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Chapter 4

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Brock University
Lydia W.L.Chen

Chapter 4: Reactions inAqueous Solutions 4.1 – General Properties ofAqueous Solutions • Asolution is a homogenous mixture of two or more substances (solute and solvent). • Asolute is the substance present in a smaller amount and the solvent is the substance present in a larger amount. • An aqueous solution is a solution in which the solute is a liquid or solid, and the solvent is water. Electrolytic Properties • All solutes that can dissolve in water can be classified as electrolytes or nonelectrolytes. o An electrolyte is a solute that, when dissolved in water, results in a solution that can conduct electricity. o Anonelectrolyte is a solute that, when dissolved in water, results in a solution that cannot conduct electricity. • Pure water is a poor conductor of electricity since it does NOT contain any dissolved solutes. However, tap water contains many dissolved solutes and is able to conduct electricity well. • When a light bulb, power source, and two electrodes are immersed in an aqueous solution, the light bulb can turn on brightly, dimly, or not even turn on, depending on the type of solute dissolved in the solution. o Strong electrolytes such as ionic compounds, strong acids, or strong bases dissociate or ionize completely (100%) when dissolved in water. When dissociated, their ions undergo hydration. Strong electrolytes cause the light bulb to shine brightly with light.  When NaCl is dissolved in H O, 2he H O mo2ecules cause NaCl to dissociate completely into its ions and then the H O molecules surround each ion in a specific manner (O ends 2 surround Na and H ends surround Cl ). This process is called hydration and prevents the Na and Cl from reacting with one another and reforming the NaCl. Then, Na get + - attracted to the negative electrode and the Cl get attracted to the positive electrode. Eventually, a pathway of ions will form from one electrode to the other electrode, allowing the current flow through the entire circuit. That lights up the bulb. o Weak electrolytes such as weak acids and weak bases do not completely ionize when dissolved in water. They cause the light bulb to be dimly lit.  When HF is dissolved in H O, 2t does not completely ionize into H and F.This provides a weak pathway for the flow of electrons between the electrodes, causing the light bulb to be dimly lit.  HF is not completely ionized since the ionization of HF stops when equilibrium is reached. That is when the rate of ionization is equal to the rate of HF reforming. + - HF ↔ H + F o Nonelectrolytes such as molecular compounds do not dissociate into ions  When molecular compounds such as methanol (CH OH) are 3laced in water, hydrogen bonds form between the O and H atoms of the hydroxyl group and the O and H atoms of the H O molecules, causing the molecule to dissolve in water. However, since molecular 2 compounds do not dissociate, their solutions do not conduct electricity. Weak acids/bases have reversible equations (↔), whereas strong acids/bases do not (→). http://highered.mcgraw- hill.com/olcweb/cgi/pluginpop.cgi? it=swf::100%25::100%25::/sites/dl/free/0072512644/117354/07_Strong_Weak_Nonelectrolytes.swf::Strong %20Electrolytes,%20Weak%20Electrolytes,%20and%20Nonelectrolytes 4.2 – Precipitation Reactions • Precipitation reactions occur in an aqueous solution and result in the formation of a precipitate. • Aprecipitate is an insoluble solid that separates from the solution. • Many precipitation reactions involve ionic compounds and are double displacement reactions. o Double displacement reactions can be called metathesis reactions. Solubility • Solubility of a solute is the maximum amount of solute that will dissolve in a given amount of solvent at a given temperature. o Chemists refer to solutes as soluble, slightly soluble, or insoluble. o Insoluble solutes are still soluble, but not that much as slightly soluble or soluble solutes. • The solubility table below can be used to determine if a compound is soluble or insoluble. o Look at anions first, and then the cation soluble/insoluble exceptions. Molecular Equations, Ionic Equations, and Net Ionic Equations • Amolecular equation is an equation in which the formulas of all the compounds are written as though all the species existed as molecules or whole units. o Pb(NO ) 3 2 (aq)2KI (aq) PbI 2 (s) 2KNO 3 (aq)olecular Equation) • Molecular equations are very useful when performing an experiment as they tell us the compounds that are reacting. • An ionic equation is an equation that shows the dissolved ionic compounds as ions. The precipitate occurs in a solid state, not dissolved, in the reaction. o Pb 2+(aq) 2NO 3 (aq) 2K +(aq)+ 2I-(aq) PbI 2 (s) 2NO 3 (aq) 2K +(aq)Ionic Equation) • The spectator ions are the ions that do not participate in the overall reaction. o They exist as ions in the solution, before the reaction starts and after it finishes. • The net ionic equation is an equation that only shows the species that take part in the reaction. It is obtained by cancelling the spectator ions in the ionic equation. 2+ - + - - + o Pb (aq) 2NO 3 (aq) 2K (aq)+ 2I (aq) PbI 2 (s) 2NO 3 (aq) 2K (aq) o Pb 2+(aq) 2I-(aq) PbI 2 (s)et Ionic Equation) • Below is a diagramic representation of this reaction: 4.3 –Acids-Base Reactions Acids and Bases • SvanteArrhenius defined acids as substances that ionizes in water to produce H and bases as substances - that ionize in water to produce OH. Acids • Acids taste sour (i.e. lemons, vinegar) • Acids cause colour changes in plant dyes. They change the colour of litmus from blue to red. • Aqueous acid solutions conduct electricity. • Acids react with metals above H on the activity series (Mg, Zn, Fe) to produce H 2 (g) o 2HCl (aq)+ Mg →(s)Cl 2(s) H 2 (g) • Acids react with metal carbonates and metal bicarbonates to produce H O 2 (l)nd CO 2 (g) o Na C2 3(s) 2HCl (aq)2 NaCl + (s) + C2 (l) 2 (g) Bases • Bases taste bitter. • Bases feel slippery (i.e. soap) • Bases cause colour changes in plant dyes. They change the colour of litmus from red to blue. • Aqueous base solutions conduct electricity. BronsteadAcids and Bases • Bronstead defined acids as proton donors, and bases as proton acceptors. o UnlikeArrhenius definition, Bronstead definition showed that acids and bases are not limited to aqueous states. • When an acid is placed in water, it ionizes and forms an anion and H . In water, the H become hydrated by H O2molecules. o Therefore, in water, H attract to many H O mol2cules, forming many ions of the form H(H O) , 2 n- where n represents the amount of water molecules the H is attracted to or hydrated by. + +  H 3 are hydrated H and are called hydronium ions. • Monoprotic Acids are acids that ionize into only one H in water. o Since the ionization of acetic acid in water is a reverse reaction, it does not completely ionize into H (double arrow).  This is why acetic acid is a weak acid and a weak electrolyte. o Since HCl and HNO ionize completely in water, they are strong acids (aq) 3 (aq) and strong electrolytes (single arrow). + • Diprotic Acids are acids that ionize into two H in water. - o H 2O 4 (aq) a strong acid and electrolyte. However HSO 4 (aq)s a weak acid electrolyte. • Triprotic Acids are acids that ionize into three H in water. o H 3O 4 (aq)H2PO 4 (aq)nd HPO 42-(aq)re all weak acids and electrolytes. Base Acid ConjugateAcid Conjugate Base - + • Since OH can accept a H and form H O, it is 2 base (conjugate base). + + • Since NH ca4 donate a H and form NH , it is an 3cid (conjugate acid). • All alkali metal hydroxides are strong bases and electrolytes. o Ba(OH) is2the only alkaline earth metal hydroxide that is soluble in water. + - • NH i3 a weak base because it reacts with H O to fo2m very few NH and OH. 4 Acid-Base Neutralizations • Aneutralization reaction is a reaction between an acid and a base. It usually produces a salt and water. StrongAcid and Strong Base Neutralization Reactions Chemical Equation: NaOH (aq) HCl (aq)→ NaCl (aq) H 2 (l) The strong acids and bases ionize completely in water. Total Ionic Equation: Na +(aq) OH (aq)+ H +(aq)+ Cl -(aq) Na +(aq) Cl (aq)+ H 2 (l) + - Na (aq)nd Cl (aq)re spectator ions. - + Net Ionic Equation: OH (aq) H (aq)→ H O 2 (l) • If we started with equal molar amounts of acid and base, we would end up with no excess acid or base. Weak Acid and Strong Base Neutralization Reactions Chemical Equation: NaOH + HCN → NaCN + H O (aq) (aq) (aq) 2 (l) HCN (aq)s a weak acid and does not ionize completely in H O2 + - + - Total Ionic Equation: Na (aq) OH (aq)+ HCN (aq)→ Na (aq) Cl (aq) H O2 (l) + Na (aq)s the spectator ions. Net Ionic Equation: OH -(aq) HCN (aq) Cl -(aq) H O2 (l) HNO 3 (aq) NH 3 (aq) NH NO 4 3 (aq) • Since NH NO4is i3 an aqueous state, it consists of NH and NO . T4e NH is com3osed of NH 4nd 3 H 2. o Therefore, water is also produced in this neutralization reaction. o HNO 3 (aq) NH 3 (aq) NH 3 (aq) H 2 +(l)3 (aq) Acid-Base Reactions Leading To Gas Formation • Certain salts such as carbonates, bicarbonates, sulphites and sulphides react with acids to form gaseous products. 4 .4 – Oxidation- Reduction Reactions • Oxidation-reduction reactions (also called redox reactions) are reactions that involve the transfer of electrons. o Many redox reactions take place in water (aqueous states), but not all do. • Half reactions are equations that show the electrons involved in a redox reaction. Every redox reaction consists of an oxidation half reaction and a reduction half reaction. o Oxidation half reactions show the loss of electrons in the redox reaction. o Reduction half reactions show the gain of electrons in the redox reaction. • Adding both half reactions gives us the chemical equation. • The reducing agent is the reactant that causes the reduction of another reactant in the reaction. In the process, it gets oxidized. • The oxidizing agent is the reactant that causes the oxidation of another reactant in the reaction. It gets reduced in the process. Chemical Equation: 2Mg + O (s) 2 (g) 2MgO (s) o In the reaction below, Mg is t(s)reducing agent 2+ - Oxidation Half Reaction: 2Mg → 2Mg (s) (s)+ 4e and O 2(g)is the oxidizing agent.
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