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Home , Aqueous solution, Ion

Aqueous

Types of Chemical Reactions

and is the dissolving medium, or . Stoichiometry

Figure 4.1: (Left) The water is Some polar. (Right) A space-filling model of the water molecule.  Water is “bent” or V-shaped.  The O-H bonds are covalent.  Water is a polar molecule.  Hydration occurs when dissolve in water.

Figure 4.3: (a) The molecule contains a polar O—H bond similar to those in the water molecule. (b) The polar water molecule interacts strongly with the A Solute polar O—H bond in ethanol. This is a case of "like dissolving like."  dissolves in water (or other “solvent”)

 changes phase (if different from the solvent)

 is present in lesser amount (if the same phase as the solvent)

1 A Solvent in Aqueous Solution Ionic Theory of Solutions  retains its phase (if different from the • Some molecular compounds dissolve but do solute) not dissociate into ions.

H O C H O (s) () →2 C H O (aq)  is present in greater amount (if the same 6 12 6 6 12 6 phase as the solute) – These compounds are referred to as nonelectrolytes. They dissolve in water to give a nonconducting solution.

Figure 4.2: Polar water interact with the positive and negative Ions in Aqueous Solution ions of a assisting in the dissolving Ionic Theory of Solutions process. • Many ionic compounds dissociate into independent ions when dissolved in water H 2 O NaCl(s)→Na+ (aq) + Cl− (aq) – These compounds that “freely” dissociate into independent ions in aqueous solution are called . – Their aqueous solutions are capable of conducting an .

Figure 4.5: When NaCl dissolves, the Na+ and Cl- ions are randomly Electrolytes dispersed in the water. Strong - conduct current efficiently

NaCl, HNO3 Weak - conduct only a small current vinegar, tap water

Non - no current flows pure water, solution

Ions in solution conduct . Strong electrolytes in solution conduct electricity!

2 Figure 4.4: Electrical Ions in Aqueous Solution conductivity Ionic Theory of Solutions of aqueous • Electrolytes are substances that dissolve in water to give an electrically conducting solutions. solution. Š Thus, in general, ionic that dissolve in water are electrolytes. Š Some molecular compounds, such as , also dissociate in aqueous solution and are considered electrolytes.. HCl (aq) º H+ (aq) + Cl- (aq)

Figure 4.6: HCl(aq) is Acids Strong acids - dissociate completely to completely produce H+ in solution ionized. hydrochloric and sulfuric

Weak acids - dissociate to a slight extent to give H+ in solution acetic and

Figure 4.8: Bases

(HC2H3O2) exists in water Strong bases - react completely with water to mostly as give OH− ions. undissociated molecules. Only a small percentage of the Weak bases - react only slightly with water molecules are to give OH− ions. ionized.

3 Figure 4.7: An aqueous Ions in Aqueous Solution Ionic Theory of Solutions solution of • Strong and weak electrolytes. sodium –A weak is an electrolyte that hydroxide. dissolves in water to give a relatively small percentage of ions.

– NH3 (g) + H2O (l) º NH4OH (aq) → + − NH4OH(aq) ← NH4 (aq) + OH (aq) – Most soluble molecular compounds are either nonelectrolytes or weak electrolytes.

Figure 4.9: The Ions in Aqueous Solution reaction Ionic Theory of Solutions: Summary • In summary, substances that dissolve in water of NH3 in are either electrolytes or nonelectrolytes. water. • Nonelectrolytes form nonconducting solutions because they dissolve as molecules.

• Electrolytes form conducting solutions because they dissolve as ions.

Ions in Aqueous Solution Working with Solutions Ionic Theory of Solutions: Summary Molar • Electrolytes can be strong or weak. • , or molarity (M), is defined as the moles of solute dissolved in • Almost all ionic substances that dissolve are strong electrolytes. one liter (cubic decimeter) of solution. moles of solute • Molecular substances that dissolve are either Molarity (M) = nonelectrolytes or weak electrolytes. liters of solution

4 Common Terms of Solution Molarity Concentration Molarity (M) = moles of solute per volume of solution in liters: Stock - routinely used solutions prepared in concentrated form.

moles of solute Concentrated - relatively large ratio of solute M ==molarity liters of solution to solvent. (5.0 M NaCl)

6 moles of HCl Dilute - relatively small ratio of solute to 3 M HCl = 2 liters of solution solvent. (0.01 M NaCl)

Figure 4.11: Working with Solutions (a) A measuring pipet is graduated and can be used to • The majority of chemical reactions measure various discussed here occur in aqueous solution. volumes of liquid accurately. – When you run reactions in liquid solutions, it is (b) a volumetric convenient to dispense the amounts of reactants (transfer) pipet by measuring out volumes of reactant solutions. is designed to measure one volume accurately.

Figure 4.10: Steps involved in the preparation of a standard Working with Solutions aqueous solution. Molar Concentration • When we dissolve a substance in a liquid, we call the substance the solute and the liquid the solvent. – The general term concentration refers to the quantity of solute in a standard quantity of solution.

5 Working with Solutions

Working with Solutions How many moles of Cl- are in 1.00 L of Molar Concentration 1.36 M FeCl3 solution? • Let’s try an example. - - – A sample of 0.0341 mol (III) , 1.36 moles FeCl3 x 3 moles Cl x 1.00 L soln = 4.08 moles Cl FeCl3, was dissolved in water to give 25.0 mL 1.00 L soln. 1 FeCl3 of solution. What is the molarity of the solution? moles of FeCl FeCl (s) ö Fe3+ (aq) + 3 Cl- (aq) –Since molarity = 3 3 liters of solution H2O 0.0341 mole of FeCl then M = 3 = 1.36 M FeCl 0.0250 liter of solution 3

Working with Solutions Working with Solutions From the equation moles of solute = Molarity Diluting Solutions Liters of solution •The molarity of a solution and its volume are inversely proportional. Therefore, Moles = M or, rearranging: L x M = moles adding water makes the solution less L concentrated. How many moles of NaCl are in 25.0 ml of 0.100 M – This inverse relationship takes the form of: NaCl solution? Moles = Moles Mi ×Vi = M f ×V f First, convert ml to L: 25.0 ml x 10-3 L/ml = 0.0250 ml – So, as water is added, increasing the final volume,

(0.0250 L soln) (0.100 moles NaCl/L soln) = 0.00250 moles Vf, the final molarity, Mf, decreases. NaCl

Figure 4.12: Procedure How many liters of 0.100 M HCl will be needed (a) About 100 mL of water is added to a 500 mL flask. to prepare 0.500 L of 0.0750 M HCl? (b) A measuring pipet is then used to transfer 375 of 0.100 M HCl solution to a volumetric flask. (c) Water is added to the flask to the calibration mark. Use MiVi = MfVf (d) The resulting solution is 0.0750 M HCl. We want to solve for Vi Mi = 0.100 M HCl Mf = 0.0750 M HCl V = 0.500 L HCl Always add f acids to water to prevent violent boiling! Solving the equation for Vi Vi = (MfVf)/ Mi = (0.0750 M x 0.500 L)/(0.100 M) = 0.375 L

6 Types of Solution Reactions Ions in Aqueous Solution  Precipitation reactions Molecular and Ionic Equations AgNO (aq) + NaCl(aq) → AgCl(s) + •A molecular equation is one in which the 3 reactants and products are written as if they NaNO3(aq) were molecules, even though they may  Acid- reactions actually exist in solution as ions.

NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) Ca(OH)2(aq) + Na2CO3(aq) → CaCO3 (s) + 2NaOH(aq)  Oxidation-reduction reactions – Note that Ca(OH)2, Na2CO3, and NaOH are all Fe2O3(s) + 2 Al(s) → 2 Fe(s) + Al2O3(s) soluble compounds but CaCO3 is not.

Simple Rules for Ions in Aqueous Solution − 1. Most (NO3 ) salts are soluble. Molecular and Ionic Equations + 2. Most alkali (group 1A) salts and NH4 are soluble. •An ionic equation, however, represents 3. Most Cl−, Br−, and I− salts are soluble (NOT Ag+, Pb2+, Hg 2+) 2 strong electrolytes as separate independent 4. Most salts are soluble (NOT BaSO , PbSO , HgSO , 4 4 4 ions. This is a more accurate representation CaSO4) 5. Most OH− salts are only slightly soluble (NaOH, KOH are of the way electrolytes behave in solution. soluble, Ba(OH) , Ca(OH) are marginally soluble) 2 2 2+ − + 2− 2− 2− 2− 3− Ca (aq) + 2OH (aq) + 2Na (aq) + CO (aq) → 6. Most S , CO3 , CrO4 , PO4 salts are only slightly soluble. 3 + − CaCO3 (s) ↓ +2Na (aq) + 2OH (aq)

Ions in Aqueous Solution Ions in Aqueous Solution Molecular and Ionic Equations Molecular and Ionic Equations • Complete and net ionic equations • Complete and net ionic equations. –A complete ionic equation is a chemical –A net ionic equation is a equation in which strong electrolytes (such as from which the spectator ions have been soluble ionic compounds) are written as removed. separate ions in solution. –A spectator is an ion in an ionic equation

Ca (NO 3 ) 2 (aq ) + K 2CO 3 (aq ) → CaCO 3 (s) + 2KNO 3 (aq ) that does not take part in the reaction. (strong) (strong) (insoluble) (strong) 2+ − + 2− Ca (aq ) + 2NO 3 (aq ) + 2K (aq ) + CO 3 (aq ) → 2+ − + 2− Ca (aq) + 2NO3 (aq) + 2K (aq) + CO3 (aq) → + − CaCO3 (s) + 2K (aq) + 2NO3 (aq) + − CaCO3 (s) + 2K (aq) + 2NO3 (aq)

7 Ions in Aqueous Solution Ions in Aqueous Solution Molecular and Ionic Equations Molecular and Ionic Equations • Complete and net ionic equations • Complete and net ionic equations – Separating the strong electrolytes into separate – Eliminating the spectator ions results in the net ions, we obtain the complete ionic equation. ionic equation. + − + − 2H (aq) + 2NO3 (aq) + Mg(OH)2 (s) → 2H (aq) + 2NO3 (aq) + Mg(OH)2(s) → 2+ − 2+ − 2H2O(l) + Mg (aq) + 2NO3 (aq) 2H2O(l) + Mg (aq) + 2NO3 (aq) + 2+ – Note that the nitrate ions did not participate in 2H (aq) + Mg(OH)2(s) → 2H2O(l) + Mg (aq) the reaction. These are spectator ions. This equation represents the “essential” reaction.

Figure 4.16: Precipitation of chloride by mixing solutions of and chloride. The

K+ and NO3- ions remain in solution.

Figure 4.17: The reaction of

KCl(aq) with AgNO3 to form Describing Reactions in Solution AgCl(s). 1. Molecular equation (reactants and products as compounds)

AgNO3(aq) + KCl(aq) → AgCl(s) + KNO3(aq)

2. Complete ionic equation (all strong electrolytes shown as ions) + − + − Ag (aq) + NO3 (aq) + K (aq) + Cl (aq) → + − AgCl(s) + K (aq) + NO3 (aq)

8 Describing Reactions in Solution Types of Chemical Reactions (continued) Acid-Base Reactions • The Arrhenius Concept 3. Net ionic equation (show only – The Arrhenius concept defines acids as components that actually react) substances that produce ions, H+, Ag+(aq) + Cl−(aq) → AgCl(s) when dissolved in water. – An example is , HNO , a molecular + − 3 K and NO3 are spectator ions. substance that dissolves in water to give H+ and - NO3 .

H2O + − HNO3 (aq) → H (aq) + NO3 (aq)

Types of Chemical Reactions Types of Chemical Reactions Acid-Base Reactions Acid-Base Reactions • The Arrhenius Concept • The Arrhenius Concept

– The Arrhenius concept defines bases as – The molecular substance ammonia, NH3, is a substances that produce hydroxide ions, OH-, base in the Arrhenius view, when dissolved in water. ← + − – An example is , NaOH, an NH3 (aq) + H2O(l)→NH4 (aq) + OH (aq) ionic substance that dissolves in water to give sodium ions and hydroxide ions. because it yields hydroxide ions when it reacts with water. H O + − NaOH(s) →2 Na (aq) + OH (aq)

Types of Chemical Reactions Types of Chemical Reactions Acid-Base Reactions Acid-Base Reactions • The Brønsted-Lowry Concept • The Brønsted-Lowry Concept – The Brønsted-Lowry concept of acids and – The Brønsted-Lowry concept defines an acid bases involves the transfer of a (H+) as the species (molecule or ion) that donates a from the acid to the base. proton (H+) to another species in a proton- – In this view, acid-base reactions are proton- transfer reaction. transfer reactions. –Abase is defined as the species (molecule or ion) that accepts the proton (H+) in a proton- transfer reaction.

9 Types of Chemical Reactions Types of Chemical Reactions Acid-Base Reactions Acid-Base Reactions • The Brønsted-Lowry Concept • The Brønsted-Lowry Concept + In the reaction of ammonia with water, The H (aq) ion associates itself with water to form + ← + − H3O (aq). NH3 (aq) + H2O(l) → NH4 (aq) + OH (aq) + + H (aq) + H2O(l) → H3O (aq) H+ This “mode of transportation” for the H+ ion is the H2O molecule is the acid because it donates a called the ion. proton. The NH3 molecule is a base, because it accepts a proton.

Types of Chemical Reactions Types of Chemical Reactions Acid-Base Reactions Acid-Base Reactions • The Brønsted-Lowry Concept • In summary, the Arrhenius concept and the Brønsted-Lowry concept are essentially the The dissolution of nitric acid, HNO3, in water is therefore a proton-transfer reaction same in aqueous solution. − + – The Arrhenius concept HNO3 (aq) + H2O(l) → NO3 (aq) + H3O (aq) acid: proton (H+) donor

- H+ base: hydroxide ion (OH ) donor

where HNO3 is an acid (proton donor) and H2O is a base (proton acceptor).

Types of Chemical Reactions Types of Chemical Reactions Acid-Base Reactions Acid-Base Reactions • In summary, the Arrhenius concept and the • Strong and Weak Acids and Bases Brønsted-Lowry concept are essentially the –A strong acid is an acid that ionizes completely same in aqueous solution. in water; it is a strong electrolyte. – The Brønsted-Lowry concept − + HNO3 (aq) + H2O(l) → NO3 (aq) + H3O (aq) acid: proton (H+) donor − + base: proton (H+) acceptor HCl(aq) + H2O(l) → Cl (aq) + H3O (aq)

– Table 4.3 lists the common strong acids.

10 Types of Chemical Reactions Acid-Base Reactions • Strong and Weak Acids and Bases –A strong base is a base that is present entirely as ions, one of which is OH-; it is a strong electrolyte. H O NaOH(s) →2 Na+ (aq) + OH− (aq) – The of Group IA and IIA elements, except for hydroxide, are strong bases. (see Table 4.3) Return to Slide 43

Types of Chemical Reactions Types of Chemical Reactions Acid-Base Reactions Acid-Base Reactions • Strong and Weak Acids and Bases • Strong and Weak Acids and Bases – You will find it important to be able to identify –A weak acid is an acid that only partially an acid or base as strong or weak. ionizes in water; it is a weak electrolyte. – When you write an ionic equation, strong acids – The hydrogen molecule, HCN, reacts and bases are represented as separate ions. with water to produce a small percentage of – Weak acids and bases are represented as ions in solution. undissociated “molecules” in ionic equations. ← − + HCN(aq) + H2O(l)→CN (aq) + H3O (aq)

Types of Chemical Reactions Types of Chemical Reactions Acid-Base Reactions Acid-Base Reactions • Strong and Weak Acids and Bases • Neutralization Reactions – One of the chemical properties of acids and bases –A weak base is a base that is only partially is that they neutralize one another. ionized in water; it is a weak electrolyte. –A neutralization reaction is a reaction of an acid – Ammonia, NH , is an example. 3 and a base that results in an and water. NH (aq) + H O(l) →← NH + (aq) + OH− (aq) 3 2 4 – The ionic compound that is the product of a neutralization reaction is called a salt.

HCN(aq) + KOH(aq) → KCN(aq) + H2O(l) acid base salt

11 Figure 4.18a: The Key Titration Terms Titrant - solution of known concentration of an acid used in titration with a base. Analyte - substance being analyzed Equivalence point - enough titrant added to react exactly with the analyte Endpoint - the indicator changes color so you can tell the equivalence point has been reached.

Figure 4.18b: Figure 4.18c: The titration The titration of an acid of an acid with a base with a base.

Quantitative Analysis Quantitative Analysis Volumetric Analysis Volumetric Analysis – After balancing the equation, we must convert the 0.0350 L (35.0 mL) to moles of H SO (using the • Consider the reaction of , 2 4 molarity of the H2SO4). H SO , with sodium hydroxide, NaOH: 2 4 • Then, convert to moles of NaOH (from the

H2SO4 (aq) + 2NaOH(aq) → 2H2O(l) + Na2SO4 (aq) balanced chemical equation). • Finally, convert to volume of NaOH solution – Suppose a beaker contains 35.0 mL of 0.175 M (using the molarity of NaOH).

H2SO4. How many milliliters of 0.250 M NaOH 0.175 mole H2SO4 2 mol NaOH 1 L NaOH soln. must be added to completely react with the (0.0350L)× × × = 1 L H2SO4 solution 1 mol H SO 0.250 mol NaOH sulfuric acid? 2 4 0.0490 L NaOH solution ( or 49.0 mL of NaOH solution)

12 Performing Calculations for Types of Chemical Reactions Acid-Base Reactions Acid-Base Reactions 1. List initial species and predict reaction. • Neutralization Reactions 2. Write balanced net ionic reaction. – The net ionic equation for each acid-base 3. Calculate moles of reactants. neutralization reaction involves a transfer of a proton. 4. Determine limiting reactant. – Consider the reaction of the strong acid , 5. Calculate moles of required reactant/product. HCl(aq) and a strong base, LiOH(aq). 6. Convert to grams or volume, as required. HCl(aq) + KOH(aq) → KCl(aq) + H2O(l)

Types of Chemical Reactions Types of Chemical Reactions Acid-Base Reactions Acid-Base Reactions • Neutralization Reactions • Neutralization Reactions – Writing the strong electrolytes in the form of – Canceling the spectator ions results in the net ions gives the complete ionic equation. ionic equation. Note the proton transfer. H + (aq ) + Cl − (aq ) + K + (aq) + OH − (aq ) → + − + − + − H (aq) + Cl (aq) + K (aq) + OH (aq) → K (aq) + Cl (aq) + H 2O(l) + − + − H (aq) + OH (aq) → H2O(l) K (aq) + Cl (aq) + H2O(l)

H+

Types of Chemical Reactions Types of Chemical Reactions Acid-Base Reactions Acid-Base Reactions • Neutralization Reactions • Acid-Base Reactions with Formation – In a reaction involving HCN(aq), a weak acid, – react with acids to form H2S, hydrogen and KOH(aq), a strong base, the product is gas. KCN, a strong electrolyte. – The net ionic equation for this reaction is Na2S + 2HCl → 2NaCl + H 2S ↑

− − – These reactions are summarized in Table 4.4. HCN(aq) + OH (aq) → CN (aq) + H2O(l)

Note the proton transfer. H+

13 Types of Chemical Reactions Acid/Base Reactions Acid-Base Reactions • Acid-Base Reactions with Gas Formation

– react with acids to form CO2, dioxide gas.

Na2CO3 + 2HCl → 2NaCl + H2O + CO2 ↑

– react with acids to form SO2, sulfur dioxide gas.

Na2SO3 + 2HCl → 2NaCl + H2O + SO2 ↑ Return to Slide 52

Figure 4.19: The reaction of solid Types of Chemical Reactions sodium and gaseous to form solid . • Oxidation-Reduction Reactions – Oxidation-reduction reactions involve the transfer of from one species to another. – Oxidation is defined as the loss of electrons. – Reduction is defined as the gain of electrons. – Oxidation and reduction always occur simultaneously.

Figure 4.13: Combination reaction. Figure 4.20: A summary of an oxidation- reduction process, in which M is oxidized and X is reduced.

Return to Slide 63

14 Figure 4.10: Reaction of iron with Types of Chemical Reactions Cu2+ (aq). Photo courtesy of • Oxidation-Reduction Reactions American – The reaction of an iron nail with a solution of Color. (II) sulfate, CuSO4, is an oxidation- reduction reaction (see Figure 4.10). – The molecular equation for this reaction is:

Fe(s) + CuSO4 (aq) → FeSO4 (aq) + Cu(s)

Return to Slide 54

Types of Chemical Reactions Types of Chemical Reactions Oxidation-Reduction Reactions • Oxidation-Reduction Reactions • Oxidation Numbers – The net ionic equation shows the reaction of – The concept of oxidation numbers is a simple 2+ iron with Cu (aq) to produce iron(II) ion way of keeping track of electrons in a reaction. and copper metal. – The oxidation number (or ) of

Loss of 2 e-1 oxidation an in a substance is the actual charge of the atom if it exists as a monatomic ion. 2+ 2+ Fe(s) + Cu (aq) → Fe (aq) + Cu(s) – Alternatively, it is hypothetical charge

Gain of 2 e-1 reduction assigned to the atom in the substance by simple rules.

Types of Chemical Reactions Types of Chemical Reactions Oxidation-Reduction Reactions Oxidation-Reduction Reactions • Oxidation Number Rules • Oxidation Number Rules Rule Applies to Statement Rule Applies to Statement 1 Elements The oxidation number of an atom in an 4 Hydrogen The oxidation number of hydrogen is +1 in element is zero. This includes molecular elements. its covalent compounds. 2 Monatomic The oxidation number of an atom in a 5 Halogens Fluorine is –1 in all its compounds. The other ions monatomic ion equals the charge of the ion. halogens are –1 unless the other element is 3 The oxidation number of oxygen is –2 in another halogen or oxygen. most of its compounds. (An exception is O in 6 Compounds The sum of the oxidation numbers of the

H2O2 and other peroxides, where the and ions in a compound is zero. The sum in a oxidation number is –1.) equals the charge on the ion.

15 Types of Chemical Reactions Types of Chemical Reactions Oxidation-Reduction Reactions Oxidation-Reduction Reactions • Describing Oxidation-Reduction Reactions • Describing Oxidation-Reduction Reactions – Look again at the reaction of iron with –A half-reaction is one of the two parts of an copper(II) sulfate. oxidation-reduction reaction. One involves the loss of electrons (oxidation) and the other Fe(s) + Cu2+ (aq) → Fe2+ (aq) + Cu(s) involves the gain of electrons (reduction). 2+ − – We can write this reaction in terms of two half- Fe(s) → Fe (aq) + 2e oxidation half-reaction reactions. 2+ − Cu (aq) + 2e → Cu(s) reduction half-reaction

Types of Chemical Reactions Types of Chemical Reactions Oxidation-Reduction Reactions Oxidation-Reduction Reactions • Describing Oxidation-Reduction Reactions • Balancing Simple Oxidation-Reduction Reactions –An oxidizing agent is a species that oxidizes another species; it is itself reduced. – At first glance, the equation representing the –A reducing agent is a species that reduces reaction of metal with silver(I) ions might another species; it is itself oxidized. appear to be balanced. Loss of 2 e- oxidation Zn(s) + Ag+ (aq) → Zn2+ (aq) + Ag(s) reducing agent Fe(s) + Cu2+ (aq) → Fe2+ (aq) + Cu(s) – However, a balanced equation must have a oxidizing agent charge balance as well as a balance. Gain of 2 e- reduction

Types of Chemical Reactions Types of Chemical Reactions Oxidation-Reduction Reactions Oxidation-Reduction Reactions • Balancing Simple Oxidation-Reduction Reactions • Balancing Simple Oxidation-Reduction Reactions

– Since the number of electrons lost in the – Adding the two half-reactions together, the oxidation half-reaction must equal the number electrons cancel, gained in the reduction half-reaction, 2+ − Zn(s) → Zn (aq) + 2e oxidation half-reaction 2+ − oxidation half-reaction + − Zn(s) → Zn (aq) + 2e ______2Ag (aq) + 2e → 2Ag(s) reduction half-reaction + − 2 Ag ( aq ) + 2 e → 2 Ag(s) reduction half-reaction Zn (s) + 2 Ag2+ (aq) º Zn2+ (aq) + 2 Ag (s) we must double the reaction involving the reduction of the silver.

16 Rules for Assigning Oxidation States

Figure 4.12: The burning of 1. Oxidation state of an atom in an element = 0 metal in chlorine. 2. Oxidation state of monatomic element = charge Photo courtesy of James Scherer. 3. Oxygen = −2 in covalent compounds (except in peroxides where it = −1) Ca (s) + Cl2 (g) º CaCl2 (s) Ca º Ca2+ + 2 e- 4. H = +1 in covalent compounds 5. Fluorine = −1 in compounds Cl + 2 e- º 2 Cl- 2 6. Sum of oxidation states = 0 in compounds

Ca (s) + Cl2 (g) º CaCl2 (s) Sum of oxidation states = charge of the ion

Return to Slide 62

Balancing by Half-Reaction Half-Reaction Method - Method Balancing in Base 1. Balance as in acid. 1. Write separate reduction, oxidation reactions. 2. Add OH− that equals H+ ions (both 2. For each half-reaction: sides!)  Balance elements (except H, O) 3. Form water by combining H+, OH−.  Balance O using H2O  Balance H using H+ 4. Check elements and charges for balance.  Balance charge using electrons

Balancing by Half-Reaction Stoichiometry Steps for reactions Method (continued) in solution. 3. If necessary, multiply by integer to equalize count.

4. Add half-reactions.

5. Check that elements and charges are balanced.

17