General Chemistry for Pharmaceutical Sciences PHARM-101 PDF

 General Chemistry for Pharmaceutical Sciences PHARM-101 Presented by Dr. Azza H. Rageh Associate Professor of Pharmaceutical Analytical Chemistry  D- Stoichiometry, Solution Concentration and Chemical Reactions 1- Reaction Stoichiometry, Solution Concentration and Ty...

 General Chemistry for Pharmaceutical Sciences PHARM-101 Presented by Dr. Azza H. Rageh Associate Professor of Pharmaceutical Analytical Chemistry  D- Stoichiometry, Solution Concentration and Chemical Reactions 1- Reaction Stoichiometry, Solution Concentration and Types of Aqueous Solutions What is meant by Stoichiometry  Stoichiometry: calculations of the quantities of reactants and products in a chemical reaction.  Stoichiometry allows us to predict the amounts of products that will form in a chemical reaction based on the amount of the reactants.  Stoichiometry also allows us to determine the amount of reactants necessary to form a given amount of product. 4 Reaction Stoichiometry  The coefficients in a balanced chemical equation specify the relative amounts in moles of each of the substances involved in the reaction:  How Much CO2 is Produced?  Example: 2 C8H18(l) + 25 O2(g) 16 CO2(g) + 18 H2O(g) 2 molecules of C8H18 react with 25 molecules of O2 to form 16 molecules of CO2 and 18 molecules of H2O. Or: 2 moles of C8H18 react with 25 moles of O2 to form 16 moles of CO2 and 18 moles of H2O. 2 mol C8H18 : 25 mol O2 : 16 mol CO2 : 18 mol H2O 5 Reaction Stoichiometry From the balanced equation of the combustion of octane: 2 C8H18(l) + 25 O2(g) 16 CO2(g) + 18 H2O(g) we can write the following stoichiometric ratio: 2 moles C8H18(l) : 16 moles CO2 (This ratio is called: The Conversion Factor) Suppose that we burn 22 moles of C8H18: the amount of CO2 produced can be calculated using the conversion factor, as follows: 22 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝐶8𝐻18 × 16 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝐶𝑂2 = 176 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝐶𝑂2 2 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝐶8𝐻18 6 Concentration of Solutions What Is a “Solution”? Solution: A homogenous mixture of two or more substances: - Solvent: material present in largest amount. - Solute: all other materials present. - Example: Consider sugar dissolved in water: - Water is the solvent. - Sugar is the solute. Concentrated solution has a relatively large proportion of solute to solvent. Dilute solution has a relatively smaller proportion of solute to solvent. 7 Concentration of Solutions  Concentration: is the amount of solute present in the solution. Concentration units: Name Units Symbol % Weight Gram solute/100 g solution %, w/w % Volume Milliliter solute/100 mL solution %, v/v % Weight per volume Gram solute/100 mL solution %, w/v Parts per million Gram solute/106 g solution ppm Parts per billion Gram solute/109 g solution ppb Gram molecular weight of solute Molarity M (Moles)/liter solution Gram formula weight of Formality F solute/liter solution Gram equivalent weight of solute Normality N /liter solution Molality Moles solute/1000 g solution m 8 Concentration of Solutions: Molarity  Molarity: is a method to express the concentration. It shows the relationship between the moles of solute and liters of solution.  is the No of gram molecular weight (moles) of solute per one litre of solution. No. of moles = weight (g)/molecular weight So, Molarity (M) = weight (g)/molecular weight × volume (L)  Unit of molarity (M) = moles of solute / liter of solution M = mol/L = mol.L-1 = molar 9 The States of Matter Concentration of Solutions: Molarity Example 1: Find the molarity of a solution that has 25.5 g KBr dissolved in 1.75 L of solution Given: 25.5 g KBr, 1.75 L solution Find: molarity, M Plan: g KBr mol KBr M L sol’n Relationships: 1 mol KBr = 119.00 g, M = moles/L Solution: Check: because most solutions are between 0 and 18 M, the answer makes sense 10 Concentration of Solutions: Molarity Example 2: How many litres of 0.125 M NaOH solution would contain 0.255 mol NaOH? Given: 0.125 M NaOH, 0.255 mol NaOH Find: liters, L Plan: mol NaOH L sol’n Relationships: 0.125 mol NaOH = 1 L solution Solution: Check: because each L has only 0.125 mol NaOH, it makes sense that 0.255 mol should require a little more than112 L  Concentration of Solutions: Molarity Example 3: Preparing 1 L of a 1 M NaCl Solution 12 Types of aqueous solution and Solubility Consider two familiar aqueous solutions: salt water and sugar water: – Salt water is a homogeneous mixture of NaCl and H2O. – Sugar water is a homogeneous mixture of C12H22O11 and H2O. As you stir either of these two substances into the water, it seems to disappear. – How do solids such as salt and sugar dissolve in water? 13 What Happens When a Solute Dissolves? There are attractive forces between the solute particles holding them together. There are also attractive forces between the solvent molecules.  When we mix the solute with the solvent, there are attractive forces between the solute particles and the solvent molecules.  If the attractions between solute and solvent are strong enough, the solute will dissolve. 14 Dissolving of Sodium Chloride in Water  Each ion is attracted to the surrounding water molecules and pulled off and away from the crystal.  Compounds such as salt that dissociate into ions when dissolved in water are called electrolytes, and the resulting solutions are able to conduct electricity. 15 Dissolving of Sugar in Water  Table sugar (sucrose, C12H22O11) molecules homogeneously mixed with water molecules (H2O).  Compounds such as sugar that don’t dissociate into ions when dissolved in water are called nonelectrolytes, and the resulting solutions do not conduct electricity. 16 Electrolytes and Nonelectrolytes Substances that dissolve in water to form solutions that conduct electricity are called Electrolytes  Solution of salt (an electrolyte)  Solution of sugar (a nonelectrolyte) 17 Electrolytes and Nonelectrolytes  Strong Electrolytes:  substances that completely ionize when dissolve in water.  They can conduct electrical current strongly.  Important Examples: Soluble ionic salts (e.g. NaCl & MgBr2 …), strong acids (e.g. HCl & HNO3) and strong bases (e.g. NaOH & Mg(OH)2).  Weak Electrolytes:  Include substances that partially ionize when dissolve in water.  They can conduct electrical current weakly.  Important Examples: weak acids (e.g. HF & CH3COOH) and weak bases (e.g. NH4OH).  Nonelectrolytes:  Include substances that do not ionize when dissolve in water.  They don’t conduct electrical current.  Important Examples: molecular substances (e.g. sugar & alcohol). 18 Electrolytes and Nonelectrolytes: A Summary Complete Ionizing in Partial Ionizing in water No Ionizing in water (no water (full dissociation) (partial dissociation) dissociation) Examples: ionic salts, Examples: molecular Examples: weak acids strong acids & strong (covalent) compounds as bases & weak bases sugars & alcohols 19 Assessment All of the following compounds are soluble in water, indicate which of them is expected to produce strong, weak or non- electrolyte solution? 1. CsCl(aq) 2. CH3OH(aq) 3. Ca(NO2)2(aq) 4. C6H12O6(aq) 5.Acetic acid, vinegar, CH3COOH(aq) (weak acid) 6. HCl(aq) (strong acid) 7. NaOH(aq) (strong base) 8. KOH(aq) (strong base) 9. HF(aq) (weak acid) 10.NH4OH(aq) (weak base) 21  D- Stoichiometry, Solution Concentration and Chemical Reactions 2- Basic types of Chemical Reactions and Reactions involving Oxidation-Reduction Basic Types of Chemical Reactions 1- Synthesis Reaction 2- Decomposition Reaction 3- Replacement Reactions 23 1- Synthesis Reactions  In a synthesis reaction, two or more reactants combine to yield one product.  These reactions are expressed in the general form: A + B ⟶ AB  An example of a synthesis reaction is the combination of iron and sulfur to form iron(II) sulfide: 8 Fe + S8 ⟶ 8FeS  Another example is simple hydrogen gas combined with simple oxygen gas to produce a more complex substance, such as water. 2 H2 + O2 ⟶ 2 H2O 24 2- Decomposition Reactions  A decomposition reaction occurs when a more complex substance breaks down into its simpler parts.  It is the opposite of a synthesis reaction, and can be written as: AB ⟶ A + B  An example of a decomposition reaction is the electrolysis of water to make oxygen and hydrogen gas: 2 H2O ⟶ 2 H2 + O2 25 3- Replacement Reactions  There are two types of the replacement reactions: A. Single replacement reaction B. Double replacement reaction A. Single replacement reaction  In this reaction, a single uncombined element replaces another element in a compound  These reactions come in the general form of: A + BC ⟶ AC + B  For example; magnesium replaces hydrogen in water to make magnesium hydroxide and hydrogen gas Mg + 2 H2O ⟶ Mg(OH)2 + H2↑ 26 3- Replacement Reactions B. Double replacement reaction  In a double replacement reaction, the anions and cations of two compounds switch places and form two entirely different compounds.  These reactions are in the general form: AB + CD ⟶ AD + CB  For example; the reaction of lead(II) nitrate with potassium iodide to form lead(II) iodide and potassium nitrate: Pb(NO3)2 + 2 KI ⟶ PbI2 ↓ + 2 KNO3 27 Types of Chemical Reactions: A Summary Representation of four basic types of chemical reactions: 1- Synthesis, 2- Decomposition, 3- Single replacement and 4- Double replacement. 28 Chemical Reactions between Ions  Combination of ions occurs through the formation of any of the following: Water 1 Weak electrolyte 2 Precipitate 3 Gas 4 Complex ion 5 29 1- Formation of Water  When metallic hydroxide, including ammonium hydroxide, is mixed with an acid , water is formed. Ex. 1 NaOH + HCl NaCl + H2O Ex. 2 NH4OH + HCl NH4Cl + H2O  These examples represent Acid-Base reactions (reactions between strong/weak base with strong/weak acid to form salt and water). 30 2- Formation of Weak Electrolyte  When a solution of a strong acid is mixed with a solution of a salt containing the anion of a weak acid, the weak acid is formed. HCl + CH3COONa NaCl + CH3COOH (weak acid)  When a solution of a strong base is mixed with a solution of ammonium salt, the weak base (ammonium hydroxide) is formed: NH4Cl + NaOH NH4OH (weak base) + NaCl 31 3- Formation of a Precipitate  Precipitation is the formation of a solid in a solution during a chemical reaction.  It usually takes place when the concentration of dissolved ions exceeds the solubility limit and forms an insoluble salt.  This process can be assisted by adding a precipitating agent or by removal of the solvent.  For example; the reaction between silver nitrate (precipitating agent) and sodium chloride to form silver chloride precipitate AgNO3 + NaCl AgCl + NaNO3 ppt 32 3- Formation of a Precipitate Another example:  If solutions of ferric chloride and sodium hydroxide are mixed, the following ions are present: Fe3+, Cl-, Na+ and OH-.  Ferric hydroxide is formed which is insoluble in water  The following reaction occurs: FeCl3+ 3 NaOH Fe(OH)3 + 3NaCl ppt 33  4- Formation of a Gas  The combination of ions may result in the evolution of a gas for two reasons because either the product is gaseous, or the product is unstable and decomposes to form a gas.  Examples of the former are: Ex. 1 2H+ + S2- H2S ↑ hydrogen sulphide gas Ex. 2 H+ + CN- HCN ↑ Hydrogen cyanide  Unstable acids formed by the combination of ions are H2CO3, H2SO3, H2S2O3, and HNO2: 2H+ + CO32- H2CO3 H 2O + CO2 ↑ 2H+ + SO32- H2SO3 H 2O + SO2 ↑ 2H+ + S2O32- H2S2O3 H 2O + S  + SO2 ↑ 2H+ + 2NO2- 2HNO2 H 2O + NO ↑ + NO2 ↑ H2CO3 carbonic acid H2SO3 sulfurous acid H2S2O3 hydrogen thiosulphate HNO2 nitrous acid 34 5- Formation of Complex Ions  Complexation reaction is the reaction between Lewis acid (electron acceptor; Metal) and Lewis base (electron donor; Ligand).  In complexation reactions, several ligands form coordinate bonds with a metal atom to form a complex.  This is achieved by donating lone pairs of electrons from the ligand (L) to the metal atom (M) to form a coordination complex (ML). L: + M ⟶ M L Coordinate bond 35 5- Formation of Complex Ions  Ligands are Lewis bases, They can be both anions and neutral molecules.  Anions that frequently form complexes are chloride (Cl-), bromide (Br-), iodide (I-), fluoride (F-), cyanide (CN-), thiocyanate (SCN-), thiosulphate (S2O32-) and oxalate (C2O42-).  Neutral molecules that frequently form complexes such as; carbon monoxide (CO), ammonia (NH3) and water.  Many cations act as metals such as (Ca2+, Mg2+, Fe2+, Fe3+, Pb2+, Cu2+, Zn2+, Al3+...etc).  The only common cations that do not usually form complexes are sodium (Na+), potassium (K+), ammonium (NH4+). Cu2++ 4 ¨NH3 [Cu(NH3)4]2+ 36 Reactions involving Oxidation-Reduction (Redox) Oxidation – reduction reactions or redox reactions are reactions in which electrons are transferred from one reactant to the other. - Oxidation: is the loss of electrons. - Reduction: is the gain of electrons.  Based on these definitions, redox reactions do not need to involve oxygen.  One cannot occur without the other.  Example on Redox Reactions:  In this reaction, a metal (which has a tendency to lose electrons) reacts with a nonmetal (which has a tendency to gain electrons). In other words, metal atoms lose electrons to nonmetal atoms. 37 Reactions involving Oxidation-Reduction (Redox) Other common redox reactions: 38 Oxidizing Agent & Reducing Agent  Oxidizing Agent (Oxidant):  A Substance that oxidizes something else. The oxidizing agent itself is reduced in the same reaction:  Reducing Agent (Reductant)  A Substance that reduces something else. The reducing agent itself is oxidized in the same reaction: 39 Assessment 1- For the following reactants, what are the products of the double displacement reaction? FeCl3(aq) + Ba(OH)2 ​(aq)→ Choose 1 answer: A- no reaction B- Fe(OH)2 ​(s) + BaCl (aq) C- FeBa (s)+ HOCl (aq) D- Fe(OH)3(s) + BaCl2 ​(aq) 2- What type of reaction is the above reaction? Choose all answers that apply: A- Oxidation-reduction reaction B- Double replacement reaction C- Neutralization reaction D- Precipitation reaction 40 Assessment 1- What is the type of the following reactions: BaCl2 + MgSO4 ⟶ BaSO4 + MgCl2 2 H2 + O2 ⟶ 2 H2 O NaOH + CH3COOH ⟶ CH3COONa + H2O AgNO3 + Kl AgI + KNO3 2 NaBr + Cl2 2 NaCl + Br2 2 Na + Cl2 2 NaCl H2CO3 H2O + CO2 2- Calculate the oxidation number of: Manganese in KMnO4 , MnSO4 , MnO2 Chromium in K2CrO4 , K2Cr2O7 Oxygen in H2O , O2 , H2O2 3- What is a decomposition reaction? What is the general equation for this reaction and give an example? 4- Compare between: Lewis acid and Lewis base Oxidant and Reductant 41 Assessment 42