Reactions in aqueous solutions Aqueous solution Solution in

Reactions in aqueous solutions • Aqueous solution: – Solution in which water is the solvent (dissolving agent). • 3 major types of chemical processes of aqueous solutions: – Precipitation reactions – Acid-base reactions – Redox reactions

• Solution: – Homogenous mixture of 2 or more substances. • Solvent: – Dissolving medium, usually present in greater quantity. • Solute: – The other substance(s) in the solution.

• Electrolyte: – A substance whose aqueous solution forms ions; conducts electricity. – Ionic compounds. • Nonelectrolyte: – Substance that does not form ions in an aqueous solution; poor conductor. – Molecular compounds.

• Ionic compounds in water: – Dissociate into its component ions.

Do not get to wrapped up in the difference between the terms ionization and dissociation. Consider them to mean the same thing, the separation of a substances ions. Equations showing ionization or dissociation

Molecular compounds in H 2 O Molecular compounds – nonmetal + nonmetal Structural integrity of molecule is usually maintained meaning no ions form (C 12 H 22 O 11) Exception: Some molecular solutes interact with water to form ions. These would be electrolytes. Examples: Acids HCl, H 2 C 3 O 2 Ammonia NH 3

Strong Electrolytes • Exists in solution completely or almost completely as ions. • All ionic compounds and a few molecular compounds. (Ex: Strong Acids)

Weak Electrolytes • Molecular compounds that produce a small concentration of ions when dissolved in H 2 O. Ex: Acetic acid (HC 2 H 3 O 2) only slightly ionizes when dissolved in water. HC 2 H 3 O 2(aq) H+(aq)+ C 2 H 3 O 2 -(aq) Weak acids are better conductors if they are dilute, as you will see in lab. Explain.

• Reactions that result in an insoluble product. • Insoluble: – Substance with solubility less than 0. 01 mol/L – Water molecules cannot overcome the attraction between the ions.

KI (aq) + Pb(NO 3)2 (aq) Pb. I 2(s) + KNO 3 (aq)

You must be able to determine whether a substance is soluble in water by simple examination of the chemical formula. To do so, you must memorize how specific polyatomic ions act in water. Not as hard as it sounds. We will focus mainly on 10 anions. This will give you the tools to predict the solubility of many compounds.

Solubility of Ionic Compounds • All acetates and nitrates are soluble in water. • All ionic compounds of alkali metals and ammonium are soluble. – (1 A goes AWAY) • Solubility rules are on your reference sheet.

Soluble Exceptions Cl- Ag+, Hg 22+, Pb 2+ Acetate ion C 2 H 3 O 2 - None Br- Ag+, Hg 22+, Pb 2+ I- Ag+, Hg 22+, Pb 2+ NO 3 - None SO 42 - Sr 2+, Ba 2+, Hg 22+, Pb 2+

Insoluble Exceptions CO 32 OHPO 43 S 2 - NH 4+ and Group 1 metals and Ba 2+, Sr 2+, Ca 2+

Equation Types • Molecular • Complete Ionic 3 • Net ionic equation

Ionic Equations • Those ions that appear on both sides of a complete ionic equation are known as Spectator Ions. • Net ionic equations do not spectator ions. include

• Exchange Reactions • Metathesis reactions • Double displacement • Double replacement

Writing Net Ionic Equations 1) Write a balanced molecular equation. 2) Rewrite the equation showing ions of strong electrolytes only. 3) Identify and cancel all spectator ions.

Acid-Base Reactions Acids: • Ionize in H 2 O, causes increase in H+ ions. • H+ ions are bare protons. • Acids are proton donors. All this acid rain is killing my complexion!

• Monoprotic Acids: (HCl, HNO 3) – Acids that can only yield one H+ per molecule upon ionization. HCl H+ + Cl-

Diprotic Acids: (H 2 SO 4) Ionization occurs in 2 steps. Only the first ionization is complete.

Is HF a weak or strong acid? weak acid Although it is a weak acid, this acid is extremely reactive because of the F- ion. Must be kept in special polypropylene container because it eats through glass. Used to etch glass. Has caused major accidents in lab.

• Substances that increase the OH- when added to water. (Na. OH) • NH 3 is a base. In water it accepts an H+ ion from HOH, leaving an OH- in solution. – NH 3 is a weak electrolyte – About 1% ionizes to form NH 4+/OH-

Strong acids and bases • Acids and bases that ionize completely in solution are strong acids and bases. • Those that only ionize partially are weak acids and bases. • You must memorize these.

Strong Acids Hydrochloric Acid – HCl Hydrobromic Acid – HBr Hydroiodic Acid – HI Nitric Acid – HNO 3 Sulfuric Acid – H 2 SO 4 Chloric Acid – HCl. O 3 Perchloric Acid – HCl. O 4

Strong Bases All group 1 Metal Hydroxides (Li. OH, Na. OH, KOH, Rb. OH, Cs. OH) Heavy Group 2 Metal Hydroxides Ca(OH)2, Sr(OH)2, Ba(OH)2

Once you memorize the strong acids and bases, you will have enough information to determine if a substance is a strong or weak electrolyte.

Example problems: KF Na 3 PO 4 NH 3 CH 3 CH 2 OH HCl NO 2 HC 2 H 3 O 2 CH 4 NH 4 Cl CH 3 Cl strong electrolyte weak electrolyte nonelectrolyte strong electrolyte nonelectrolyte

Acid + Base Neutralization • Products of a neutralization reaction have none of the properties of an acid or a base. • An acid reacts with a metal hydroxide to form a salt plus water.

Neutralization Reactions • Acid + Base (Metal Hydroxide) Salt + Water • HCl(aq) + Na. OH(aq) Na. Cl(aq) + H 2 O(l) • H+ + Cl- + Na+ + OH- Na+ + Cl- + H 2 O(l) H+ + OH- H 2 O(l)

Write the net ionic equation for the following reaction. It might help to first write the molecular equation, and then the complete ionic equation, followed by the net ionic equation. Potassium Hydroxide + Sulfuric Acid • Ionic equation: • Net Ionic equation:

Neutralization Reaction of Weak Acid *Remember, only strong electrolytes are written as ions. * Acetic Acid + Sodium Hydroxide HC 2 H 3 O 2(aq) + Na. OH(aq) Na. C 2 H 3 O 2(aq) + H 2 O(l) Weak acid strong base soluble salt water HC 2 H 3 O 2 + Na+ + OH- Na+ + C 2 H 3 O 2 - + H 2 O(l) HC 2 H 3 O 2(aq) + OH (aq) C 2 H 3 O 2 (aq)+ H 2 O(l)

Acid/Base Rx’s with gas formation • Other bases besides OH- react with H+ to form molecular compounds. Two common -2 -2 bases are CO 3 and S. • Carbonates and bicarbonates react with acid to form CO 2.

Hydrochloric acid + Sodium Sulfide 2 HCl (aq) + Na 2 S(aq) H 2 S(g) + 2 Na. Cl(aq) 2 H+ (aq) + S 2 -(aq) H 2 S(g) Hydrochloric acid + Sodium Hydrogen Carbonate HCl (aq) + Na. HCO 3(aq) Na. Cl(aq) + H 2 CO 3(aq) H 2 O(l) + CO 2(g) HCl (aq) + Na. HCO 3(aq) Na. Cl(aq) + H 2 O(l) + CO 2(g) H+ (aq) + HCO 3 -(aq) H 2 O(l) + CO 2(g)

Reactions in which electrons are transferred between substances

Use of Oxidation numbers in determining redox reactions is basically a bookkeeping method for keeping track of electrons You must be able to identify an oxidation-reduction reaction. But first, we must learn the rules for assigning oxidation #’s to different species.

Rules for oxidation numbers 1) Atoms in elemental form are 0. 2) Monatomic ion; charge of the ion is its oxidation number. 3) Nonmetals; usually negative numbers. a. ) oxygen = -2 unless a peroxide = -1 b. ) Hydrogen +1 with nonmetals, -1 with metals c. ) Halogens (-1) unless bonded to oxygen (+) in a polyatomic ion (Ex: Cl. O 3 -; Cl = +5) 4) Sum of oxidation numbers must = 0 5) Most electronegative (furthest to right and up) element gets a negative charge. See pages 128 – 129 for more on this.

1) Atoms in elemental form are 0. Examples Ag Pb Cl 2 Oxidation # = 0 for 7 diatomic elements and for all other elements when by themselves.

2) Monatomic ion-charge of the ion is its oxidation number. Examples Ag. Cl Ag = +1 Pb. I 2 Pb = +2 Fe 2 O 3 Fe = +3 Cl = -1 I = -1 O = -2

3) Nonmetals; usually negative numbers. a. ) oxygen = -2 unless a peroxide = -1 b. ) Hydrogen +1 with nonmetals, -1 with metals c. ) Halogens (-1) unless bonded to oxygen (+) in a polyatomic ion (Ex: Cl. O 3 -; Cl = +5) Pb. O H 2 S KI Examples oxygen = -2 Na 2 O 2 hydrogen = +1 Na. H iodine = -1 KIO 2 oxygen = -1 hydrogen = -1 iodine = + 3

Determine Oxidation # of element red element in each of the following: Mn. O 2 +4 Br 2 0 KMn. O 4 HCl. O 4 +7 +7 Br. O 2 - H 2 SO 4 +3 +6 Br. O 3 - PO 33 - +5 +3 Ca. H 2 -1 SO 42+6 Na 2 S -2 Mg(NO 3)2 +5

Again, oxidation reduction reactions occur when there is a transfer of electrons from one species to another in a reaction. • If one reactant gains electrons another must lose electrons. • Reduction is always accompanied by oxidation.

Oxidation-Reduction Reactions • An atom that becomes more positively charged is oxidized. – This is due to loss of e-. • The gain of electrons by an atom is called reduction.

Two mnemonics for remembering which substance is undergoing oxidation and which is undergoing reduction? OIL -- RIG Oxidation Involves Loss -- Reduction Involves Gain “Leo the lion says Ger” Loss of electrons oxidation -- Gain of electrons reduction

Many metals react with O 2 in the air to form metal oxides. Metals lose electrons to oxygen. 2 Fe + O 2 2 Fe. O As Fe is oxidized (loses e-), oxygen is reduced (gains e-). Reduction is gain

2 Fe + oxidation O 2 2 Fe. O reduction

Oxidation of metals by acids and salts • Reaction of a metal with either an acid or metal salt follows general form of: A + BX AX +B • Single displacement reaction

+2+6 -2 0 Cu. SO 4(aq) + Zn(s) Zn. SO 4(aq) + Cu(s) Reduced Oxidized What are the products? What are the charges on each species? What is oxidized and what is reduced?

For the following reactants: 1) Write the reaction that occurs. 2) Identify what is being oxidized and reduced. Magnesium + Hydrochloric Acid Aluminum + Cobalt(II) Nitrate

Mg(s) + HCl (aq) oxidation Mg. Cl 2 (aq) + H 2(g) reduction

Al(s) + Co(NO 3)2 (aq) Al(NO 3)3(aq) + Co(s) oxidation reduction

Types of Redox Reactions • Combination (synthesis) • Decomposition • Displacement – hydrogen, metal, halogen • Disproportionation (When an element is simultaneously oxidized and reduced). • Ex: H 2 O 2 H 2 O + O 2

Activity Series • List of metals in order of decreasing ease of oxidation. • Alkali and alkaline earth metals are at the top. (active metals) • Gold, Silver, Platinum, and palladium are considered to be (noble metals) because they resist oxidation.

Using activity series to predict reactions • Activity series can be used to predict reactions between metals and metal salts or acids. • Any metal listed on the series can be oxidized by the ions of elements below it on the list.

Using activity series of metals, which metals from the list below can be oxidized by H+? Ni Al Cu Pb Ag Mg Au

This should be a review before handing out All reaction types worksheet. Must edit all reaction type worksheet. It has to many weak acids reacting with strong bases, which is confusing for students. Make sure most have reactions (see note in folder on handout). Chemical Reaction Types • Decomposition • Synthesis • Single Replacement • Precipitation • Neutralization • Combustion Types of Redox Reactions Types of Double Replacement Reactions

What are the 7 Diatomic Elements H 2 – hydrogen N 2 – nitrogen O 2 – oxygen F 2 - fluorine Cl 2 – chlorine Br 2 – bromine I 2 – Iodine

Synthesis A + B AB Examples H 2 (g) + O 2 (g) H 2 O (g) Mg (s) + O 2 (g) Mg. O (s) Na (s) + Cl 2 (g) Na. Cl (s)

Decomposition AB A + B Examples Na. Cl (s) Na (s) + Cl 2 (g) elec KCl. O 3 (s) KCl (s) + O 2 (g)

Single Replacement A + BC AC + B Examples Na (s) + HOH (l) Na. OH (aq) + H 2 (g) sodium replaces hydrogen in water Cl 2 (g) + Na. Br (aq) Br 2 (l) + Na. Cl (aq) chlorine replaces bromine in sodium bromide

Double Replacement (Metathesis) AB + CD AD + BC Examples Ag. NO 3 (aq) + Na. Cl (aq) Ag. Cl (s) + Na. NO 3 (aq) HCl (aq) + Na. OH (aq) Na. Cl (aq) + H 2 O (l) NH 4 Cl (aq) + Na. OH (aq) NH 3 (g) + H 2 O (l) + Na. Cl (aq) Blue color for the products represents the driving force which allows the chemical reaction to occur.

Combustion hydrocarbon + oxygen carbon dioxide + water Examples CH 4 (g) + O 2 (g) CO 2 (g) + H 2 O (l) C 3 H 8 (g) + O 2 (g) CO 2 (g) + H 2 O (l) CH 3 OH (g) + O 2 (g) CO 2 (g) + H 2 O (l) CO 2 + H 2 O

Neutralization Reactions Strong Acid + Strong Base Salt + Water Example HCl (aq) + Na. OH (aq) Na. Cl (aq) + H 2 O (l) H+ + Cl- + Na+ + OH- Na+ + Cl- + H 2 O(l) H+ + OH- H 2 O

Concentration (Molarity) • Concentration — the amount of solute per unit of solution. • Molarity (M) — expresses concentration:

Calculate the molarity of a solution that contains 20. 0 g copper(I) chloride and has a total volume of 300. 0 m. L. Must have the correct units.

Given: 20. 0 grams Cu. Cl; 300. 0 ml solution Need: Moles Cu. Cl; L of solution 20. 0 g Cu. Cl 1 mol Cu. Cl 99. 0 g Cu. Cl 300. 0 m. L 1000 m. L = 0. 202 mol Cu. Cl = 0. 3000 L

0. 673 M Cu. Cl

How many grams of Na. Cl are needed to make 2. 5 L of 0. 20 molar solution? Given: 0. 20 M solution; 2. 5 L solution Need: grams of Na. Cl (L) Must calculate # of moles, and then convert it into grams. mol = ML

Mol = ML Mol = 0. 20 mol Na. Cl x 2. 5 L = 0. 50 mol Na. Cl L 0. 50 mol Na. Cl 55. 84 g Na. Cl 1 = mol Na. Cl 28 g Na. Cl

Dilution of Stock Solutions • Chemicals are purchased in concentrated form. They need to be diluted for most lab use. • Formula for dilution: Mi Vi = M f Vf

How much stock (12 M) HCl (aq) is required to make 200. 0 m. L of 3 M HCl (aq)? Mi Vi = M f Vf Mi = 12 M Vf = 200. 0 m. L Mf = 3 M Vi = ? m. L Must rearrange equation above to solve for initial volume.

50. m. L Measure 150 m. L of water in a beaker. Slowly add 50. 0 m. L of 12 M HCl for a final volume of 200. 0 m. L.

Two things to note: 1) Always add concentrated acid to water, and not the reverse to avoid unwanted splashing due to the heat generated. 2) When diluting a solution, the amount of solute doesn’t change, only the final volume.

If diluting a solution other than acids, start with initial volume of concentrated solution, and then dilute with distilled water until you have the desired volume. You try one!

We want to prepare 500. m. L of 1. 00 M acetic acid from a 17. 5 M stock solution of acetic acid. What volume of the stock solution is required? Mi Vi = M f Vf Mi = 17. 5 M Vf = 500. m. L Mf = 1. 00 M Vi = ? m. L

28. 6 m. L Pour 471. 4 m. L of distilled water into a beaker. Slowly pour the 28. 6 ml of acid into the water and swirl. Fill the container with distilled water to 500. m. L.

There may be times when you must consider the concentration of ions in a solution. (You must consider the subscripts for this) Mg. Cl 2 Mg 2+ + 2 Cl- In a solution of 0. 25 M Mg. Cl 2 you have: M of Mg 2+ = 0. 25 M M of Cl- = 2 x 0. 25 M = 0. 50 M What is the concentration of each ion in the following? 0. 15 Na 3 P M of Na+ = 0. 45 M M of P 3 - = 0. 15 M

Titrations • Determining the concentration of an unknown solution. • Use a 2 nd solution of known concentration (standard solution) that undergoes a reaction with the unknown solution. • Use the ratios in the balanced equation along with the M = mol/L equation to determine molarity of unknown.

• The point at which the two solutions are stoichiometrically equal is known as the equivalence point. – The reaction is complete and no excess reactant is present. – How do we know when this occurs during the reaction?

• In acid base reactions dyes known as indicators are used. – Phenolphthalein is colorless in acid solution, and pink in basic solution. – End point is reached when a drop of the base remains pink. There is no acid for this drop to react with and the solution is now basic.