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Chapter 4

Chapter 4. Reactions in Aqueous Solution. Solutions. A solution is a homogeneous mixture composed of a solute and a solvent. The solute is dissolved in the solvent.

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Chapter 4

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  1. Chapter 4 Reactions in Aqueous Solution

  2. Solutions A solutionis a homogeneous mixture composed of a solute and a solvent. The solute is dissolved in the solvent. A solute is the substance present in smaller amount in a solution. It can also be thought of as the dissolved material. A solvent is the substance present in larger amount in a solution. It can also be thought of as the dissolving material. A solution can be gaseous (air), solid (alloy), or liquid (seawater). In this chapter, we will only discuss aqueous solutions, in which the solute initially is a liquid or a solid and the solvent is water.

  3. Molarity and Dilution Molarity (M) is the number of moles of the solute per liters of solution. Molarity (M) = moles of solute = mol Liters of Solution L

  4. When Diluting use: M1V1 = M2V2

  5. Hydration Water is a very effective solvent for ionic compounds. Although water is an electrically neutral molecule, it has a positive (H atoms) and negative (O atoms) region, or “positive and negative poles”. This is why water is called a polar solvent. When an ionic compound, such as NaCl, dissolves in water, the three dimensional network of ions in the solid is destroyed. The Na+ and the Cl- ions are separated from one another and they undergo hydration. Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner. Let’s see this in action.

  6. NaCl(s) H2O Na+(aq) + Cl-(aq) The term Dissociation means that the compound breaks up into cations and anions like in the above equation for salt. Solid NaCl, salt, is an ionic compound and breaks up into Na+ and Cl-, cations and anions when dissolved in water. The Na+ ions are attracted to the negative electrode and the Cl- anions are attracted to the positive electrode. This movement sets up an electric current that is equivalent to the flow of electrons along a metal wire. Because NaCl conducts electricity, we say that NaCl is an electrolyte. Pure Water contains very few ions and therefore we call it a nonelectrolyte. The above equation also shows that all of the salt has dissociated into ions and there is no undissociated NaCl left over in the solution. This would be the same as saying that salt is a very strong electrolyte.

  7. Ionic vs. Molecular Compounds Dissolving in Water • As we have already talked about in earlier chapters, Molecular Compounds dissolve in water but they break apart into molecules floating around in the water and therefore they do not have a charge. This is why they are called Molecular Compounds! • Ionic Compounds break apart into ions when they are dissolved in water. The ions have a charge in the solution and are good electrolytes! Remember all bases are ionic except ammonia! • Exception: Acids and the weak base ammonia (NH3) are considered molecular compounds but they do break apart into ions in water: NH3 + H2O D NH4+ + OH- HCl + H2O D H3O+ + Cl- • Also NOTE: Dissolving in water does not make something a strong electrolyte (think of sugar = molecular and dissolves in water, it is not a strong electrolyte!)

  8. All solutes that dissolve in water fit into one of 2 categories: electrolytes and nonelectrolytes. An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolytedoes not conduct electricity when dissolved in water.

  9. Always USE g arrow Always USE D arrow Strong Electrolyte Weak Electrolyte Nonelectrolyte CH3COOH HF HNO2 NH3 H2O* (all of the above loose an H+ ion when dissociated) *Pure water is an extremely weakelectrolyte (NH2)2CO (urea) CH3OH (methanol) C2H5OH (ethanol) C6H12O6 (glucose) C12H22O11 (sucrose) • HCl • HBr • HI • HNO3 • HClO4 • H2SO4* • All 1A: (LiOH, NaOH, • KOH, RbOH, CsOH) • 2A: Ba(OH)2 & Sr(OH)2 • Ca(OH)2 • Ionic Compounds • H2SO4 has 2 ionizable H+ ions, the second form, HSO4- is a weak electrolyte. The strong/weak parts of this chart should be memorized because it will help you to memorize your strong/weak acids and bases.

  10. Ionization of Acids and Bases We use the term Ionization in order to describe the separation of acids and bases into ions. In order to determine whether or not you have a strong acid or base, you see whether the acid or base dissociates completely in water. If it does completely ionize in water, then it is considered to be a strong acid or base. This would also make it a Strong Electrolyte. To symbolize a strong electrolyte you use a single arrow g. Such as: HCl(aq)g H+(aq) + Cl-(aq) When an acid or base does not completely ionize in water, it is a weak acid/base. These are also called Weak Electrolytes. Acetic acid is a weak acid and we represent the ionization of acetic acid with a double arrowDto show that it is a reversible reaction, or that the reaction can occur in both directions. CH3COOH(aq) D CH3COO-(aq) + H+(aq)

  11. Solubility and Precipitates Solubility is a term that is used in order to describe the amount of solute that will dissolve in a solvent. Most mixtures have a certain amount of solute that will dissolve inside of a solvent. Once the maximum amount of solute particles have been added to a solvent, the solute particles will no longer dissolve in the solvent and they will turn into the solid form of the particle. This change from liquid to solid is called a precipitate. If something is considered to be insoluble in water, then it means it has a solubility of less than .01 M in water at 25 degrees Celsius.

  12. Solubility Rules for Common Ionic Compounds in Water at 25 oC Soluble Compounds Exceptions Insoluble CompoundsExceptions Compounds containing alkali metal ions (Li+, Na+, K+, Rb+, Cs+) and the ammonium ion (NH4+) Nitrates (NO3-), bicarbonates (HCO3-), and chlorates (ClO3-) Halides (Cl-, Br-, I-) Sulfates (SO42-) Memorize these Rules!! Halides of Ag+, Hg22+, and Pb2+ Sulfates of Ag+, Ca2+, Sr2+, Ba2+, Hg2+, and Pb2+ Carbonates (CO32-) Phosphates (PO43-) Chromates (CrO42-) Sulfides (S2-) Hydroxides (OH-) Compounds containing alkali metal ions and the ammonium ion Compounds containing alkali metal ions, ammonium ion and the Ba2+ ion

  13. Molecular Equations and Ionic Equations Pb(NO3)2(aq) + 2NaI(aq)g PbI2(s) + 2NaNO3(aq) The above equation is considered a molecular equation because the formulas of the compounds are written as though all species existed as molecules or whole units. A molecular equation is useful because it identifies the reagents, if we wanted to bring about this reaction in the lab. However, a molecular equation does not accurately describe what actually is happening at the microscopic level. An ionic equation is used to represent what is occurring on the microscopic level. The ionic equation shows dissolved species as free ions. The above molecular equation would have an ionic equation such as: Pb2+(aq) + 2NO3- (aq) + 2Na+(aq) + 2I-(aq)g PbI2(s) + 2Na+(aq) + 2NO3- (aq)

  14. Spectator Ions and Net Ionic Equations An ionic equation includes spectator ions. Spectator ions are ions that are not included involved in the overall reaction. Spectator ions appear on both sides of the equation and are unchanged in the chemical reaction, therefore they can be canceled. In the previous ionic equation: Pb2+(aq) + 2NO3- (aq) + 2Na+(aq) + 2I-(aq)g PbI2(s) + 2Na+(aq) + 2NO3- (aq) You would cancel the 2NO3- and the 2Na+ spectator ions. The net ionic equation, which shows only the species that actually take part in the reaction, would be: Pb2+(aq) + 2I-(aq)g PbI2(s) Note: If everything in a net ionic equation is a spectator ion (and cancels on both sides) then you say: No Reaction Occurred!!

  15. Acidsand Bases Arrhenius defined an acid as producing a H+ ion when dissolved in water. He also defined a base as producing an OH- ion when dissolved in water. This definition was limiting because it only applied to aqueous solutions. The chemists Johannes Bronsted and Thomas Lowry proposed a more broad definition. Their definition does not require that an acid be in an aqueous solution and includes more than just protons and hydroxide ions for acids and bases. A Bronsted-Lowry acidis a proton (H+) donor. A Bronsted-Lowry baseis a proton (H+) acceptor.

  16. Hydrochloric Acid Water Hydronium Ion Chloride Ion Ammonia Water Ammonium Ion Hydroxide Ion

  17. Monoprotic, Diprotic and Triprotic Acids When each unit of acid only produces one hydrogen ion upon ionization (or hydronium ion), that acid is said to be a monoprotic acid. Examples are: hydrochloric (HCl), nitric (HNO3) and acetic acid (CH3COOH).These are common acids you should memorize. HNO3(aq)g H+ + NO3-(aq) (Strong Acid/Electrolyte) When each unit of acid produces 2 hydrogen ions upon ionization (or hydronium ions), that acid is said to be a diprotic acid. Example: Sulfuric Acid (H2SO4). This is also a common acid that you should memorize. H2SO4(aq)g H+(aq)+ HSO4-(aq) (Strong Acid/Electrolyte) HSO4-(aq) DH+(aq) + SO42-(aq) (Weak Acid/Electrolyte) When each unit of acid produces 3 hydrogen ions upon ionization (or hydronium ions), that acid is said to be a triprotic acid. Example: Phosphoric Acid (H3PO4). This is also a common acid that you should memorize. Let’s try and represent the triprotic acid of phosphoric acid on our own.

  18. Acid-Base Neutralization A neutralization reaction is a reaction between an acid and a base. Generally, aqueous acid-base reactions produce water and a salt, which is an ionic compound made up of a cation other that H+ and an anion other that OH- or O2-: acid + base g salt + water HCl(aq) + NaOH(aq)g NaCl(aq) + H2O(l) All salts are strong electrolytes. The substance we know as table salt, NaCl, is a familiar example. However, since both the acid and the base are strong electrolytes, they are completely ionized in solution. The ionic equation is: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)g Na+(aq) + Cl-(aq) + H2O(l) Therefore, the net ionic equation is: H+(aq) + OH-(aq)g H2O(l) Both Na+ and Cl- ions are spectator ions. If we had started the reaction with equal molar amounts of the acid and the base, at the end of the reaction we would have only a salt and no leftover acid or base. This is a characteristic of acid-base neutralization reactions.

  19. Acid-Base Titration In titration, a solution of accurately known concentration, called a standard solution, is added gradually to another solution of unknown concentration, until the chemical reaction between the two solutions is complete. The point at which the acid has completely reacted with or has been neutralized by the base is called the equivalence point. The endpoint is the point at which the solution should change in color due to the indicator, which changes color at or near the equivalence point. Many times, solid sodium hydroxide is not pure because of its absorbance of water from the air. Therefore, when we make up solutions of sodium hydroxide in the lab, we are unsure of it’s exact molarity and we must standardize the solution. Standardizing the solution is when we use the acid-base titration technique in order to verify what the molarity of the unknown solution is.

  20. Diprotic Titrations How many milliliters of a 0.610 M NaOH solution are needed to neutralize 20.0 mL of a 0.245 M H2SO4 solution? The equation for the reaction is: 2NaOH(aq) + H2SO4(aq) g Na2SO4(aq) + 2H2O(l) Notice that we need twice the amount of sodium hydroxide in order to neutralize the diprotic sulfuric acid. This is because: H2SO4(aq)g H+(aq)+ HSO4-(aq) In the above equation, one mole of OH- ions will neutralize one mole of the H+ ions. HSO4-(aq) DH+(aq) + SO42-(aq) In the above equation, one mole of OH- ions will neutralize one mole of the H+ ions. Therefore, we need a total of 2 moles of NaOH in order to neutralize one mole of H2SO4 Let’s Finish this problem on our own….

  21. Gas Forming Reactions Many reactions release a gaseous product. Although a wide variety of these gas-forming reactions occur, some of the most important gases produced in reactions are the following: Acid/Base Rxns that form Gases: S2- (Sulfides) and CO32-(Carbonate), HCO3-(Bicarbonate) ions react with acids to form gases: CO2 Gas formation: A 2 Step Reaction: 1 step – Any CO32- or HCO3- ion reacting with an acid gives Carbonic Acid (H2CO3): Example: HCl(aq) + NaHCO3(aq) g NaCl(aq) + H2CO3(aq) 2 step – Carbonic Acid is unstable and will decompose immediately into CO2 gas and liquid H2O: Example: H2CO3(aq)g CO2(g) + H2O(l) H2S Gas (smells like rotten eggs) formation: forms when an acid (like HCl) reacts with a Metal Sulfide (like Na2S): Example: 2HCl(aq) + Na2S(aq)g H2S(g) + 2NaCl(aq) H2 Gas forms in a single replacement (REDOX) reaction when an acid reacts with a metal: Example: Ca(s) + 2HCl(aq)g CaCl2(aq) + H2(g) Other Rxns that form Gases: O2 formed in many ways, one example: Electrolysis of water NO2 formed when in air when lightning hits to supply energy:2NO(g) + O2(g)g 2NO2(g)

  22. Rules for Assigning Oxidation States • The oxidation state of an atom in an element is 0. Atoms in their elemental form are 0, example H2 is the elemental form of Hydrogen therefore in the H2 molecule, each H = 0 (all diatomic atoms are the same) or in P4 each P = 0 or in S8 each S = 0. • In a neutral molecule, the sum of the oxidation numbers of all the atoms must equal 0. • The oxidation state of a monatomic ion is the same as its charge. • In its compound, fluorine is always assigned an oxidation sate of -1. • Oxygen is usually assigned an oxidation of -2 in its covalent compounds, such as CO, CO2, SO2. Exceptions to this rule includes peroxides (compounds containing the O22- group), where each oxygen is assigned an oxidation state of -1, as in hydrogen peroxide (H2O2), and OF2 in which oxygen is assigned a +2 oxidation state. • In its covalent compounds with nonmetals, hydrogen is assigned an oxidation state of +1. For example HCl. When hydrogen is bonded to a nonmetal in a binary compound, it is assigned an oxidation state of -1. For example LiH. • For an ion, the sum of the oxidation states must equal the charge of the ion. For example, the sum of the oxidation states must equal -2 in CO32-. • When Halogens combine with Oxygen, then halogens have a + charge.

  23. Oxidation-Reduction Reactions Redox Oxidation-Reduction Reactions are considered electron-transfer reactions. Also known as Redox Reactions. Remember: LEO goes GER Loss of Electrons means Oxidized & Gain of Electrons means Reduced If you say an element is Oxidized, then it is called a Reducing agent because it donates electrons to another element. If you say an element is Reduced, then it is called an Oxidizing agent because it accepts electrons from another element.

  24. Let’s take a look at the formation of calcium oxide (CaO) from calcium and oxygen: 2Ca(s) + O2(g)  2CaO(s) You should recognize this is a REDOX reaction by checking the individual oxidation charges of each atom/ion and verifying that Loss of Electrons & Gain of Electrons is occurring: Loss of electrons: 2Ca  2Ca2+ + 4e- Calcium is therefore being Oxidized Therefore Calcium is a good Reducing Agent! Gain of electrons: O2 + 4e- 2O2- While Oxygen is being Reduced! Therefore Oxygen is a good Oxidizing Agent! Remember: LEO goes GER

  25. MOST ACTIVE Single Replacement Reactions are one type of REDOX Reactions! • Single Replacement Reactions are one type of REDOX reactions (We will cover other types of REDOX reactions later – but you should be able to recognize a REDOX reaction by checking the oxidation charges of atoms/ions). • You must use the ACTIVITY SERIES to determine if a single replacement reaction will actually occur. The Activity Series for Metals is on the right. The Activity Series for nonmetals is on the left. • An element can only replace another element that is less active than itself. Sort of like the top boss can kick out anyone below him and use their office. Examples: 3Mg + 2 FeCl3g 2 Fe + 3 MgCl2 Cl2 + 2KI g I2 + 2KCl Cl2 + KF g No Reaction! F2 Cl2 Br2 I2 LEAST ACTIVE

  26. Redox Reactions In Net Ionic Equation Form Cu(s) + 2NO3-(aq) + 4H+ 2NO2(g) + Cu2+(aq) + 2H2O(l) 2Ag+(aq) + Cu(s) 2Ag(s) + Cu2+(aq) 2NO(g) + O2(g) 2NO2(g) 2Fe(s) + 3Cl2(g) 2FeCl3(s)

  27. Balancing Redox Reaction in Acidic Solutions • Write both the reduction and oxidation half-reactions • a) Balance elements (except H & O) b) Balance oxygen using H2O c) Balance hydrogen using H+ d) Balance the charges using e-s 3. Make the number of electron equal in both half reactions 4. Add half reactions together and cancel out identical species 5. Check to see that elements and charges are balanced

  28. Balancing Redox Reaction in Basic Solutions • 1-5 same as for acidic solutions 6. Add OH- ions to both sides of the equations that equal the number of H+ ions 7. Sides containing both H+ and OH- ion will for water (eliminate H2O molecules) 8. Check that elements and charges are balanced.

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