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Chemical Reactions. Effects of chemical reactions:. Chemical reactions rearrange atoms in the reactants to form new products. The identities and properties of the products are completely different from that of the reactants.

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effects of chemical reactions
Effects of chemical reactions:
  • Chemical reactions rearrange atoms in the reactants to form new products.
  • The identities and properties of the products are completely different from that of the reactants.
  • Production of gases and color changes are signs of chemical reactions.
energy and reactions
Energy and Reactions
  • Energy must be ADDED to BREAK bonds.
  • Energy is RELEASED when bonds are FORMED.
  • Chemical energy is CONSERVED in chemical reactions.
exo vs endo
Exo- vs. Endo-
  • EXOTHERMIC REACTIONS: release energy(More energy is released as the products form bonds than is absorbed to break the bonds in the reactants.)
  • ENDOTHERMIC REACTIONS: absorb energy
chemical equations
Chemical Equations

Chemical equations are used to represent or describe chemical reactions.

For example when hydrogen H2 burns, it reacts with oxygen, O2, in the air to form water. We write the chemical equation for this reaction as follows:

2H2 + O2 —> 2H2O

chemical equations1
Chemical Equations

An equation shows…

Formulas of reactants

Formulas of products

Molar ratios of all compounds in the reaction.

chemical equations2
Chemical Equations

We read the (+) sign as “reacts with” and the arrow (—>) as “produces” or “yields”.

2H2 + O2 —> 2H2O

Reactants

Products

to show physical states of each substance
To show physical states of each substance:
  • (s) or  - solid
  • (l) - liquid
  • (g) or  - gas
  • (aq) - aqueous
    • aqueous means dissolved in water
to show physical states of each substance1
To show physical states of each substance:
  • Consider the reaction of iron with oxygen to form iron (III) oxide, or rust.
  • Fe(s) + O2(g)  Fe2O3(s) (unbalanced)
coefficients subscripts
Coefficients & Subscripts

COEFFICIENTS: numbers in front of compound that represents the number of molecules/moles of that compound

SUBSCRIPTS: small numbers that help define the compound.

2H2SO4

Coefficient

Subscript

slide11

H2O One molecule of water

2H2O Two molecules of water

H2O2 One molecule of Hydrogen Peroxide

slide12

During a chem. rxn.; atoms are rearranged (NOT created or destroyed!)

  • Chemical equations must be balanced to show the relative amounts of all substances.
  • Balanced means: each side of the equations has the same # of atoms of each element.
  • CH4 + O2 —> H2O + CO2Unbalanced
  • CH4 + 2O2 —> 2H2O + CO2 Balanced
in order to balance
In order to balance…
  • Write correct formulas for all reactants and products
  • Reactants  Products
  • Count the number of atoms of each element in reactants & products.
  • Balance one at a time using coefficients.
  • Check for balance
  • Are the coefficients in the lowest possible ratio?
balancing equations
Balancing Equations

NOTE: When balancing equations, you may change coefficients as much as you need to, but you may never change subscripts because you can’t change what substances are involved.

slide15

Balancing equations involves a great deal of “trial and error” at first,

but there are some tricks…

slide16

For example…..

Sodium metal reacts with water to produce sodium hydroxide and hydrogen gas.

Na + H2O —> NaOH + H2

Note that on the product side (right side) there are an odd number of hydrogens. On the reactant side (left side) there is an even number. This implies there must be an even coefficient in front of the NaOH. Lets start with 2

slide17

_Na + _H2O —> 2NaOH + _H2

Now lets balance sodium; we need a 2 in front of the Na…

2Na + _H2O —> 2NaOH + _H2

Now consider hydrogen…

2Na + 2H2O —> 2NaOH + H2

slide18

2Na + 2H2O —> 2NaOH + H2

Check to see if it balances…

2 Na on the left 2 Na on the right

4 hydrogen 2 + 2 = 4 hydrogen

2 oxygen 2 oxygen

the equation is balanced.

examples
Examples

CuCl2(aq) + Al(s)  Cu(s) +AlCl3(aq)

3CuCl2(aq) + 2Al(s) 3Cu(s) +2AlCl3(aq)

(3:2:3:2)

examples1
Examples

Propane, C3H8, burns in oxygen, O2, to form carbon dioxide and water.

C3H8 + O2CO2 + H2O

Balance C – then H – then O

C3H8 + 5O23CO2 + 4H2O

(1:5:3:4)

examples2
Examples

Pentane, C5H12, burns in oxygen, O2, to form carbon dioxide and water.

C5H12 + O2CO2 + H2O

Balance C – then H – then O

C5H12 + 8O25CO2 + 6H2O

(1:8:5:6)

examples3
Examples

Silver nitrate reacts with copper to produce silver and copper (II) nitrate.

AgNO3 + Cu Ag + Cu(NO3)2

2AgNO3 + Cu 2Ag + Cu(NO3)2

(2:1:2:1)

examples4
Examples

Phosphorus reacts with oxygen gas to produce diphosphorus pentoxide.

P + O2 P2O5

4P + 5O2 2P2O5

(4:5:2)

examples5
Examples

C7H14 + O2CO2 + H2O

Balance C – then H – then O

C7H14 + 10½O27CO2 + 7H2O

2C7H14 + 21O214CO2 + 14H2O

(2:21:14:14)

slide25

Types of

Reactions

types of chemical reactions
Types of Chemical Reactions
  • Synthesis / Combination
  • Decomposition
  • Single Replacement
  • Double Replacement
  • Combustion
synthesis combination reactions
Synthesis / Combination Reactions

Definition: Reaction where two or more substances react to form a single substance.

A + BAB

Examples:

2K(s) + Cl2(g) 2KCl(s)

SO2(g) + H2O(l) H2SO3(aq)

decomposition reactions
Decomposition Reactions

Definition: Reaction where a single compound is broken down into two or more products.

AB A + B

Examples:

2H2O(l) 2H2(g) + O2(g)

CaCO3 CaO + CO2

single replacement reactions
Single-Replacement Reactions

Definition: Reaction where atoms of one element replace atoms of a second element in a compound.

XA + B BA + X

Note: A reactive metal will replace any metal listed below it in the activity series. Generally, nonmetal replacement is limited to the halogens. The activity of the halogens decreases as you go down Group 7A of the periodic table. See handout.

Examples:

2AgNO3 + MgMg(NO3)2+2Ag

Mg+LiNO3 no reaction

slide30

Li

K

Ca

Na

Mg

Al

Zn

Fe

Pb

(H)*

Cu

Hg

Ag

Activity Series

Increasing Activity

Any element will replace any element below it.

*Metals from Li to Na will replace H from acids and water; from Mg to Pb they will replace H from acids only

for example
For Example…

Ca + MgO  CaO + Mg

The Ca will replace the Mg because Ca is more active than Mg. That is to say…Ca is above Mg on the activity list.

double replacement reactions
Double-Replacement Reactions

Definition: Reaction that involves an exchange of positive ions between two compounds.

XA + BYBA + XY

Note: These reactions generally take place between two ionic compounds in aqueous solution, and are often characterized by one of the products coming out of solution in some way.

Examples:

2NaCN(aq)+H2SO4(aq) 2HCN(g)+Na2SO4(aq)

Na2S(aq)+Cd(NO3)2(aq) CdS(s)+2NaNO3(aq)

combustion reactions
Combustion Reactions

Definition: Reaction where an element or a compound reacts with oxygen, often producing energy in the form of heat and light.

Examples:

CH4+2O2 CO2+2H2O + heat + light

2Mg(s)+O2(g) 2MgO(s)

combustion of hydrocarbons
Combustion of Hydrocarbons

If the reactant is a hydrocarbon, the products are always carbon dioxide and water.

CH4 + 2O2  CO2 + 2H2O

ionic equations
Ionic Equations
  • When a soluble substance is dissolved in water, the substance often breaks into ions. This solution is said to be an aqueous solution.
  • Pb(NO3)2(aq) Pb2+ + 2NO3-
  • NaI(aq)  Na+ + I-
ionic equations1
Ionic Equations
  • Consider the reaction…
    • Pb(NO3)2(aq) + NaI(aq) PbI2(s) + NaNO3(aq)
  • What is really going on is…
    • Pb2+ + NO3- + Na+ + I- PbI2(s) + Na+ + NO3-
  • Note that the Na+ ion and the NO3- ion are not reacting. They are said to be spectator ions.
net ionic equations
Net Ionic Equations
  • It is often useful to write an equation showing only the species that are actually reacting. This is called a net ionic equation. It does not show the spectator ions.

Pb2+ + NO3- + Na+ + 2I- PbI2(s) + Na+ + NO3-

becomes….

Pb2+ + 2I- PbI2(s)

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