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Chemical Reactions

Chemical Reactions. Understanding the mechanism of reactions is important to our understanding of processes such as acid rain, corrosion, stain removers. . Chemistry is about reactions with molecules colliding and forming new molecules.

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Chemical Reactions

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  1. Chemical Reactions Understanding the mechanism of reactions is important to our understanding of processes such as acid rain, corrosion, stain removers. Chemistry is about reactions with molecules colliding and forming new molecules. A number of reactions can be classified as "types”, allowing some general predictions on outcomes of reactions. Dissolution Reactions Precipitation Reactions Acid-Base Reactions Oxidation-Reduction Reactions

  2. Dissolution Reactions In dissolution reactions two or more compounds disperse into each other to form a homogenous phase. The starting compounds could be different phases (e.g. a solid and a liquid), but the outcome of dissolution is a homogenous phase; a SOLUTION In dissolution reactions, the compound of lower concentration is called the SOLUTE and the higher concentration component is the SOLVENT.

  3. During dissolution the solvent interacts with the solute such that for the solute, the interactions between the solute and solvent dominate over the solute-solute interactions and solvent-solvent interactions. Dissolution reactions are considered to be intermediate between a chemical and physical process. In terms of it being considered to be a chemical process, solute-solute interactions are broken up and replaced by solute-solvent interactions. On the other hand the solution that results cannot be expressed as a chemical formula and hence the outcome of dissolution cannot be represented as a typical chemical equation.

  4. Note: for dissolution reactions, the solvent need not be water, nor necessarily a liquid Other examples of common liquid solvents are , but benzene (C6H6), acetone (CH3COCH3), carbon tetrachloride (CCl4), methanol (CH3OH) To write an equation for a dissolution reaction the solvent is left out and the change in state of the solute denoted. For example: dissolving sucrose in water C12H22O11(s)  C12H22O11 (aq) s - solid aq - aqueous solution.

  5. Dissolution of Ionic Compounds Most ionic compounds dissolve easily in water. As we have seen ionic compounds, like NaCl, have rigid lattices defined by the oppositely charged ions. For the ionic compound to dissolve in water, the water molecules must overcome the strong interactions that exist between the oppositely charged species so that ion-water (solute-solvent) interactions dominate over ion-ion interactions (solute-solute)

  6. The negative end of the water molecule (O) interacts with the positive ions in the crystal and the positive end of water (H) interacts with the negative end When a water molecule encounters an ion it orients itself so that the appropriate "side" of the water molecule interacts with the ion (for negative ions, the H points toward the ion and for positive ions, the O).

  7. Having oriented itself in this way, the water molecule essentially pulls this ion out of the crystal lattice. Other water molecule surround this one ion and screens the ion from the oppositely charged ions in the crystal. Hence the water molecules, solvate the ion, and the solvated ion then moves through the solution. By this process the ionic compound dissolves in water and is said to DISSOCIATE INTO ITS IONS. Dissolution reactions for ionic compounds are written as: NaCl(s)  Na+(aq) + Cl-(aq)

  8. Dissolution of Covalent compounds Neutral molecules do not have charges by which they can interact with the solvent. However, they can interact with the solvent through their polarity. Since the polarity of a molecule is due to a charge separation within the molecule, solute molecules that are polar can dissolve in polar solvents in much like the way ions dissolve in water (a polar solvent).

  9. For polar solutes, the polar solvent molecules orient themselves around the solute molecules so that the more positive end of the solvent is oriented towards the more negative solute molecule and the more negative end of the solvent orients itself towards the more positive end of the molecule. In this way the solute molecules are solvated replacing the solute-solute interactions by solute-solvent interactions. The solute molecules remain intact, but each solute molecule is solvated by solvent molecules. Dissolution of molecular compounds can be written as: C12H22O11(s)  C12H22O11 (aq)

  10. For a solute molecule to go “into solution”, the solvent molecules must solvate the solute molecule so that the solute-solvent interaction dominate over the solute-solute interactions. Hence, solute molecules dissolve in solutions of the same polarity - “like dissolves like”. So polar solutes dissolve in polar solvents, non-polar solutes dissolve in non-polar solvents Soaps or surfactants are “designed” so that the soap molecule dissolves in water, yet can interact with a non-polar oil molecule in a grease stain

  11. a) Dissolution of an ionic compound b) Dissolution of a covalent compound

  12. Solubilities Ethanol and water dissolve in each other and are said to be miscible. As more ethanol is added, at some point the ethanol concentration is larger than the water concentration and the solute and solvent switch. Water and ethanol have infinite solubilties in the other. Solubility is defined as the amount of a solute that can dissolve in a fixed amount of solvent, at a given temperature Solubility varies with temperature - generally higher the temperature larger is the solubility.

  13. As NaCl is added to water, a point is reached when the NaCl does dissolve, but remains as a solid in the salt solution. The point at which the NaCl stops dissolving in the salt solution, defines the solubility of NaCl in water, at that temperature. The solution is said to be saturated with NaCl. For a liquid solute dissolved in a liquid solvent, at the saturation point, a new layer is formed, with the new layer contains the solute with some of the original solvent dissolved.

  14. battery bulb + - + deionized, pure, water Electrolytes and Non-Electrolytes

  15. battery bulb + - Na+ + Cl- deionized water + NaCl NaCl(s) --> Na+(aq) + Cl-(aq) When the battery is turned on the Na+ ions flow toward the negative plate (anode) and the Cl- ions to the positive plate (cathode).

  16. The flow of ions constitutes a current. The circuit is now complete, current flows through the circuit, and the bulb turns on. NaCl is called an electrolyte. Electrolyte : a compound which when dissolved in a solvent dissociates to form ions in solution. Typically electrolytes are ionic compounds since they dissolve in solution to form ions. Example: K2SO4 Some covalent compounds (like acids and bases) can dissociate in solution to form ions.

  17. Electrolytes are characterized as being strong or weak. The strength of an electrolyte depends on the degree to which the compound dissociates in water to form ions. Hence ionic compounds like NaCl and K2SO4 which dissociate completely in water are strong electrolytes. Weak electrolytes do not dissociate extensively in water- consequently the conductance of a solution of a weak electrolyte in low. Non-electrolytes do not dissociate in solution to form ions and hence their solutions do not conduct electricity.

  18. 2 KI(aq) + Pb(NO3)2(aq) --> PbI2 (s) + 2 KNO3 (aq) Precipitation Reactions

  19. The reaction between the KI and Pb(NO3)2 results in the formation of PbI2 which has a very low solubility in water and forms a solid precipitate. 2 KI(aq) + Pb(NO3)2(aq) --> PbI2 (s) + 2 KNO3 (aq) For compounds insoluble in water, the attraction between the oppositely charged ions in the solid crystal are too strong to be overcome by solvent water molecules. KNO3, being soluble in water, exists in solution as K+ and NO3- ions.

  20. The reaction 2 KI(aq) + Pb(NO3)2(aq) --> PbI2 (s) + 2 KNO3 (aq) is also called a METATHESIS reaction. In a metathesis reaction atoms or groups of atoms are switched. In this example K+ and Pb2+ switch anions. NOTE: Metathesis reactions do not have to result in precipitation.

  21. Precipitation reactions result when 1) an amount exceeding the compound’s solubility in a particular solvent is added to the solvent 2) Removal of solvent - example by evaporation. 3) Changing the solvent - since solubilities vary from solvent to solvent, changing the solvent can result in precipitation 4) Changing the temperature - solubilities vary with temperature. Cooling a solution of a compound can result in the compound precipitating out of solution.

  22. Predicting Precipitation Reactions To determine if a product of a reaction between two ionic compounds will result in a precipitate being formed, check the solubilities of the compounds formed by the reaction.

  23. Looking at the table Example: If aqueous solutions of AgNO3 and NaCl are mixed will a precipitate form? If yes, identify the precipitate. AgNO3 (aq) + NaCl (aq) --> AgCl (?) + NaNO3 (?) AgCl is insoluble, NaNO3 is soluble AgNO3 (aq) + NaCl (aq) --> AgCl (s) + NaNO3 (aq)

  24. Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) --> AgCl(s) + Na+ + NO3-(aq) Net Ionic Equations AgNO3(aq) + NaCl(aq) --> AgCl(s) + NaNO3(aq) If we re-write the above equation in terms of the species that actually exist in solution, the equation is: This is the complete IONIC equation In writing the complete ionic equation we see that Na+ and NO3-exist in the same form on both sides of the equation, whereas the Ag+ and Cl- have reacted to form solid AgCl.

  25. Both Na+ and NO3- are called SPECTATOR IONS since they are present in solution but do not participate directly in the chemical reaction. They “provide” the Ag+ and Cl- ions which reacted, but they themselves do not directly participate in the chemical reaction. Hence, the complete ionic equation can be re-written to show only those species which are directly involved in the chemical reaction Ag+(aq) + Cl-(aq) --> AgCl(s) This is the NET IONIC equation.

  26. Note: Net ionic equations, as with any chemical equation must be balanced, in terms of mass and charge. So for a reaction between Pb2+(aq) and I-(aq) to form PbI2(s) must be written as: Pb2+(aq) + 2 I-(aq) --> PbI2(s)

  27. aqueous Calcium chloride: CaCl2(aq) aqueous sodium carbonate: Na2CO3(aq) Next, determine what the products of the reaction will be and which product is the precipitate. Example: Write the net ionic equation for the precipitation reaction that occurs when aqueous solutions of calcium chloride and sodium carbonate are mixed. First write the chemical formulas of the reactants The products of this reaction are: NaCl and CaCO3. From the solubility table, determine if NaCl and CaCO3 are soluble in water or not. Write the net ionic equation - make sure it is balanced in terms of mass and charge.

  28. Hence, the equation for the reaction is: CaCl2(aq) + Na2CO3(aq) --> NaCl(aq) + CaCO3 (s) CaCl2(aq) + Na2CO3(aq) --> 2 NaCl(aq) + CaCO3 (s) Ca2+(aq)+ 2Cl-(aq) + 2 Na+ + CO32-(aq) --> 2 Na+(aq) + 2Cl-(aq) + CaCO3 (s) Canceling the spectator ions, gives the net ionic equation Ca2+(aq) +CO32-(aq) --> CaCO3 (s) Check that the ionic equation is balanced in terms of mass and charge

  29. Acid-Base Reactions Acids and bases are probably one of the more commonly encountered compounds Acids are found in fruit (citric acid), vinegar (acetic acid), in our stomachs (HCl), and in acid rain (H2SO4, HNO3). Examples of bases - aqueous solutions of ammonia (used in household cleanser); antacids like “Tums” and “Rolaids” Acids and bases are also electrolytes they dissociate in water to form ions.

  30. Arrhenius Acids and Bases Arrhenius acids are compounds which in aqueous solution dissociate to form H+ ions. (The H+ ion is also referred to as a proton since a H+ ion does not have an electron and the charge is due to the single proton in the nucleus) Hence acids are H+ donors or proton DONORS.

  31. Examples: HCl(aq) --> H+(aq) + Cl-(aq) HNO3(aq) --> H+(aq) + NO3-(aq) Acids like HCl and HNO3 are called MONOPROTIC acids since every molecule of HCl or HNO3 produces one H+ Acids like HCl and HNO3 completely dissociate in water These acids are are STRONG ACIDS, and hence also strong electrolytes

  32. HSO4-(aq) H+(aq) + SO42-(aq) The fact not all the HSO4- ions in solution dissociate is denoted by the arrows pointing in both direction H2SO4(aq) --> H+(aq) + HSO4-(aq) The HSO4- formed can dissociate further producing a H+. However, HSO4- is not as strong an electrolyte as H2SO4 and not all the HSO4- ions dissociate. Dissociation of HSO4-is “incomplete”

  33. H2SO4 is a strong electrolyte since it completely dissociates in water and hence is a strong acid. HSO4-is a weak electrolyte since it does not completely dissociate in water and hence a weak acid. An aqueous solution of H2SO4 contains H+, HSO4-and SO42- ions. H2SO4 is called a DIPROTIC acid, since each molecule of H2SO4 can produce up to 2 H+ ions. Polyprotic acids - produce more that 2 H+ /molecule of acid

  34. NH3(g) + H2O(l) NH4+(aq) + OH-(aq) NaOH(s) --> Na+(aq) + OH-(aq) Compounds that do not contain OH- can still be bases as long as when the dissolve in water the chemical reaction that results produces OH- ions. For example, if NH3 is dissolved in water, it can react with the water to produce OH- Arrhenius Bases An Arrhenius base is a compound that when dissolved in water dissociates to produce an OH- (hydroxide) ion.

  35. In this example, note that the reaction is not complete as indicated by the double arrows. NH3 is a weak electrolyte and hence a weak base. Also, note that NH3 received a H+ from water, and so in this reaction the water acts as an acid giving up an H+ to NH3.

  36. Water itself can dissociate as follows: H2O(l) H+(aq) + OH-(aq) producing both H+ and OH- ions. However the extent to which pure water dissociates is very,very small, and a very small number of water molecules dissociate. In fact, pure water is not considered to be an electrolyte and does not conduct electricity. Compounds, like water, that can exhibit both acidic and basic properties are called AMPHOTERIC. Other examples of amphoteric compounds are amino acids which exhibit both acidic and basic properties.

  37. Reactions between acids and bases When solutions of an acid and a base are mixed, a NEUTRALIZATION reaction occurs. The products of the neutralization reaction have neither acidic nor basic properties. Example: HCl(aq) + NaOH(aq) --> H2O(l) + NaCl(aq) In general: acid + base --> salt + water

  38. The complete ionic equation for the reaction between HCl and NaOH is H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) --> H2O(l) + Na+(aq) + Cl-(aq) Therefore, the net ionic equation is H+(aq) + OH-(aq) --> H2O(l)

  39. Problem Milk of magnesia is essentially Mg(OH)2 and is insoluble in water. Adding HCl to a suspension of milk of magnesia dissolves it leaving behind a clear solution. Write the overall equation and the net ionic equation. Overall equation Mg(OH)2(s) + 2 HCl(aq) --> MgCl2(aq) + 2 H2O(l) Complete ionic equation Mg(OH)2(s) + 2H+(aq) + 2Cl-(aq) --> Mg2+(aq) + 2Cl- (aq) + 2H2O(l) Net ionic equation Mg(OH)2(s) + 2H+(aq) -> Mg2+(aq) + 2H2O(l)

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