Lecture 4- Aqueous Reactions And Solution Stoichiometry

Chemistry 111 General Chemistry Dr. Rabih O. Al-Kaysi Ext: 47247 Email: [email protected]

Lecture 4 Aqueous Reactions and Solution Stoichiometry

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1 - Solution Composition • Solutions are homogenous mixtures of two or more substances: • Solute: present in smallest amount and is the substance dissolved in the solvent. • Solvent: present in the greater quantities and is used to dissolve the solute. • Example: NaCl dissolved in Water (water = Solvent and NaCl = Solute) • Change concentration by using different amounts of solute and solvent. • Molarity: Moles of solute per liter of solution.

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2 - Concentrations of Solutions Formula for Molarity moles of solute Molarity = volume of solution in liters • The most widely used way of quantifying concentration of solutions in chemistry. Molarity is generally represented by the symbol M and defined as the number of moles of solute dissolved in a liter of solution.

Class Guided Practice Problem • Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate, Na2SO4, in enough water to form 125 mL of solution.

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Determining Mass using Molarity • Key concept: If we know: molarity and liters of solution, we can calculate moles (and mass) of solute.

Class Guided Practice Problem • How many grams of C6H12 O6 are required to make 100 mL of 0.278 M C6H12 O6?

Class Practice Problem • How many grams of NaCl are required to make a 1 L of 0.500 M NaCl?

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3 - Concentration of Diluted Solutions

• Solutions are routinely prepared in stock solutions form. •

Example: 12 M HCl

• Solution of lower concentrations are prepared by adding more solvent (e.g., water), a process called dilution. • We recognize that the number of moles are the same in dilute and concentrated solutions. •

Hence, moles solute before dilution = moles solute after dilution

• So: Mdilute Vdilute = Mconcentrated Mfinal Vfinal

Vconcentrated

or = Minitial Vinitial

Class Guided Practice Problem • How much 3.0 M H2SO4 would be required to make 500 mL of 0.10 M H2SO4?

Class Practice Problem • How many milliliters of 5.0 M K2Cr2O7 solution must be diluted in order to prepare 250 mL of 0.10 M solution?

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4 - General Properties of Aqueous Solutions

Electrolytic Properties • Aqueous solutions, solutions in water, have the potential to conduct electricity. • The ability of the solution to conduct depends on the number of ions in solution. • There are three types of solution: • Strong electrolytes, • Weak electrolytes, and • Nonelectrolytes.

General Properties of Aqueous Solutions Electrolytic Properties

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General Properties of Aqueous Solutions Molecular Compounds in Water • Molecular compounds in water (e.g., CH3OH): no ions are formed. • If there are no ions in solution, there is nothing to transport electric charge.

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General Properties of Aqueous Solutions Ionic Compounds in Water

• • • •

Ions dissociate in water (NaCl). In solution, each ion is surrounded by water molecules. Transport of ions through solution causes flow of current. Other substances that are not ionic compound dissociate in water to form ions. • For example: (HCl)

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General Properties of Aqueous Solutions Strong and Weak Electrolytes

• Strong electrolytes: completely dissociate in solution. For example:

HCl(aq)

H+(aq) + Cl-(aq)

• Weak electrolytes: produce a small concentration of ions when they dissolve. • These ions exist in equilibrium with the unionized substance. • For example:

HC2H3O2(aq)

H+(aq) + C2H3O2-(aq)

General Properties of Aqueous Solutions Some General Terms • Acids - substances that able to ionize in solution to form hydrogen ion (H+) and increase the concentration of H+ in the solution. • For example, HCl dissociate in water to form H+ and Cl- ions.

• Bases - are substances that can react with or accept H+ ions. • For example, OH- will accept H+ from HCl forming H2O.

• Salts - are ionic compounds that can be formed by replacing one or more of the hydrogen ions of an acid by a different positive ion. • For example, NaCl instead of HCl.

General Properties of Aqueous Solutions Identifying Strong and Weak Electrolytes • Most salts are strong electrolytes (NaCl, CaCO3. • Most acids are weak electrolytes. However, HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4 are strong acids. • The common strong bases are the hydroxides, Ca(OH)2, of the alkali metals and the heavy alkaline earth metals. • Most other substances are nonelectrolytes.

5 - Precipitation Reactions Exchange (Metathesis) Reactions • When two solutions are mixed and a solid is formed, the solid is called a precipitate. • Metathesis reactions involve swapping ions in solution: AX + BY → AY + BX. HCl + NaOH → NaCl + H2O • Metathesis reactions will lead to a change in solution if one of three things occurs: – an insoluble solid is formed (precipitate), – formation of either a soluble weak or nonelectrolytes, – an insoluble gas is formed.

Class Practice Problem • Write a balanced equation for the reaction between phosphoric acid, H3PO4, and potassium hydroxide, KOH. •H3PO4 + 3KOH → 3H2O + K3PO4 •AX + BY → AY + BX

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Precipitation Reactions Writing Reaction Equations • Ionic equation: used to highlight reaction between ions. • Molecular equation: all species listed as molecules: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq) • Complete ionic equation: lists all ions: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → H2O(l) + Na+(aq) + Cl-(aq) • Net ionic equation: lists only unique ions: H+(aq) + OH-(aq) → H2O(l)

Class Guided Practice Problem • Write the net ionic equation for the reactions that occur when solutions of KOH and Co(NO3)2 are mixed.

•2OH- + Co2+ → Co(OH)2

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6 - Acid-Base Reactions • Acids with one acidic proton are called monoprotic (e.g., HCl). • Acids with two acidic protons are called diprotic (e.g., H2SO4). • Acids with many acidic protons are called polyprotic.

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Acid-Base Reactions Neutralization Reactions and Salts • Neutralization occurs when a solution of an acid and a base are mixed: • HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq) • Notice we form a salt (NaCl) and water. • Salt = ionic compound whose cation comes from a base and anion from an acid. • Neutralization between acid and metal hydroxide produces water and a salt.

Acid-Base Reactions Acid-Base Reactions with Gas Formation • Sulfide and carbonate ions can react with H+ in a similar way to OH−. 2HCl(aq) + Na2S(aq) → H2S(g) + 2NaCl(aq) 2H+(aq) + S2- (aq) → H2S(g) HCl(aq) + NaHCO3(aq) → NaCl(aq) + H2O(l) + CO2(g)

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7 - Oxidation-Reduction Reactions Oxidation and Reduction • When a metal undergoes corrosion it loses electrons to form cations: Ca(s) +2H+(aq) → Ca2+ (aq) + H2(g) • Oxidized: atom, molecule, or ion becomes more positively charged. – Oxidation is the loss of electrons.

• Reduced: atom, molecule, or ion becomes less positively charged. – Reduction is the gain of electrons.

Oxidation-Reduction Reactions Oxidation and Reduction

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Oxidation-Reduction Reactions Oxidation Numbers • Oxidation number for an ion: the charge on the ion. • Oxidation number for an atom: the hypothetical charge that atom would have if it was an ion. • Oxidation numbers are assigned by a series of rules: 1. If the atom is in its elemental form, the oxidation number is zero. E.g., Cl2, H2, P4. 2. For a monoatomic ion, the charge on the ion is the oxidation state.

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Oxidation-Reduction Reactions Oxidation Numbers 3. Nonmetal usually have negative oxidation numbers: a) Oxidation number of O is usually –2. The peroxide ion, O22- , has oxygen with an oxidation number of –1. b) Oxidation number of H is +1 when bonded to nonmetals and –1 when bonded to metals. c) The oxidation number of F is –1. 3. The sum of the oxidation numbers for the atom is the charge on the molecule (zero for a neutral molecule).

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Oxidation-Reduction Reactions Oxidation of Metals by Acids and Salts • Metals are oxidized by acids to form salts: Mg(s) +2HCl(aq) → MgCl2(aq) + H2(g) • During the reaction, 2H+(aq) is reduced to H2(g). • Metals can also be oxidized by other salts: Fe(s) +Ni2+ (aq) → Fe2+ (aq) + Ni(s) • Notice that the Fe is oxidized to Fe2+ and the Ni2+ is reduced to Ni.

Oxidation-Reduction Reactions Activity Series • Some metals are easily oxidized whereas others are not. • Activity series: a list of metals arranged in decreasing ease of oxidation. • The higher the metal on the activity series, the more active that metal. • Any metal can be oxidized by the ions of elements below it.

Solution Stoichiometry and Chemical Analysis Titrations

Solution Stoichiometry and Chemical Analysis Titrations • Suppose we know the molarity of a NaOH solution and we want to find the molarity of an HCl solution. • We know: –

molarity of NaOH, volume of HCl.

• What do we want? –

Molarity of HCl.

• What do we do? –

Take a known volume of the HCl solution, measure the mL of NaOH required to react completely with the HCl.