Types of Chemical Reactions

Transcription

1 Types of Chemical Reactions 1) Combination (Synthesis) Reaction 2) Decomposition 3) Single Replacement 4) Double Replacement 5) Combustion 6) Oxidation-Reduction (Redox) Combination (Synthesis) Reactions 2 or more substances combine to make a NEW, SINGLE substance. For examples, 1. 1S(s) + 1O 2 (g) 1SO 2 (g) 2. 2S(s) + 3O 2 (g) 2SO 3 (g) 3. 2Fe(s) + 3S(s) 1Fe 2 S 3 (s)

2 Decomposition Reactions A SINGLE compound is broken down into 2 or more products. Sometimes the products of decomposition reactions are difficult to predict. Examples 1. 2H 2 O(l) 2H 2 (g) + 1O 2 (g) 2. 2HgO(s) 2Hg(l) + 1O 2 (g) 3. 2KClO 3 (s) 2KCl(s) + 3O 2 (g) Single Replacement Reactions Atoms of an element replace the atoms of a second element in a compound.

3 Whether one metal will displace another metal from a compound can be determined by the relative reactivities of the two elements. Examples 1. Mg(s) + Zn(NO 3 ) 2 (aq) Mg(NO 3 ) 2 (aq) + Zn(s) 2. Mg(s) + 2AgNO 3 (aq) Mg(NO 3 ) 2 (aq) + 2Ag(s) 3. Mg(s) + LiNO 3 (aq) No Reaction Double Replacement Reactions These reactions involve an exchange of positive ions between 2 compounds. These reactions often have 1 product coming out of solution in some way (i.e. usually forming a precipitate).

4 Negative Ions (Anions) + Positive Ions (Cations) = Solubility of compounds in water any anion + alkali ions (Li +,Na +,K +,Rb +,Cs +,Fr + ) = Soluble any anion + hydrogen ion [H + (aq)] = Soluble any anion + ammonium ion (NH + 4 ) = Soluble nitrate - NO 3 + any cation = Soluble acetate (C 2 H 3 O - 2 ) + any cation = Soluble Chloride (Cl - ), Bromide (Br - ), Iodide (I - ) Sulfate (SO 4 2- ) + silver (Ag+ ), lead (Pb 2+ ), mercury (Hg 2+ ), copper (Cu + ), thallium (Tl + = ) Insoluble + any other cation = Soluble calcium (Ca 2+ ), strontium (Sr 2+ ), barium + (Ba 2+ ), silver (Ag + ), lead (Pb 2+ ), radium = Insoluble (Ra 2+ ) + any other cation = Soluble alkali ions (Li +,Na +,K +,Rb +,Cs +,Fr + ), alkali earth metals Sulfide + S 2- (Be 2+,Mg 2+,Ca 2+,Sr 2+,Ba 2+,Ra 2+ = Soluble ), and H + + (aq), NH 4 + any other cation = Insoluble alkali ions (Li +,Na +,K +,Rb +,Cs +,Fr + ), Hydroxide + OH - H + (aq),nh + 4,Sr 2+,Ba 2+,Ra 2+,Tl + = Soluble + any other cation = Insoluble Phosphate, PO 4 3- Carbonate, CO 3 2-, Sulfite, SO alkali ions (Li +,Na +,K +,Rb +,Cs +,Fr + ), H + (aq),nh 4 + = Soluble + any other cation = Insoluble

5 Examples 1. Na 2 S(aq) + Cd(NO 3 ) 2 (aq) 2NaNO 3 (aq) + CdS(s) 2. 2NaCN(aq) + H 2 SO 4 (aq) 2HCN(g) + Na 2 SO 4 (aq) 3. BaCl 2 (aq) + K 2 CO 3 (aq) BaCO 3 (s) + 2KCl(aq) Combustion Reactions These reactions involve an element or a compound reacting with oxygen, often producing energy in the form of heat and/or light. Example 2Mg(s) +O 2 (g) 2MgO(s)

6 Hydrocarbons (compounds containing hydrogen and carbon burn producing the products of CO 2 + H 2 O Examples 1) 2C 6 H 6 (l) +15O 2 (g) 12CO 2 (g) + 6H 2 O(g) 2) 1C 2 H 5 OH(l) + 1O 2 (g) 2CO 2 (g) + 3H 2 O(g)

7 OXIDATION REDUCTION REACTIONS Common - involves an exchange of electrons between 2 species Also know as REDOX reactions Many chemical reactions are actually redox reactions One species loses e- (oxidized) One species gains e- (reduced) The species that loses electrons The species that gains electrons OXIDIZED REDUCED

8 Example: Redox Rxn. Zinc(s) + HC1. The net ionic eqn: Zn(s) + 2H + (aq) Reduced Zn 2+ (aq) + H 2 (g) Oxidized Zn (s) Zn 2 + (aq) + 2e - (oxidation half-rxn) 2H + (aq) + 2e - H 2 (g) (reduction half-rxn) Note: 1) No net in the # of e - 2) Oxidation and reduction must occur together.

9 Oxidizing agent - species that accepts e - Reduction agent - species that donates e - Oxidation Number Charge of the element Rules to determine oxidation #: 1) The ox. # of an element in an elementary state = 0 Example: Na 0, O 2 0 2) Ox. # of an element in a monotomic ion = charge of that ion

10 3) Certain compounds always have the same ox. # (ex. Groups 1&2) Many vary, some notable exceptions are: Oxygen: usually O 2-, but it can be O 1- as in H 2 O 2. Hydrogen: Usually H 1+, but it can be H 1- as in NaH. 4) The sum of the ox. # in a neutral compd. = 0 * In a polyatomic ion, the ox. # is equal to the charge of the ion.

11 Example: What is the ox. # of "S" in K 2 SO 3? K +1 0=2((+1) + x + 3(-2) 0-2 0=+2 + x = x = x

12 BALANCING REDOX EQUATIONS Mg(s) + C1 2 (g) MgC1 2 (s) Mg is oxidized 0 +2 C1 is reduced 0-1 P 4 (s) + 5O 2 (g) 2P 2 O 5 (s) P is oxidized 0 +5 O is reduced 0-2 Oxidation in ox. # Reduction in ox. #

13 Often redox rxns. cannot be balanced by balancing the # of atoms. In these situations, the Half-Reaction Method is used (in aqueous solutions). These are the steps: 1) Split the rxn. into 2 half-rxn with (reduction & oxidation) 2) Balance one half-rxn with respect to both atoms & charge. 3) Balance the other half-rxn. 4) Combine the two half-rxn to eliminate electrons.

14 Example: Balance the following: Fe +3 (aq) + C1-1 (aq) Fe(s) + C1 2 (g) Split rxn. Fe +3 (aq) Fe(s ) Fe +3 (aq)+3e - Fe(s) C1-1 (aq) C1 2 (g) 2C1-1 (aq) C1 2 (g)+2e - 2 Fe +3 (aq)+3e - Fe(s) 3 2C1-1 (aq) C1 2 (g)+2e - ) 2Fe +3 (aq)+6e - 2Fe(s) 6Cl - (aq) 3Cl 2 (g)+6e - 2Fe +3 (aq) + 6Cl - (aq) 3Cl 2 (g) + 2Fe(s)