Check to see that the coefficients are in the lowest possible ratio. For example, 12:6:4 should be reduced to 6:3:2.


Writing and Balancing Chemical Equations

Strategies for balancing equations:

1. EVEN/ODD

Example:  Write the equation for the reaction of iron and oxygen to form iron (III) oxide
1.  Make sure you have the correct formula for all reactants and products.
 iron                 oxygen 

iron (III) oxide


2.  Write the reactants on the left and the products on the right with a yield sign in between.
         Fe +         O2 ------------> ______________

3. Equation is not balanced with an even number of oxygen atoms on the left of the equation and odd number of O atoms on the right.

4.  Balance using coefficients 

In the reaction,

              Fe       +     ________      --------->    _________

the number of oxygen atoms on the left must be even. Begin by making the number of oxygen atoms on the right even, then balance the iron atoms.

5.  double-check
6.  lowest possible ratio

Example:  Write an equation for the formation of ammonia from the reaction of nitrogen and hydrogen.

Substances:
 nitrogen  _______            hydrogen  _______

ammonia      ____________

 
Skeleton equation - chemical equation that does not  indicate the relative amount of reactants and products

                           _____________  +  ____________  -->     _____________

Equation is not balanced

Draw a box around reactants to remind yourself that you cannot change the subscripts; the boxes are off-limits - you cannot change anything inside the box


__________________________________________

use the even/odd strategy

__________________________________________

Double-check number of atoms

Check that coefficients are in the lowest possible ratio


Example:  Write the equation showing that sodium sulfide is formed from sodium and sulfur. 
Substances:
sodium  ______          sulfur  ______

sodium sulfide

Skeleton equation - chemical equation that does not indicate 
the relative amount of reactants and products


      _______ + ________ ------->  ____________


Equation is not balanced
 
      ________ + ________ ---------> _________


Double check number of atoms on both sides of the equation

Check that coefficients are in lowest possible ratio


2. INTACT POLYATOMIC ION

Example:   Write the equation for the formation of iron (III) acetate and lead from iron and lead (II) acetate
Substances:
1) lead ______

iron _______


iron (III) acetate




lead (II) acetate

 

Skeleton equation: 



Equation not balanced



the atoms in the polyatomic ion, acetate, can be balanced 
as a group

also use even/odd strategy






Is the equation balanced?

Are the coefficients in the lowest possible ratio?


Example:  Sodium chloride and potassium nitrate react to form sodium nitrate and potassium chloride.

Substances:

sodium chloride  _______________                             sodium nitrate ________________

potassium nitrate _______________                            potassium chloride  _____________

Equation:

 

 





3.  COMBUSTION OF HYDROCARBONS 

The number of hydrogen atoms in the reactants must be divisible by four.

Example:  Write an equation for the combustion of acetylene to form carbon dioxide and water.
Substances:
CO2
O2
C2H2
H2
Skeleton equation:
 __________ + _______ ---------> ________ + _________

Balance: 1) Make the number of hydrogen atoms on the reactant side - a number divisible by 4

              2)  Balance carbon atoms        3)    Balance oxygen atoms

Begin by making the number of hydogen atoms on the left 4 by choosing a coefficient of 2 for __________; the equation then balances easy 


___________ + _________ -------> _________ + __________


Is the equation balanced?
Are the coefficients in the lowest possible ratio?

When the number of hydrogen atoms is already divisible by 4, this technique is not needed, as in:

__________ + _________--------> _________ + _________

Read Section 8.2

Types of Chemical Reactions

I    Combination
II   Decomposition
III Single-Replacement
IV Double-Replacement
V   Combustion

I.   Combination Reaction

Two or more substances combine to form a single product
Reactants are elements or compounds

A. Group A metal cations combine with non-metal anions to form an ionic compound
Examples:
2 Ca (s) + O2 (g) ------------> 2 CaO(s)

4 Al (s) + 3 O2 (g) ------------> 2 Al2O3 (s)
Practice:
zinc   +    nitrogen  -->

 

sodium  +   chloride  -->



B. Two non-metals react with each other - or - 
     a transition metal reacts with a non-metal

Example:  non-metal and non-metal

2 Cl2 (g)  +  O2        ------->      2 Cl2O (g) dichlorine monoxide 

Cl2 (g)   +   2 O2 (g)  ------>       2 ClO2 (g) chlorine dioxide

Example:  transition metal and non-metal

2 Cu (s) + O2 (g) ----------> 2 CuO (s) copper (II) oxide

4 Cu (s) + O2 (g) ----------> 2 Cu2O (s) copper (I) oxide


C. Non-metal oxides react with water to form acids 
(compounds that produce hydrogen ions, H+, when
combined with water)

Example:        CO2 (g) + H2O (l) ---------> H2CO3 (aq) carbonic acid

Practice:       SO3   +    H2O    --->   
                    N2O5      +     H2O   --->



D. Metallic oxides react with water to form bases
(compounds which contain hydroxide ions, OH-)

Example:      Na2O (s) + H2O (l) ----------> 2 NaOH (aq)

Practice:         Li2O    +     H2O   -->
                      BaO       +      H2O    -->


II. Decomposition Reactions

In decomposition reactions, it is more difficult to determine the products. Compounds that contain only 2 kinds of elements (binary compounds) will separate into those two elements.
A.  Binary compounds separate into two elements

Example:     2  Ag2O (s)       -->      4 Ag (s)      +    O2 (g)
Practice:

B.  Some acids will decompose into water and a nonmetal oxide (if heated)

Example:     H2CO3 (aq)     ----->    CO2   +    H2
Practice:

C.  Some bases will decompose into water and a metal oxide (if heated)

Example:    Ba(OH)2  (aq)   -->      BaO (s)     +    H2O (l)

Practice:

Compounds containing three or more elements are less predictable.

III. Single-Replacement Reactions

One element replaces a second element in a compound; the element that begins the reaction alone replaces another
element that begins the reaction as part of a compound;  the element that ends up alone in the products is the one
that was originally part of the reactants

A.  An active metal will replace the metallic ion in a compound of a less active metal from a 
compound 

Example:    Cu (s)    +     AgNO3 (aq)   --->      Ag (s)    +      CuNO3 (aq)

Practice:      Fe (s) + CuSO4 (aq)          --->

B.  Active metals such as Zn, Pb, and Al will replace H in acids to give a salt (ionic compound) and hydrogen gas

Example:

C.  Halogens will replace less active halogens.

Example:


IV. Double-Replacement Reactions

Positive ions exchange places between two compounds
Three indications of reaction
1)   precipitate forms
2)   gas given off
3)   water produced when acid mixed with base

AX + BY --------> AY + BX

Example:

_________ (aq) + __________ (aq) --------> __________ (s) + 2 NaCl (aq) 

Practice:

2 KI (aq) + _________  (aq) --------> 2 ________ (aq) + _________ (s)


V. Combustion
Substance combines with oxygen
Products of complete combustion are carbon dioxide and water 
Exothermic

CxHy + _______ -------> ________+ ________

Example:   the combustion of octane in gasoline


2 _________ (l) + 25 _______ (g) ----> 16 ______ (g) + 18 _____ (g) 


Example:      the combustion of heptane: _________(g) + _____ (g) ------>


_________ (g) + ________ (g) -------> ________ (g) + _______(g)

HOMEPAGE