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Types of Chemical Reactions

Synthesis (combination) reactions – two or more substances combine to form a single substance. A + B = AB A combine with to form ionic compounds. (To get the correct formula you must know the charges of the cations and anions that the and form). 2K(s) + Cl2(g)  2KCl(s)

When two nonmetals react in a synthesis reaction, there is more than one possible product. – you must be given the product name S(s) + O2(g)  SO2(g) or 2S(s) + 3O2  2SO3(g)

When a and a nonmetal react in a synthesis reaction, there may be more than one possible product because the transition metal could form more than one cation. – you must be given at least the charge on the cation. Fe(s) + S(s)  FeS(s) (II) 2Fe(s) + 3S(s)  Fe2S3 Iron (III) sulfide.

Nonmetal (nonmetal with ) react with to produce an (H+ compound). SO2(g) + H2O(l)  H2SO3(aq)

Metallic oxides (metal with oxygen) react with water to give a (OH- compounds). Use ionic charges to write formula of product. CaO(s) + H2O(l)  Ca(OH)2(aq)

A metal and nonmetal oxide combine to form a . CO2(g) + Na2O(cr)  Na2CO3(s)

Decomposition Reactions – a single compound is broken down into two or more products and usually require energy (heat, light or electricity) to take place.

AB A + B

When a binary (2 elements only) compound breaks down, the products will be those 2 elements. electricity

H2O(l) H2(g) + O2(g) . When some are heated, they decompose to form water and nonmetal oxide

H2CO3(aq) CO2(g) + H2O(l)

When some metal (metal combined with OH-)are heated, they decompose to form a metallic oxide and water.

Ca(OH)2 CaO(s) + H2O(g)

2- When some metallic (metal combined with CO3 ) are heated, they decompose to form a metallic oxide and dioxide.

Li2CO3(s) Li2O(s) + CO2(g)

- When metallic chlorates (metal combined with ClO3 )are heated, they decompose to form metallic and oxygen.

KClO3(s) 2KCl(s) + 3O2(g)

Single Replacement/Displacement Reactions – one element replaces a second element in a compound. A + BC  B + AC (If A is a metal) or A + BC  C + BA (If A is a nonmetal)

Whether one metal will displace another metal from a compound can be determined by the relative reactivities of the two metals. An activity series lists the reactivities of some metals.

Activity Series of Metals Name Symbol Li

Potassium K Ca Na Mg Aluminum Al Zn Iron Fe

Decreasing Reactivity Pb () (H)* Cu Hg Ag *Metals from Li to Na will replace H from acids and water; from Mg to Pb they will replace H from acids only.

A nonmetal can also replace another nonmetal from a compound, usually a . The activity of the decreases as you go down group 17 of the . F Cl Br I At

Combustion Reactions – an element or compound reacts with oxygen, often producing energy as heat and light.

Commonly involves hydrocarbons (compounds that only contain H and C) . Complete combustion – forms carbon dioxide and water. CxHy + O2  CO2 + H2O . Incomplete combustion – reaction runs out of oxygen, then elemental carbon and carbon may be additional products.

Combustion reactions between elements and oxygen also exist. 2Mg(s) + O2(g)  2MgO(s)

Double Replacement/Displacement Reactions – Involve an exchange of two anions between two reacting ionic compounds. AB + CD  AD + CB For a double- displacement reaction to occur, one of the following is usually true: 1. One product precipitates out of . 2. One product is a gas that bubbles out of the mixture. 3. One product is a molecular compound, usually water.

To describe double displacement reactions more clearly we use net ionic equations.

Net Ionic Equations Most ionic compounds when dissolved in water dissociate, or separate, into their anions and cations.

Molecular equation: AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)

1+ 3- When AgNO3 is dissolved in water, it separates into Ag cations and NO anions. The other aqueous compounds dissociate also.

Complete Ionic Equation: + - + - + - Ag (aq) + NO3 (aq) + Na (aq) + Cl (aq)  AgCl(s) + Na (aq) + NO3 (aq)

Ions that appear on both sides of the reaction are not directly involved in the reaction and are called spectator . These ions may be canceled out of the reaction.

Net Ionic Equation: + - Ag (aq) + Cl (aq)  AgCl(s)

 Note: When writing a balanced net ionic equation, you must balance the charges as well as the .

Precipitation Reactions The that forms after a reaction is called a precipitate. To decide if the product is a precipitate or not you must follow the rules.

Solubility Rules for Ionic Compounds

Soluble Insoluble (dissolve, dissociate, separate) Stays a solid in solution - (NO3 ) - Chlorate (ClO3 ) salts + + + metal salts (Na , K ) and salts (NH4 ) salts (Cl-) and Br-, I- Except ones with Ag, Hg, and Pb 2- salts (SO4 ) Except ones with Pb, Ag, Hg, Ba, Sr, and Ca Sulfides (S2-) 3- (PO4 ) 2- Chromates(CrO4 ) 2- Carbonates (CO3 ), Except ones with Na, K and Ca Hydroxides (OH-) *salt is used to mean

Acid/ Base Reactions/Neutralization Reactions

Acids – a substance that produces H+ ions () when it is dissolved in water. Strong acids completely dissociate in water. The strong acids are HCl, HNO3, and H2SO4.

Bases/ – a substance that produces ions (OH-) in water. Strong bases completely dissociate in water. The strong bases are NaOH and KOH.

In the reaction of a strong acid and a strong base, one product is always water and the other is always a salt (ionic compound) that remains dissolved in the water. Therefore, the net ionic equation for all strong acid/strong base reactions is always: + - H (aq) + OH (aq)  H2O(l)

Oxidation-Reduction () Equations – involves the transfer of .

Electrolyte – Carries a current through water. Strong – almost all of the dissociate into ions, carries a strong current. Weak – some may dissociate, but mostly the stays intact in water, carries very little current. Oxidation - Reduction

Combustion Synthesis Decomposition Double displacement Single displacement

Activity series Formation of metals of a gas Precipitation Acid/Base Activity Series of halogens