Chapter 5:Classification and Balancing of Chemical Reactions

• Law of conservation of mass states that matter can neither be created nor destroyed in a chemical reaction.

• The bonds between atoms in the reactants are rearranged to form new compounds in chemical reactions, but @@none of the atoms disappear and no new ones are formed.@@

• Chemical equations must be balanced equations, meaning that the numbers and kinds of atoms must be the same on both sides of the reaction arrow.

• Most chemical equations can be balanced by following the four step approach:

  • STEP 1: Write an unbalanced equation, using the correct formulas for all given reactants and products.
  • STEP 2: Add appropriate coefficients to balance the numbers of atoms of each element.
  • STEP 3: Check the equation to make sure the numbers and kinds of atoms on both sides of the equation are the same.
  • STEP 4: Make sure the coefficients are reduced to their lowest whole-number values.

• Ionic reactions can be classified into three types:

  1)precipitation reactions

  2)acid-base neutralization reactions

  3)oxidation-reduction reactions.

• Precipitation reactions are processes in which an @@insoluble solid called a precipitate forms when reactants are combined in aqueous solution@@. Most precipitations take place when the anions and cations of two ionic compounds change partners.

• To predict whether a precipitation reaction will occur upon mixing aqueous solutions of two ionic compounds, you must know the solubilities of the potential products—how much of each compound will dissolve in a given amount of solvent at a given temperature.

  • If a substance has a ==low solubility== in water, then it is ==likely to precipitate== from an aqueous solution. If a substance has a ^^high solubility^^ in water, then ^^no precipitate^^ will form.

• General Rules on Solubility

  RULE 1: A compound is probably soluble if it contains one of the following cations:

  • Group 1A cation: Li +, Na +, K+, Rb+, Cs+
  • Ammonium ion: NH4+

• RULE 2: A compound is probably soluble if it contains one of the following anions:

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  • Halide: Cl-, Br -, and I

  except Ag+, Hg2+, and Pb2+ compounds

  • Nitrate 1NO3 -2, perchlorate 1ClO4 -2, acetate 1CH3CO2 -2, and sulfate 1SO4 2-2
  • except Ba2+, Hg2+, and Pb2+ sulfates.

• Acid-base neutralization reactions are processes in which an acid reacts with a base to yield water plus an ionic compound called a salt.

• When acids and bases are mixed in the correct proportion, ^^both acidic and basic properties disappear because of a neutralization reaction^^.

  • The most common kind of neutralization reaction occurs between an acid (generalized as HA) and a metal hydroxide (generalized as MOH) to yield water and a salt.
  • The H + ion from the acid combines with the OH- ion from the base to give neutral H2O, whereas the anion from the acid 1A-2 combines with the cation from the base 1M+2 to give the salt.

• An oxidation is defined as the loss of one or more electrons by an atom, and a reduction is the gain of one or more electrons. Thus, an oxidation-reduction reaction, or redox reaction, is one in which electrons are transferred from one atom to another.

• Oxidation and reduction %%always occur together.%% Whenever one substance loses an electron (is oxidized), another substance must gain that electron (be reduced).

  • . The substance that gives up an electron and causes the reduction is called a reducing agent.
  • The substance that gains an electron and causes the oxidation is called an oxidation agent.

• Reducing agent :

  • Loses one or more electrons
  • Causes reduction
  • Undergoes oxidation
  • Becomes more positive (less negative) (May gain oxygen atoms)

• Oxidizing agent :

  • Gains one or more electrons
  • Causes oxidation
  • Undergoes reduction
  • Becomes more negative (less positive) (May lose oxygen atoms).

• The alkali metals and alkaline earth metals are the most powerful reducing agents (electron donors).

  • This is due in part to the fact that alkali metals and alkaline earth metals have low ionization energies.

• The reactive non-metals at the top right of the periodic table have the ==highest ionization energies== and are extremely weak reducing agents but powerful oxidizing agents (electron acceptors).

• In reactions involving metals and non-metals, metals tend to lose electrons while non-metals tend to gain electrons.

• In reactions involving non-metals, the “more metallic” element (farther down and/or to the left in the periodic table) tends to lose electrons, and the “less metallic” element (up and/or to the right) tends to gain electrons

• Corrosion is the deterioration of a metal by oxidation, such as the rusting of iron in moist air. The economic consequences of rusting are enormous: it has been estimated that up to one-fourth of the iron produced in the United States is used to replace bridges, buildings, and other structures that have been destroyed by corrosion.

• Combustion is the burning of a fuel by rapid oxidation with oxygen in air.

• Respiration is the process of breathing and using oxygen for the many biological redox reactions that provide the energy required by living organisms.

• Bleaching makes use of redox reactions to decolorize or lighten coloured material.

• Metallurgy, the science of extracting and purifying metals from their ores, makes use of numerous redox processes.

  • Worldwide, approximately 800 million tons of iron are produced each year by reduction of the mineral hematite, Fe2O3, with carbon monoxide.

• Oxidation number (or oxidation state), which indicates whether the atom is neutral, electron-rich, or electron-poor.

• Oxidation numbers do not necessarily imply ionic charges.

  • An atom in its elemental state has an oxidation number of 0.
  • A monatomic ion has an oxidation number equal to its charge.
  • In a molecular compound, an atom usually has the same oxidation number it would have if it were a monatomic ion.
  • The sum of the oxidation numbers in a neutral compound is 0.

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