Types of Chemical Equations PDF

Types of Chemical Equations Water, on the other hand, is an extremely weak electrolyte and exists almost entirely in the form of H2O molecules. So, when a solution of sodium hydroxide is mixed with a solution of hydrochloric acid the expression showing the ionic and molecular species invo...

Types of Chemical Equations Water, on the other hand, is an extremely weak electrolyte and exists almost entirely in the form of H2O molecules. So, when a solution of sodium hydroxide is mixed with a solution of hydrochloric acid the expression showing the ionic and molecular species involved would be: Na+ + OH- + H+ + Cl- → Na+ + Cl- + H2O It is clear that both Na+ and Cl- do not altered in the reaction. Therefore, the actual change which takes place is expressed by Net Ionic Equation : H+ + OH- → H2O 115 How to write Net Ionic Equation? 1) Write the balanced molecular equation. 2) Write the ionic equation showing the strong electrolytes. 3) Determine if there is any precipitate from the solubility rules. 4) Cancel the similar ions on both sides of the ionic equation. 116 How to write Net Ionic Equation? Example 1: Write the net ionic equation of the reaction of sodium hydroxide with hydrochloric acid. The molecular equation: NaOH + HCI → NaCl + H2O The ionic equation showing the strong electrolytes: Na+ + OH- + H+ + Cl- → Na+ + Cl- + H2O No precipitate is formed; so cancel similar ions: Na+ + OH- + H+ + Cl- → Na+ + Cl- + H2O Net ionic equation will be: H+ + OH- → H2O 117 How to write Net Ionic Equation? Example 2: Write the net ionic equation of the reaction of a solution of silver nitrate with hydrochloric acid. The molecular equation: AgNO3 + HCI → AgCl  + HNO3 The ionic equation showing the strong electrolytes: Ag+ + NO3- + H+ + Cl- → AgCl + NO3- + H+ Silver chloride (white precipitate) is formed; then cancel the similar ions: Ag+ + NO3- + H+ + Cl- → AgCl  + NO3- + H+ Net ionic equation will be: Ag+ + Cl- → AgCl  118 Chemical Bonds Compounds are made of atoms held together by bonds. Chemical bonds are forces of attraction between atoms. The bonding attraction comes from attractions between protons and electrons of bonded atoms. Bonds can form between atoms of the same element, or between atoms of different elements. Chemical bonds form because they lower the potential energy between the charged particles that compose atoms. 119 Types of Chemical Bonds Chemical bonds can be classified into three types, depending on the types of atoms involved in the bonding: Ionic bond Covalent bond Intramolecular Force Metallic bond 120 Types of Chemical Bonds: The Ionic Bond Ionic bond: results when electrons have been transferred between atoms, resulting in oppositely charged ions that attract each other. Generally formed when metal atoms bond to nonmetal atoms. Method: electron transfer. 121 Types of Chemical Bonds: The Covalent Bond Covalent bond: results when two atoms share some of their electrons: Generally formed when nonmetal atoms bond together Shared electrons hold the atoms together by attracting nuclei of both atoms. Method: electron sharing Multiple Covalent Bonds: ✓ Single covalent bond: A covalent bond formed by sharing one electron pair (2eˉ). Represented by a single line: H−H ✓ Double covalent bond: formed by sharing two electron pairs (4eˉ). Represented by a double line: O=O ✓ Triple covalent bond: formed by sharing three electron pairs (6eˉ). Represented by a triple line: N≡N 122 Types of Chemical Bonds: The Coordinate Bond Coordinate bond (also called a dative covalent bond): is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom. Coordinate bond is a sharing of lone pair of electrons from one atom called donor (Lewis base) to another atom called acceptor (Lewis acid). Lewis acid: electron pair acceptor e.g. H+, AlCl3, FeBr3, BF3. Lewis base: electron pair donor e.g. compounds containing heteroatoms (O, S, N) e.g. NH3, H2O. 123 Representing Valence Electrons with Dots (Lewis Structures) Lewis Structures: simple diagrams to visualize the number of valence electrons in atoms of main-group elements by dots. The dots are placed around the element’s symbol with a maximum of two dots per side. Each dot represents one valence electron. ✓ Remember: the number of valence electrons for main group element is equal to the group number of the element (except for helium, which is in group 8A but has only two valence electrons). ✓ Note: While the exact location of dots is not critical, here we first place dots singly before pairing (except for helium which always has two paired dots) 124 Representing Valence Electrons with Dots (Lewis Structures) The electron configuration of Oxygen is as follows: Its Lewis structure is as follows: 125 Representing Valence Electrons with Dots (Lewis Structures) Lewis structure for all period 2 elements: Practice: Draw the Lewis dot structure of a phosphorus atom. Solution: Since phosphorus is in Group 5A in the periodic table, it has 5 valence electrons. Represent these as five dots surrounding the symbol for phosphorus: 126 Lewis Structures: For Covalent Bonding Hydrogen and oxygen have the following Lewis structures: In water, hydrogen and oxygen share their valence electrons so that each hydrogen atom gets a duet and the oxygen atom gets an octet. 127 Lewis Theory Predicts That Hydrogen Should Exist as H2 The individual Hydrogen atoms has the following Lewis structure: When two hydrogen atoms share their valence electrons, they each get a duet, a stable configuration for hydrogen. Lewis theory predicts that elemental hydrogen exists as a diatomic molecule (H2). 128 Lewis Structures: Double and Triple Covalent Bonds Oxygen exists as a diatomic molecule (O2): Nitrogen exists as a diatomic molecule (N2): 129 Assessment 1. Write the Lewis structure for each atom or ion: a. Al b. sodium ion c. magnesium ion d. chloride ion 2. Use Lewis structures to explain why each element occurs as diatomic molecules: a. hydrogen b. bromine c. oxygen d. nitrogen 3. Write the Lewis structure for each compound: a. PH3 b. SCl2 c. HI d. CH4 e. NaF f. CaO g. SrBr2 h. K2O 4. Determine whether a bond between each pair of atoms would be nonpolar covalent, polar covalent, or ionic. a. Br & Br b. C & Cl c. Mg & I d. Sr & O 5. Order these compounds in order of increasing carbon–carbon bond strength and in order of decreasing carbon–carbon bond length: HC≡CH , H2C═CH2 , H3C─CH3 130  B- Molecules, Compounds and Chemical Bonds 4- Intermolecular Forces and Bond Polarity Intermolecular Forces Intermolecular forces: are the attractive forces that exist between molecules. In contrast to intermolecular forces, intramolecular forces hold atoms together in a molecule. Generally, intermolecular forces are much weaker than intramolecular forces. It usually requires much less energy to evaporate a liquid than to break the bonds in the molecules of the liquid. From the chemistry point of view: The strength of the intermolecular forces (IMF) determines whether a compound has a high or low melting point and boiling point, and thus if the compound is a solid, liquid, or gas at a given temperature. IMF influences the solubility of substances in various solvents. IMF affects the rate and outcome of chemical reactions by influencing how reactant molecules come together and interact. 132 Intermolecular Forces in Covalent Molecules There are three different types of intermolecular forces in covalent molecules, presented in order of increasing strength: London dispersion forces (also called van der Waals forces) Dipole–dipole interactions (also called van der Waals forces) Hydrogen bonding Intermolecular Forces: London Dispersion Forces London dispersion forces: are very weak interactions due to the momentary changes in electron density in a molecule. Temporary dipole All covalent compounds exhibit London dispersion forces. These intermolecular forces are the only intermolecular forces present in nonpolar compounds. The strength of these forces is related to the size of the molecule. ▪ The larger the molecule, the larger the attractive force between two molecules, and the stronger the intermolecular forces. 134 Intermolecular Forces: Dipole–dipole interactions Dipole–dipole interactions: are the attractive forces between the permanent dipoles of two polar molecules. For example, the carbon–oxygen bond in formaldehyde, H2C=O, is polar because oxygen is more electronegative than carbon. This polar bond gives formaldehyde a permanent dipole, making it a polar molecule. The dipoles in adjacent formaldehyde molecules can align so that the partial positive and partial negative charges are close to each other. These attractive forces due to permanent dipoles are much stronger than London dispersion forces. 135 Intermolecular Forces: Hydrogen Bonding Hydrogen bonding is a strong type of dipole-dipole interaction which occurs when a hydrogen atom bonded to O, N, or F, is electrostatically attracted to an O, N, or F atom in another molecule. Hydrogen bonding is only possible between two molecules that contain a hydrogen atom bonded to a very electronegative atom—that is, oxygen, nitrogen, or fluorine These forces are weaker than intramolecular bonds, but are much stronger than other intermolecular forces, causing these compounds to have high boiling points. Hydrogen bonds are the strongest of the three types of intermolecular forces 136 Intermolecular Forces: Types of Hydrogen Bonding Intermolecular hydrogen bond: Refers to reaction between two same or different molecules. Occurs when hydrogen locates between two electronegative groups (N, S, O). Intramolecular hydrogen bond: Refers to hydrogen bonds within the same molecule. Stronger than intermolecular hydrogen bonds. Higher boiling and melting points. Intramolecular hydrogen bonding of salicylaldehyde 137 Intermolecular Forces: Types of Hydrogen Bonding Examples of Intra- and Intermolecular Hydrogen Bonds Intermolecular Forces: Ion-dipole Forces Ion–dipole forces: are the electrostatic attractions between a charged ion and a dipole. They are common in solutions and play an important role in the dissolution of ionic compounds, like NaCl, in polar solvent such as water. The strength of ion–dipole interactions is directly proportional to: 1. the charge on the ion 2. the magnitude of the dipole of polar molecules. 139 Why are intermolecular forces important in pharmacy? IMF forces are of great importance in the field of pharmacy, as they play a crucial role in various aspects of pharmaceutical science and drug development. Here are some key ways in which IMF are important in pharmacy: 1- Drug solubility and formulation: The solubility of a drug molecule in certain solvent is largely determined by the IMF between the drug and the solvent. Understanding these forces is crucial for designing effective drug formulations. 2- Drug absorption and bioavailability: IMF influence the ability of a drug to be absorbed from the gastrointestinal tract into the bloodstream. For example, hydrogen bonding and ionic interactions can affect the permeability of drug molecules across biological membranes. The bioavailability of a drug, which is the fraction of the administered dose that reaches the systemic circulation, is also affected by intermolecular forces during absorption, distribution, and metabolism. Why are intermolecular forces important in pharmacy? 3- Drug stability and shelf life: IMF such as hydrogen bonding and van der Waals interactions, can stabilize the structure of drug molecules and influence their resistance to degradation. Understanding the IMF involved in drug-excipient and drug-drug interactions is crucial for developing stable pharmaceutical formulations with an adequate shelf life. 4- Drug delivery and targeting: IMF play a role in the design of drug delivery systems, such as liposomes, nanoparticles, and polymeric carriers, which aim to improve the targeting and controlled release of drugs. 5- Protein-drug interactions: Many drugs exert their therapeutic effects by interacting with specific proteins, such as enzymes, receptors, or transport proteins. The strength and nature of these protein-drug interactions are determined by IMF. Intermolecular Forces: Problems 1- What types of intermolecular forces are present in each compound: (a) HCl; (b) C 2H6 (ethane); (c) NH3? N.B. London dispersion forces are present in all covalent compounds. Dipole–dipole interactions are present only in polar compounds with a permanent dipole. Hydrogen bonding occurs only in compounds that contain an O – H, N – H, or H – F bond. a. HCl has London forces like all covalent compounds. HCl has a polar bond, so it exhibits dipole–dipole interactions. HCl has no H atom on an O, N, or F, so it has no intermolecular hydrogen bonding. b. C2H6 is a nonpolar molecule since it has only nonpolar C – C and C – H bonds. Thus, it exhibits only London forces. c. NH3 has London forces like all covalent compounds. NH3 has a net dipole from its three polar bonds, so it exhibits dipole–dipole interactions. NH3 has a H atom bonded to N, so it exhibits intermolecular hydrogen bonding.142 Intermolecular Forces: Problems 2. What types of intermolecular forces are present in each molecule? a. Cl2 b. HCN c. HF d. CH 3Cl e. H2 3. Which of the compounds in each pair has stronger intermolecular forces? a. CO2 or H2O b. CO2 or HBr c. HBr or H2O d. CH4 or C2H6B 4. Determine which compound can form inter or intramolecular hydrogen bonding: a. 7.30 e e 143 Electronegativity and Bond Polarity Electronegativity (EN): is the ability of an atom (in a molecule) to attract the bond electrons to itself. ✓ is higher for nonmetals; and lower for metals ✓ The greater the difference in electronegativity (ΔEN), the more polar the bond. δ+ and δ- in polar molecular compounds represent the partial positive and negative charges, to differentiate them from the full charge (+ or -) on ions in ionic compounds. 144 Electronegativity Values for Elements (Unitless) 145 Electronegativity and Bond Types 146 Electronegativity and Bond Types: Examples - Example: Based on the of values of electronegativity (EN) of elements, which bond is more polar: (B-Cl) or (C-Cl)? Answer: - The ΔEN of Cl and B = 3.0 - 2.0 = 1.0 - The ΔEN of Cl and C = 3.0 - 2.5 = 0.5 ✓ Hence, the B-Cl bond is more polar. - Exercise 1: Which of the following bonds is the most polar? (a) H−F (b) Se−F (c) N−P (d) Ga−Cl - Exercise 2: Predict the type of each bond (use the table of EN values): (a) H−Br (b) O−O (c) H−O (d) S−O 147 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 than 112 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 Electrolytes Solutions in Pharmacy Electrolyte Use Magnesium sulfate A drug used to treat convulsions during pregnancy, nephritis in children, magnesium deficiency, and tetany. + Ammonium chloride Expectorant in cough syrups. The ammonium ion (NH4 ) in the body plays an important role in the maintenance of acid-base balance. Sodium chloride An ingredient found in a variety of nutritional products as a source of electrolytes and water. Sodium acetate A compound used for electrolyte replenishment and total parenteral nutrition (TPN) therapy. Magnesium chloride An ionic compound and source of magnesium used for electrolyte replenishment and conditions associated with magnesium deficiencies. Potassium chloride A potassium salt used to treat hypokalemia. Calcium chloride An ionic compound used for the treatment of hypocalcemia and hyperkalemia, and as an antidote to magnesium intoxication due to overdosage of magnesium sulfate. 20 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 H2O + CO2 ↑ 2H+ + SO32- H2SO3 H2O + SO2 ↑ 2H+ + S2O32- H2S2O3 H2O + S  + SO2 ↑ 2H+ + 2NO2- 2HNO2 H2O + 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 H2O 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