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Reactions - PowerPoint PPT Presentation

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Reactions. CH.3 Balancing Reactions Reaction Types. A Chemical Reaction. Reactants Products Types of Chemical reactions: There are many types of reactions but most can be classified into a few simple reactions types. How do you know a chemical reaction has occurred? .

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Reactions l.jpg



Balancing Reactions

Reaction Types

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A Chemical Reaction

  • Reactants Products

  • Types of Chemical reactions:

  • There are many types of reactions but

  • most can be classified into a few simple

  • reactions types.

  • How do you know a chemical reaction has occurred?

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Chemical equations

  • Chemist’s shorthand to describe a reaction.

  • It shows:

    • All reactants and products

    • The state of all substances

    • Any conditions used in the reaction

      • CaCO3 (s) CaO (s) + CO2 (g)

Reactant Products

A balanced equation shows the relationship

between the quantities of all reactants and products.

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Balancing chemical equations

  • Each side of a chemical equation must have the same number of each type of atom.

    • CaCO3 (s) CaO (s) + CO2 (g)

    • Reactants Products

      • 1 Ca 1 Ca

      • 1 C 1 C

      • 3 O 3 O

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Balancing chemical equations

  • Step 1 Count the number of atoms of each element on each side of the equation.

  • Step 2 Determine which atom numbers are not balanced.

  • Step 3 Balance one atom at a time by using coefficients in front of one or more substances.

  • Step 4 Repeat steps 1-3 until everything is balanced.

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Chemical Changes

Temperature Changes


Gas formation


Color changes

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Types of Chemical Reactions

  • Reactions involving electron transfer

    • synthesis or combination

    • decomposition

    • Electrochemical or single replacement

    • Combustion of hydrocarbons.

  • Reactions that involve rearrangement but not necessarily involve electron transfer

    • Metathesis or Double Substitution

    • Precipitation reactions and

    • Acid Base Reactions

  • Polymerization reaction

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Properties ofaqueous solutions

  • There are two general classes of solutes.

  • Electrolytic

    • ionic compounds in polar solvents

    • dissociate in solution to make ions

    • conduct electricity

    • may be strong (100% dissociation) or weak (less than 100%)

  • Nonelectrolytic

    • do not conduct electricity

    • solute is dispersed but does not dissociate

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Dissolving ionic compounds

  • When an ionic solid dissolves in water, the solvent removes ions from the crystal.

  • NaCl + H2O  Na+(aq) + Cl- (aq)

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Dissolving covalent compounds

  • Covalent compounds do not dissociate.

  • C6H12O6 C6H12O6 (aq)

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Ionic equations

  • When ionic substances dissolve in water, they dissociate into ions.

  • AgNO3 Ag++ NO3-

  • KClK+ + Cl-

  • When a reaction occurs, only some of the ions are actually involved in the reaction.

  • Ag++ NO3- +K+ + Cl- AgCl(s) + K+ + NO3-

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Ionic equations

  • To help make the reaction easier to see, we commonly list only the species actually involved in the reaction.

  • Molecular equation

  • KCl + AgNO3AgCl(s) + KNO3

  • Full ionic equation

  • Ag++ NO3- +K+ + Cl- AgCl(s) + K+ + NO3-

  • Net ionic equation

  • Ag++Cl-AgCl(s)

  • NO3- and K+ are referred to as spectator ions.

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Some simple solubility rules

  • All acids are soluble.

  • All Na+, K+ and NH4+ salts are soluble.

  • All nitrate and acetate salts are soluble.

  • All chlorides except AgCl and Hg2Cl2 are soluble. PbCl2 is slightly soluble.

  • All sulfates are soluble except PbSO4, Hg2SO4, SrSO4 and BaSO4. Ag2SO4 and CaSO4 are slightly soluble.

  • All sulfides are insoluble except those of the Group IA (1), IIA (2) and ammonium sulfide.

  • All hydroxides are insoluble except those of the group IA(1) and Ba(OH)2. Sr(OH)2 and Ca(OH)2 are slightly soluble.

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Metathesis Reactions

  • Precipitation reactions

    • the formation of a solid upon mixing two solutions.

  • Gas formation

    • The formation of a gas when two mixtures are reacted

  • Acid Base titration

    • The neutralization of acids and bases to form a salt in water.

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  • NaCl + AgNO3 = AgCl (ppt) + NaNO3

  • Pb(NO3)2 + K2CrO4 = PbCrO4 (ppt) + KNO3

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Gas Formation

  • Vinegar + Baking soda = Carbon dioxide

  • CH3COOH + NaHCO3 = NaCH3COO + H2O + CO2

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Acid Base Neutralization

  • Vinegar + Ammonia = Salt + water


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Common REDOX Reactions

  • Combustion

  • Corrosion

  • Photosynthesis

  • Kreb’s Cycle

  • Synthesis and Decomposition

  • Single Replacement

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Oxidation numbers and the periodic table

  • Some observed trends in compounds.

  • Metals have positive oxidation numbers.

  • Transition metals typically have more than one oxidation number.

  • Nonmetals and semimetals have both positive and negative oxidation numbers.

  • No element exists in a compound with an oxidation number greater than +8.

  • The most negative oxidation numbers equals 8 - the group number

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Oxidation Numbers

  • 1. An atom in its elemental state has an oxidation number of 0.

  • 2. An atom in a monatomic ion has an oxidation number identical to its charge.

  • 3. An atom in a polyatomic ion or in a molecular compound usually has the same oxidation number it would have it were a monatomic ion.

  • 4. The sum of the oxidation numbers is 0 for a neutral compound and equal to the net charge for a polyatomic ion.

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Identifying oxidation-reduction reactions.

  • Oxidation-Reduction - REDOX

  • A chemical reaction where there is a net change in the oxidation number of one or more species.

  • Both an oxidation and a reduction must occur during the reaction.

Mg (s) + Cl2 (g) MgCl2 (s)

Here the oxidation number of Mg has changed from

zero to +2. Cl has changed from zero to -1.

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REDOX reactions

  • Oxidation

  • An increase in oxidation number.

  • Reduction

  • A decrease in oxidation number.

  • If the oxidation number of any element changes in the course of a reaction, the reaction is oxidation-reduction.

  • Example.

  • 2 Fe(NO3)3 (aq) + Zn(s) 2 Fe(NO3)2 (aq) + Zn(NO3)2 (aq)

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Half Reactions

  • The reactions that show the loss or gain of electrons.

  • 2 Mg + O2 2 MgO

  • What are the oxidation states of each of the atoms in the above reaction?

  • Mg  Mg+2

  • O2  2 O2-

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  • 2 Mg + O2 2 MgO

  • Mg  Mg+2

  • 12 p, 12 e 12 p, 10 e

  • Mg  Mg+2 + 2 e

  • The loss of electrons

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Single replacement reaction

  • Where one element displaces another in a chemical compound.

  • H2 + CuO Cu + H2O

  • In this example, hydrogen replaces copper.

  • This type of reaction always involves oxidation and reduction (REDOX).

  • Since one species is replacing another, there are no spectator ions.

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Synthesis Reaction

  • The formation of a new compound from several substances

    • 2 H2 + O2 = 2 H2O

    • 2 Na + Cl2 = 2 NaCl

  • CaO (lime water) + CO2 = CaCO3 (s)

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Decomposition Reactions

  • HgO = Hg + O2

  • H2O2= H2O + O2

  • with MnO2 catalyst

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Combustion process and chemical processes

  • The burning of a substance in the presence of oxygen.

  • Gasoline burns when oxygen is taken into your engine, the products of the reaction are carbon dioxide and water

  • C8H18 + O2 = CO2 + H2O

  • Fe + O2 = Fe2O3 (rust)