Chemical Reactions

• A chemical reaction is a chemical change where one or more substances are destroyed and new substances are created.

Parts of a Chemical Reaction

• Reactants are substances that are destroyed during the chemical change; bonds break.
• Products are substances created by the chemical change; new bonds form.
• Reactants yield products; the arrow (→) represents "yields".

Chemical Reaction Symbols

• (s) = solid
• (l) = liquid
• (g) = gas
• (aq) = aqueous solution, the substance is dissolved in H₂O
• "+" separates two or more reactants or products
• "→" yield sign separates reactants from products

Evidence of a Chemical Reaction

• Evolution of light or heat
• Temperature change to the surroundings, increase or decrease
• Formation of a gas, bubbling or an odor, not including boiling
• Color change due to the formation of a new substance
• Formation of a precipitate; a new solid forms from two aqueous solutions.

Word Equations

• Statements indicate the reactants and products in a chemical reaction, for example: Iron (s) + chlorine (g) → iron (III) chloride (s).
• This is read as "Solid iron and chlorine gas react (combine) to produce solid iron (III) chloride".

Skeleton Equations

• Skeleton equations use chemical formulas rather than words to identify reactants and products of a chemical reaction.
• Example equation is: 2Fe(s) + 3 Cl2(g) → 2FeCl3 (s).
• Skeleton equations are not yet "balanced" by coefficients.
• The numbers preceding the chemical formulae are coefficients, used to balance the reaction.
• Subscripts are the numbers within the chemical formulae.
• The balanced reaction 6 Na (s) + Fe2O3 (s) → 3 Na₂O (s) + 2 Fe (s) can be read as either: -"6 atoms of solid sodium reacts with 1 formula unit of solid iron (III) oxide yields 3 formula units of solid sodium oxide and 2 atoms of solid iron".
• "6 moles of solid sodium reacts with 1 mole of solid iron (III) oxide yields 3 moles of solid sodium oxide plus 2 moles of solid iron".
• Chemical reactions can never be read in terms of grams, only in terms of particles or groups of particles (moles).

Conservation of Mass

• Antoine Lavoisier discovered that matter cannot be created nor destroyed; this is the law of conservation of mass.
• In the reaction H2 (g) + O2 (g) → H2O (1), the equation isn't balanced; one oxygen atom appears to be missing.
• The balanced equation, according to conservation of mass, should be 2H2 (g) + 1O2 (g) → 2H2O (1).
• Coefficients represent the number of molecules or moles of each reactant or product.

Properties Conserved During Reactions

• Conserved properties include mass, types of atoms, and number of each atom.
• Non-conserved properties include color, physical state (solid, liquid, gas), volume, and number of moles of reactants/products.

Types of Chemical Reactions

• There are five basic types: synthesis, decomposition, single replacement, double replacement, and combustion.

Synthesis Reaction

• (A + B → AB) two or more simple substances (the reactants) combine to form a more complex substance (the product).
• Example equation is 2Mg + O2 → 2MgO.
• Types of synthesis equations include:
• Element A + Element B → Compound, example equation: 2Na(s) + Cl2 (g) → 2NaCl(s)
• Element + Compound A → Compound B, example equation: O2(g) + 2SO2(g) → 2SO3(g)
• Compound A + Compound B → Compound C, example equation: CaO(s) + H₂O(I) → Ca(OH)2 (s)
• Metallic and nonmetallic elements react to form ionic compounds.
• The resultant compound should be charge balanced by the criss-cross method, for example 4Li + O2 → 2Li₂O.
• Nonmetals react with each other to form covalent (molecular) compounds, able to draw a valid Lewis Structure for the product.
• For example, 2H2 + O2 → 2H₂O or H2 + O2 → H2O2.

Decomposition Reactions

• (AB → A + B) a more complex substance (the reactant) breaks down into two or more simple parts (products).
• Synthesis and decomposition reactions are opposites.
• Example equation is 2H₂O → 2H2 + O2.
• Decomposition of a compound produces two or more elements and/or compounds.
• Products are always simpler than the reactant.
• Gases are often produced (H2, N2, O2, CO2, etc.) in the decomposition of covalent compounds.
• Ionic compounds may be decomposed into pure elements by using electricity (electrolysis).
• Electrolysis is how pure metals are obtained from salts.

Single Replacement Reactions

• (A + BC → B + AC) a single uncombined element replaces another element in an ionic compound based on its reactivity.
• Example equations include: K + Na2SO4 → K2SO4 + Na and Na + K2SO4 → No Reaction.
• Examine the reaction Zn + CuSO4 → ZnSO4 + Cu; this reaction occurs.
• Examine the reaction Cu + ZnSO₄ → No Reaction.
• Zinc will replace Cu in solution, but not vice versa.
• Elements with higher reactivities replace elements with lower reactivities during a single-replacement reaction, but not vice-versa.

Double Replacement Reactions

• (AB + CD → AD + CB) parts of two aqueous ionic compounds switch places to form two new compounds.
• Reactions have two reactants and two products.
• Example equation: AgNO3 + NaCl → AgCl + NaNO3.
• Not all double replacement reactions occur.
• The reaction occurs if a solid precipitate is produced, a gas is produced, or water is produced.
• If none of the above are produced and both products are (aq), then there is no reaction (NR).
• Examples of equations:
• Pb(NO3)2 (aq) + 2Nal (aq) → Pbl₂ (s) + 2NaNO3 (aq) (precipitate forming)
• HCI (aq) + NaOH (aq) → NaCl (aq) + H2O (I) (water-forming, acid-base, neutralization)
• CaCO3 (s) + 2HCl (aq) → CaCl2 (aq) + H2CO3 (gas-forming)
• Use the solubility rules to determine if one of the products of a double replacement reaction will be a precipitate.
• Soluble compounds break down when put in water, in water NaCl → Na¹+ and Cl1-; NaCl has dissolved, is soluble, and forms an aqueous solution (aq).
• Insoluble compounds do NOT break down when put in water: CaCO3 does NOT break down into Ca²+ and CO32- ions, CaCO3 stays as a solid, (s) or (ppt).

Combustion Reactions

• Involves oxygen (O2) as a reactant, combining with another substance.
• Combustion reactions are exothermic and involve the complete combustion of a hydrocarbon which produces CO2 and H₂O.
• Incomplete combustion of a hydrocarbon will produce CO and possibly C (black carbon soot).
• Example equations
• CH4 + 2O2 => CO2 + 2H2O (complete combustion – blue flame)
• CH4 + 1.50₂ => CO + 2H2O (incomplete combustion – yellow flame)
• CH4 + O2 => C + 2H2O soot) (incomplete combustion – yellow flame
• Any synthesis reaction which involves O2 as a reactant is also considered to be a combustion reaction; for example, 2Mg + O2 → 2MgO (metal oxide).
• Called the combustion of magnesium or the synthesis of magnesium oxide, always produces a metal oxide using the criss-cross method.

Classifying Equations

• C4H8 + 6O2 → 4CO2 + 4H2O is combustion of a hydrocarbon.
• HCI + NaOH → H2O + NaCl is double replacement, water forming.
• 2KNO3(s) → 2KNO2(s) + O2(g) is decomposition.
• 2Ag + S → Ag₂S is synthesis.
• MgCO3(s) → MgO(s) + CO2(g) is decomposition.
• Cl2 + 2KBr → 2KCl + Br2 is single replacement.

Counting Atoms Rules

• Coefficients propagate to the right through the entire compound, whether or not parentheses are present.
• Subscripts affect only the element to the left of the subscript, unless a subscript occurs to the right of a parentheses which propagates to the left through the parentheses.
• When a coefficient and subscript “meet”, you must multiply the two.

Balancing Chemical Reactions Rules

• Change the coefficients, but never the subscripts!
• Coefficients must reduced to represent the lowest possible numbers.
• It is OK to use fraction coefficients, but you must get rid of them in the end (multiply through by denominator).
• Save certain elements until the end (do other elements first); these include H, C, O.
• Do a final balance check for each element!