Chemical Reactions
6.1 Evidence for a Chemical Reaction
- Chemical reactions: often give visual signal. represented by writing a chemical equation.
6.2 Chemical Equations
- Chemical equation: the chemicals present before the reactant(the reactants) are shown to the left of an arrow and the chemicals formed by the reactant (the product) are shown to the right of an arrow.
- in a chemical reaction, atoms are neither created nor destroyed
6.3 balancing Chemical Equations
- The identities (formulas) of the compounds must never be changed in balancing a chemical equation.
- Coefficients: integers that help balance the equation.
ex. 1
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Unbalanced equation:
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Reactions in Aqueous Solutions
7.1 Predicting Whether a Reaction will occur
- 4 common “driving forces” for reaction
- Formation of a solid(ex precipitation)
- Formation of water(ex. combustion)
- Transfer of electron( ex. oxidation- reduction)
- Formation of gas(ex. combustion)
7.2 Reactions in which a Solid Forms
- Precipitation: one driving force for a chemical reaction is the formation of a solid, a process called precipitation. The solid that forms is called a precipitate, and the reaction is known as a pre citation reaction.
- Strong electrolyte: when each unit of a substance that dissolves in water produces separated ions, the substance is called a strong electrolyte.
- 4 common “driving forces” for reaction
- Formation of a solid(ex precipitation)
- Formation of water(ex. combustion)
- Transfer of electron( ex. oxidation- reduction)
- Formation of gas(ex. combustion)
Solubility Rules:
- 1. Most nitrate( NO3 -1) are soluble
- 2. Most salts of Na+,K+ and NH4+ are soluble
- 3. Most chloride salts are soluble.
- exception: Ag+,Pb2+ , Hg2 +2
- 4. Most sulfates salts are soluble
- exception:Ba2+, Pb2+ , Ca2+
- 5. Most hydroxide (OH-)compounds are insoluble
- Exception- ions in rule number 2 and Ba/Ca
- 6.Most sulfides, carbonate, phosphate are insoluble
- Exception- ions in rule number 2
- 1. Most nitrate( NO3 -1) are soluble
- 2. Most salts of Na+,K+ and NH4+ are soluble
- 3. Most chloride salts are soluble.
- exception: Ag+,Pb2+ , Hg2 +2
- 4. Most sulfates salts are soluble
- exception:Ba2+, Pb2+ , Ca2+
- 5. Most hydroxide (OH-)compounds are insoluble
- Exception- ions in rule number 2 and Ba/Ca
- 6.Most sulfides, carbonate, phosphate are insoluble
- Exception- ions in rule number 2
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- There are 3 ways to write chemical equations
- 1. Molecular equation
- No charges are showing. Only compounds
- AgNO3 (aq)+ NaCl(aq) -> AgCl(s) + NaNO3(aq)
- II. Complete ionic equations
- All (aq) compounds will disassociate “break” into their ions
- Ag+(aq) +NO3- (aq)+ Na+(aq) +Cl-(aq) -> AgCl(s) + Na+ (aq) +NO3-(aq)
- III. Net ionic equation
- Find the solid and then write the ions in the reactants that makes up the product
- Ag+(aq) + Cl-(aq)-> AgCl(s)
- There are 3 ways to write chemical equations
- 1. Molecular equation
- No charges are showing. Only compounds
- AgNO3 (aq)+ NaCl(aq) -> AgCl(s) + NaNO3(aq)
- II. Complete ionic equations
- All (aq) compounds will disassociate “break” into their ions
- Ag+(aq) +NO3- (aq)+ Na+(aq) +Cl-(aq) -> AgCl(s) + Na+ (aq) +NO3-(aq)
- III. Net ionic equation
- Find the solid and then write the ions in the reactants that makes up the product
- Ag+(aq) + Cl-(aq)-> AgCl(s)
Types of Reactions:
- 1. Combustion.
- -CH4+O2 -> CO2 + H2O
- -Always look for in the reactant a carbon source and oxygen gas and in the product CO2
- -CH4+O2 -> CO2 + H2O
- 2. Synthesis
- -Adding elements and/or compounds to form a new single compound
- -Example: H2 +O2-> H2O
- 3. Decomposition
- -Breaking a compound into simpler parts
- -Example: H2O-> H2 +O2
- 4. Single displacement
- -One element/compound replaces another one.
- note only a cation can replace a cation and an anion can only replace an anion
- -Example: Zn +FeCl2-> Fe + ZnCl2
- 5. Double displacement
- -Elements/compounds switches partner to form 2 new compound.
- -Example: AgNO3 (aq)+ NaCl(aq) -> AgCl(s) + NaNO3(aq)
- 6. Acid/ base reaction
- -In the reactant, there must be both an acid(H) and a base(OH).
- -Example: HNO3 + NaOH->H2O + NaNO3
- 7. Precipitation reaction:
- -Formation of a solid when two or more aqueous solution are combined
- -Example: AgNO3 (aq)+ NaCl(aq) -> AgCl(s) + NaNO3(aq)
- 1. Combustion.
- -CH4+O2 -> CO2 + H2O
- -Always look for in the reactant a carbon source and oxygen gas and in the product CO2
- -CH4+O2 -> CO2 + H2O
- 2. Synthesis
- -Adding elements and/or compounds to form a new single compound
- -Example: H2 +O2-> H2O
- 3. Decomposition
- -Breaking a compound into simpler parts
- -Example: H2O-> H2 +O2
- 4. Single displacement
- -One element/compound replaces another one.
- note only a cation can replace a cation and an anion can only replace an anion
- -Example: Zn +FeCl2-> Fe + ZnCl2
- 5. Double displacement
- -Elements/compounds switches partner to form 2 new compound.
- -Example: AgNO3 (aq)+ NaCl(aq) -> AgCl(s) + NaNO3(aq)
- 6. Acid/ base reaction
- -In the reactant, there must be both an acid(H) and a base(OH).
- -Example: HNO3 + NaOH->H2O + NaNO3
- 7. Precipitation reaction:
- -Formation of a solid when two or more aqueous solution are combined
- -Example: AgNO3 (aq)+ NaCl(aq) -> AgCl(s) + NaNO3(aq)