Chapter 3 Chemical Reactions - PDF Free Download

Chapter 3 Chemical Reactions Jeffrey Mack California State University, Sacramento

Chemical Reactions Reactants: Zn + I 2 Product: ZnI 2

Chemical Reactions Evidence of a chemical reaction: Gas Evolution Temperature Change Color Change Precipitation (insoluble species forms) In general, a reaction involves a rearrangement or change in oxidation state of atoms from reactants to products.

Chemical Equations Chemical Equations show: the reactants and products in a reaction. the relative amounts in a reaction. Example: 4 Al(s) + 3 O 2 (g) 2 Al 2 O 3 (s) The numbers in the front are called stoichiometric coefficients The letters (s), (g), (l) and (aq) are the physical states of compounds.

Reaction of Phosphorus with Cl 2 Notice the stoichiometric coefficients and the physical states of the reactants and products.

Reaction of Iron with Cl 2 Notice the stoichiometric coefficients and the physical states of the reactants and products.

4 Al(s) + 3 O 2 (g) 2 Al 2 O 3 (s) This equation states that: or... Chemical Equations 4 Al atoms + 3 O 2 molecules react to form 2 formula units of Al 2 O 3 4 moles of Al + 3 moles of O 2 react to form 2 moles of Al 2 O 3

Chemical Equations Law of the Conservation of Matter Because the same number of atoms are present in a reaction at the beginning and at the end, the amount of matter in a system does not change. 2HgO(s) 2 Hg(l) + O 2 (g)

Chemical Equations Since matter is conserved in a chemical reaction, chemical equations must be balanced for mass! In other words, there must be same number of atoms of the each kind on both sides of the equatoin. Lavoisier, 1788

Balancing Chemical Reactions Steps in balancing a chemical reaction using coefficients: 1. Write the equation using the formulas of the reactants and products. Include the physical states (s, l, g, aq etc ) 2. Balance the compound with the most elements in the formula first using integers as coefficients. 3. Balance elements on their own last. 4. Check to see that the sum of each individual elements are equal on each side of the equation. 5. If the coefficients can be simplified by dividing though with a whole number, do so.

Balancing Chemical Equations: Example balance last C 2 H 6 + O 2 CO 2 + H 2 O 2 C s & 6 H s 2 O s 1 C & 2 O s 2 H s & 1 O

Balancing Chemical Equations: Example balance last C 2 H 6 + O 2 CO 2 + H 2 O 2 C s & 6 H s 2 O s 1 C & 2 O s 2 H s & 1 O balance H first C 2 H 6 + O 2 CO 2 + H 2 O 3 This side will always have an even # of O-atoms This side has an odd # of O-atoms

Balancing Chemical Equations: Example balance last C 2 H 6 + O 2 CO 2 + H 2 O 2 C s & 6 H s 2 O s 1 C & 2 O s 2 H s & 1 O balance H first 2 C 2 H 6 + O 2 CO 2 + H 2 O 3

Balancing Chemical Equations: Example balance last C 2 H 6 + O 2 CO 2 + H 2 O 2 C s & 6 H s 2 O s 1 C & 2 O s 2 H s & 1 O balance H first 2 C 2 H 6 + O 2 CO 2 + H 2 O balance C next 2C 2 H 6 + O 2 CO 2 + 6H 2 O 4 3

Balancing Chemical Equations: Example balance last C 2 H 6 + O 2 CO 2 + H 2 O 2 C s & 6 H s 2 O s 1 C & 2 O s 2 H s & 1 O balance H first 2 C 2 H 6 + O 2 CO 2 + H 2 O balance C next 2C 2 H 6 + O 2 CO 2 + 6H 2 O balance O 7 2C 2 H 6 + O 2 4CO 2 + 6H 2 O 4 3

Balancing Equations Al(s) + Br 2 (l) Al 2 Br 6 (s)

Balancing Equations: Practice C 3 H 8 (g) + O 2 (g) CO 2 (g) + H 2 O(g) B 4 H 10 (g) + O 2 (g) B 2 O 3 (g) + H 2 O(g)

Balancing Equations: Practice Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water. Write the balanced chemical equation for this reaction.

Balancing Equations: Practice Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water. Write the balanced chemical equation for this reaction. _ Mg(OH) 2 (s) + _ HCl(aq) _ MgCl 2 (aq) + _ H 2 O(l)

Chemical Equations: Review What Scientific Principles are used in the process of balancing chemical equations? What symbols are used in chemical equations: gasses: liquids: solids: aqueous species in solution: What is the difference between P 4 and 4P in an eq.? In balancing a chemical equation, why are the reactant and product subscripts not changed?

Chemical Equilibrium When writing chemical reactions one starts with: Reactants products N 2 (g) + 3H 2 (g) 2NH 3 (g) Some reactions can also run in reverse: 2NH 3 (g) N 2 (g) + 3H 2 (g) Under these conditions, the reaction can be written: 3H 2(g) N 2(g) 2NH 3(g) Double arrows indicate Equilibrium.

Chemical Equilibrium Once equilibrium is achieved, reaction continues, but there is no net change in amounts of products or reactants.

Classifying Compounds Salts (ionic compounds): Composed of a metal and non metal element(s). Acids: Arrhenius definition Produce H + (aq) in water Examples: HCl, HNO 3, HC 2 H3O 2 Bases: Arrhenius definition Produce OH (aq) in water Examples: NaOH, Ba(OH) 2, NH 3

Classifying Compounds Molecular Compounds: Covalently bonded atoms, not acids, bases or salts. Compounds like alcohols (C 2 H 5 OH) or table sugar (C 6 H 12 O 6 ) These never break up into ions.

Classifying Compounds Classify the following as ionic, molecular, acid or base. Compound Na 2 SO 4 Ba(OH) 2 H 3 PO 4 CH 4 P 2 O 5 NH 3 HCN Type

Classifying Compounds Classify the following as ionic, molecular, acid or base. Compound Na 2 SO 4 Ba(OH) 2 H 3 PO 4 CH 4 P 2 O 5 NH 3 HCN Type ionic base acid molecular molecular base acid

Reactions in Aqueous Solutions Aqueous Solutions: Water as the solvent Solution = solute + solvent That which is dissolved (lesser amount) That which is dissolves (greater amount) There are three types of aqueous solutions: Those with Strong Electrolytes Those with Weak Electrolytes & those with non-electrolytes

Reactions in Aqueous Solutions Many reactions involve ionic compounds, especially reactions in water aqueous solutions. KMnO 4 in water K + (aq) + MnO 4- (aq)

Ionic Compounds (CuCl 2 ) in Water

Strong Electrolyte When ions are present in water, the solutions conduct electricity! Ions in solution are called ELECTROLYTES Examples of Strong Electrolytes: HCl (aq), CuCl 2 (aq) and NaCl (aq) are strong electrolytes. These dissociate completely (or nearly so) into ions. Strong Electrolytes conduct electricity well.

Strong Electrolytes HCl(aq), CuCl 2 (aq) and NaCl(aq) are strong electrolytes. These dissociate completely (or nearly so) into ions.

Weak Electrolytes Acetic acid ionizes only to a small extent, it is a weak electrolyte. CH3CO 2H(aq) CH3CO 2(aq) H (aq) Weak electrolytes exist in solution under equilibrium conditions. The small concentration of ions conducts electricity poorly. Weak electrolytes exit primarily in their molecular form in water.

Weak Electrolytes Weak electrolytic solutions are characterized by equilibrium conditions in solution: When acetic acid dissociates, it only partially ionizes. HC H O (aq) H (aq) + C H O (aq) + 2 3 2 2 3 2 95% 5% The majority species in solution is acetic acid in its molecular form. When writing a weak electrolyte in solution, one NEVER breaks it up into the corresponding ions! HC H O (aq) H (aq) + C H O (aq) + 2 3 2 2 3 2

Weak Electrolytes Acetic acid ionizes only to a small extent, so it is a weak electrolyte. CH 3 CO 2 H(aq) CH 3 CO 2- (aq) + H + (aq)

Non-Electrolytes Some compounds dissolve in water but do not conduct electricity. They are non-electrolytes. Examples include: sugar ethanol ethylene glycol Non-electrolytes do not dissociate into ions!

Species in Solution: Electrolytes Strong electrolytes: Characterized by ions only (cations & anions) in solution (water). Conduct electricity well Weak electrolytes: Characterized by ions (cations & anions) & molecules in solution. Conduct electricity poorly Non-electrolytes: Characterized by molecules in solution. Do not conduct electricity

Solutes in Aqueous Solutions

Solubility Rules How do we know if a compound will be soluble in water? For molecular compounds, the molecule must be polar. We will discuss polarity later, for now I will tell you whether or not a molecular compound is polar For ionic compounds, the compound solubility is governed by a set of SOLUBILITY RULES! You must learn the basic rules on your own!!!

Water Solubility of Ionic Compounds If one ion from the Soluble Compound list is present in a compound, then the compound is water soluble.

Types of Reactions in a Solution Precipitation Reactions: A reaction where an insoluble solid (precipitate) forms and drops out of the solution. Acid base Neutralization: A reaction in which an acid reacts with a base to yield water plus a salt. Gas forming Reactions: A reaction where an insoluble gas is formed. Reduction and Oxidation Reactions (RedOx): A reaction where electrons are transferred from one reactant to another.

EXCHANGE: Precipitation Reactions EXCHANGE Gas-Forming Reactions REACTIONS EXCHANGE Acid-Base Reactions REDOX REACTIONS

Chemical Reactions in Water EXCHANGE REACTIONS Pb(NO 3 ) 2 (aq) + 2 KI(aq) PbI 2 (s) + 2 KNO 3 (aq) The anions exchange places between cations. A precipitate forms if one of the products in insoluble.

Precipitation Reactions The driving force is the formation of an insoluble solid called a precipitate. Pb(NO 3 ) 2 (aq) + 2 KI(aq) 2 KNO 3 (aq) + PbI 2 (s) BaCl 2 (aq) + Na 2 SO 4 (aq) BaSO 4 (s) + 2 NaCl(aq) Precipitates are determined from the solubility rules.

Precipitation Reactions Which species is the precipitate? Pb(NO 3 ) 2 (?) + 2KI(?) 2KNO 3 (?) + PbI 2 (?)

Precipitation reactions Which species is the precipitate? Pb(NO 3 ) 2 (?) + 2KI(?) 2KNO 3 (?) + PbI 2 (?) From the solubility rules: All nitrate salts are soluble, therefore: Pb(NO 3 ) 2 (aq) + 2KI(?) 2KNO 3 (aq) + PbI 2 (?)

Precipitation Reactions Which species is the precipitate? Pb(NO 3 ) 2 (?) + 2KI(?) 2KNO 3 (?) + PbI 2 (?) From the solubility rules: All nitrate salts are soluble, therefore: Pb(NO 3 ) 2 (aq) + 2KI(?) 2KNO 3 (aq) + PbI 2 (?) All potassium salts are soluble, therefore: Pb(NO 3 ) 2 (aq) + 2KI(aq) 2KNO 3 (aq) + PbI 2 (?)

Net Ionic Equations Molecular Equation: all species listed as formula units or in molecular form. reactants products Note all states of each reactant or product by: (s), (l), (g) or (aq) Ionic Equation: All soluble (aq) species present are listed as ions. Leave all (s), (l) or (g) species as is. They do not dissociate into ions Net Ionic Equation: From the ionic equation, cancel out any species that appear on either side of the equation. These are known as the spectator ions and they are never part of a net ionic equation!

Writing Net Ionic Equations Molecular Equation: Pb(NO 3 ) 2 (aq) + 2KI(aq) 2KNO 3 (aq) + PbI 2 (s)

Writing Net Ionic Equations Molecular Equation: Pb(NO 3 ) 2 (aq) + 2KI(aq) 2KNO 3 (aq) + PbI 2 (s) Total Ionic Equation: Pb 2+ (aq) + 2NO 3 (aq) + 2K + (aq) + 2I (aq) 2K + (aq) + 2NO 3 (aq) + PbI 2 (s)

Writing Net Ionic Equations Molecular Equation: Pb(NO 3 ) 2 (aq) + 2KI(aq) Total Ionic Equation: Pb 2+ (aq) + 2NO 3 (aq) + 2K + (aq) + 2I (aq) 2KNO 3 (aq) + PbI 2 (s) Never break up any (s), (l) or (g) or molecular (aq) species! 2K + (aq) + 2NO 3 (aq) + PbI 2 (s)

Acids & Bases Arrhenius Definition: An acid is any substance that increases the H + (aq) concentration in an aqueous solution. HX(aq) H + (aq) + X (aq) A base is any substance that increases the OH (aq) concentration in an aqueous solution. MOH(aq) M + (aq) + OH (aq)

Acids and Bases Brönsted-Lowry: An acid is any substance that donates H + (aq) to another species in an aqueous solution. HX(aq) + H 2 O(l) H 3 O + (aq) + X (aq) H 3 O + (aq) = H + (aq) A base is any substance that accepts an H + (aq) in an aqueous solution. H + (aq) + NH 3 (aq) NH 4+ (aq)

Acids

Strong Acids Examples: Strong acids are almost completely ionized in water. (strong electrolytes) HX (aq) (X = Cl, Br & I) HNO 3 (aq) HClO 4 (aq) H 2 SO 4 (aq)* hydro ic acid nitric acid perchloric acid sulfuric acid * Only the 1 st H is strong, sulfuric acid dissociates via: H 2 SO 4 (aq) H + (aq) + HSO 4 (aq)

Acids An acid: H 3 O + in water

Weak Acids Examples: Weak Acids are incompletely ionized in water. (weak electrolytes) Weak acids are governed by dynamic equilibrium. HC 2 H 3 O 2 (aq) nitrous acid hydrosulfuric acid hydrogen sulfate ion acetic acid (vinegar) HNO 2 (aq) H 2 S (aq) HSO 4 (aq) Weak acids are always written in their molecular form. See you text and home work for more examples.

Strong Bases Bases: A base is a substance that produces OH (aq) ions in water by dissociation in water: H2O( ) NaOH(s) Na (aq) ΟΗ aq Strong bases are almost completely ionized in aqueous solution. (Strong electrolytes) Examples: Hydroxides of Group 1 (MOH(aq) where M = Li, Na, K ect ) and Ca, Sr, Ba.* *Ca(OH) 2, Sr(OH) 2 & Ba(OH) 2 are slightly soluble, but that which dissolves is present as ions only.

Bases Base: OH - in water NaOH(aq) Na + (aq) + OH - (aq) NaOH is a strong base

Weak Bases Weak Bases: NH 3 acts as a base by reacting with water: NH 3 (aq) + H 2 O(l) NH 4+ (aq) + OH (aq) Ammonia can also accept H + from an acid: NH 3 (aq) + H + (aq) NH 4+ (aq)

Ammonia, NH3

Reactions of Acids & Bases: Acid-Base Neutralization Acid + Base Salt + Water (usually) HA (aq) + MOH(aq) MA(aq) + HOH(l) Strong acid - Strong base neutralization: HBr(aq)/KOH(aq) Molecular Equation: HBr(aq) + KOH(aq) KBr (aq) + H 2 O(l) Total Ionic Equation: H + (aq) + Br (aq) + K + (aq) + OH (aq) Net Ionic equation: / / / / H + (aq) + OH (aq) K + (aq) + Br (aq) + H 2 O(l) H 2 O (l)

Acid-Base Reactions The driving force is the formation of water. NaOH(aq) + HCl(aq) NaCl(aq) + H 2 O(liq) Net ionic equation OH - (aq) + H 3 O + (aq) 2 H 2 O(l) This applies to ALL reactions of STRONG acids and bases.

Reactions of Acids & Bases: Acid-Base Neutralization Reactions of weak acids and strong bases: Molecular Equation: HC 2 H 3 O 2 (aq) + NaOH(aq) NaC 2 H 3 O 2 (aq) + H 2 O(l) Total Ionic Equation: HC 2 H 3 O 2 (aq) + Na + (aq) + OH (aq) Leave in molecular form / / Na + (aq) + C 2 H 3 O 2 (aq) + H 2 O(l) Net Ionic: HC 2 H 3 O 2 (aq) + OH (aq) C 2 H 3 O 2 (aq) + H 2 O(l)

Non-Metal Acids Nonmetal oxides can form acids in aqueous solutions: Examples: CO 2 (aq) + H 2 O(s) H 2 CO 3 (aq) SO 3 (aq) + H 2 O(s) H 2 SO 4 (aq) Both gases come from the burning of fossil fuels.

Bases Metal oxides form bases in aqueous solution CaO(s) + H 2 O(l) Ca(OH) 2 (aq) CaO in water. Indicator shows solution is basic.

Gas-Forming Reactions

Gas-Forming Reactions Metal carbonate salts react with acids to the corresponding metal salt, water and carbon dioxide gas. 2HCl(aq) + CaCO 3 (s) CaCl 2 (aq) + H 2 CO 3 (aq) decomposes Similarly: H 2 O(l) + CO 2 (g) HCl(aq) + NaHCO 3 (s) NaCl(aq) + H 2 O(l) + CO 2 (g) acid base salt water Neutralization!!!

Gas-Forming Reactions Group I metals: Na, K, Cs etc.. react vigorously with water 2K(s) + 2H 2 O(l) 2KOH(aq) + H 2 (g) Metals & acid: Some metals react vigorously with acid solutions: Zn(s) + 2H + (aq) Zn 2+ (aq) + H 2 (g)

Gas-Forming Reactions CaCO 3 (s) + H 2 SO 4 (aq) 2 CaSO 4 (s) + H 2 CO 3 (aq) Carbonic acid is unstable and forms CO 2 & H 2 O H 2 CO 3 (aq) CO 2 + water (The antacid tablet contains citric acid + NaHCO 3 )

Oxidation-Reduction Reactions Thermite reaction: Fe 2 O 3 (s) + 2Al(s) 2Fe(s) + Al 2 O 3 (s)

Oxidation-Reduction Reactions REDOX = reduction & oxidation O 2 (g) + 2 H 2 (g) 2 H 2 O(l)

Oxidation-Reduction Reactions Oxidation involves a reactant atom or compound losing electrons. Reduction involves a reactant atom or substance gaining electrons. Neither process can occur alone that is, there must be an exchange of electrons in the process. The substance that is oxidized is the reducing agent The substance that is reduced is the oxidizing agent oxidized reduced Mg(s) + 2H + (aq) Mg 2+ (aq) + H 2 (g) reducing agent oxidizing agent

Chemists use oxidation numbers to account for the transfer of electrons in a RedOx reaction. Oxidation numbers are the actual or apparent charge on atom when alone or combined in a compound. 1. The atoms of pure elements always have an oxidation number of zero. Examples: Oxidation Numbers Mg(s) Hg(l) I 2 (s) O 2 (g) All have an oxidation number of zero (0)

Oxidation Numbers 2. If an atom is charged, then the charge is the oxidation numbers. Examples: Ion Oxidation Number Mg 2+ (aq) Cl (aq) Sn 4+ (s) 2 Hg 2 (aq) +2 1 +4 +2/2 = +1 for each Hg atom

Oxidation Numbers 3. In a compound, fluorine always has an oxidation numbers of 1. 4. Oxygen most often has an oxidation number of 2.» *When combined with fluorine, oxygen has a positive O.N.» *In peroxide, the O.N. is 1. 5. In compounds, Cl, Br & I are 1 (Except with F and O present) 6. In compounds, H is +1, except as a hydride (H : 1)

Oxidation Numbers Examples: compound Oxidation Numbers HF(g) H = +1 F = 1 H 2 O(l) H = +1 O = 2 OF 2 (g) O = +2 F = 1 Na 2 O 2 (s) Na = +1 O = 1 HCl(g) H = +1 Cl = 1 NaH(l) Na = +1 H = 1

Oxidation Numbers Most common oxidation numbers:

Oxidation Numbers 7. For neutral compounds, the sum of the oxidation numbers equals zero. For a poly atomic ion, the sum equals the charge. Examples: +2 + 2 ( 1) = 0 MgCl 2 3 + 4 (+1) = +1 NH 4

Oxidation Numbers Determine the oxidation number of iron in the following compound:? + 3 ( 1) = 0 Fe(OH) 3 Iron must have an oxidation number of +3!

Recognizing a Redox Reaction In a RedOx reaction, the species oxidized and the species reduced are identified by the changes in oxidation numbers : Oxidation numbers: +1 0 + 2+ 2Ag (aq) + Cu(s) 2Ag(s) + Cu (aq) 0 +2 Oxidation numbers: Since silver goes from +1 to zero, it is reduced. Since copper goes from zero to +2, it is oxidized. The reaction is balanced for both mass and charge.

Practice: Identify the species that is Oxidized and Reduced by assigning oxidation numbers in the following reaction. 3CH (g) Cr O (aq) 8H (aq) Answer: 2 4 2 7 3CH OH(l) 2Cr (aq) 4H O(l) 3 3 2

Practice: Identify the species that is Oxidized and Reduced by assigning oxidation numbers in the following reaction. 3CH (g) Cr O (aq) 8H (aq) 2 4 2 7 3CH OH(l) 2Cr (aq) 4H O(l) 3 3 2 Answer: The carbon in methane (CH 4 ) is oxidized ( 4 to 2)

Redox Reactions

Oxidation-Reduction Reactions Iron gains 3 electrons (+3 to 0) oxidation number change. It is Reduced. Carbon loses 2 electrons (+2 to +4) it is Oxidized.

Redox Reactions REDOX = reduction & oxidation Corrosion of aluminum 2 Al(s) + 3 Cu 2+ (aq) 2 Al 3+ (aq) + 3 Cu(s)

Redox Reactions Cu(s) + 2 Ag + (aq) Cu 2+ (aq) + 2 Ag(s) In all reactions if a species is oxidized then another species must also been reduced

Redox Reactions Cu(s) + 2 Ag + (aq) Cu 2+ (aq) + 2 Ag(s)

Electron Transfer in a Redox Reaction 2Ag + (aq) + Cu(s) e e Cu 2+ (aq) + 2Ag(s) Two electrons leave copper. The silver ions accept them. The copper metal is oxidized to copper (II) ion. The silver ion is reduced to solid silver metal.

Redox Reactions in Our World Batteries Corrosion Fuels Manufacturing metals

Examples of Redox Reactions Metal + halogen 2 Al + 3 Br 2 Al 2 Br 6

Examples of Redox Reactions Nonmetal (P) + Oxygen P 4 O 10 Metal (Mg) + Oxygen MgO

Examples of Redox Reactions Metal + acid Mg + HCl Mg = reducing agent H + = oxidizing agent Metal + acid Cu + HNO 3 Cu = reducing agent HNO 3 = oxidizing agent

Reviewing What You ve Learned You have the following items available to you: Deionized water, ph paper, test tubes various metal nitrate salts, common acid and base solutions. Suggest a simple test or set of tests for identifying the unknown substances. Use proper terminology and write balanced chemical equations where applicable. Justify your answers thoroughly.

Reviewing What You ve Learned How would you determine whether or not a test tube containing a clear colorless solution is water or sulfuric acid? Given a white powder that my be silver chloride or sodium chloride. Whether a compound is silver nitrate or sodium nitrate.3