Lesson Notes By Weeks and Term v3 - Senior Secondary 2
Ionic Theory
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Ionic Theory
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Subject: Chemistry
Class: Senior Secondary 2
Term: 1st Term
Week: 5
Theme: Chemistry And Industry
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These ions are very stable and difficult to oxidize. If OH− ions are present (from water), OH− will be discharged in preference to these complex ions. Halides (Cl−, Br−, I−) and Hydroxide (OH−): In dilute solutions, OH− is preferentially discharged over halides. In concentrated solutions, the higher concentration of halide ions can override the series, leading to their preferential discharge over OH−.
General Rule for Anions: The anion that is lower in the electrochemical series (or less electronegative, more easily oxidized) is preferentially discharged at the anode. 2.
6. Factors Affecting the Preferential Discharge of Ions While the electrochemical series provides a general guideline, several factors can influence which ion is discharged preferentially:
1. Position of the Ion in the Electrochemical Series (Nature of the Ion): This is the primary factor. As explained above, cations lower in the series and anions lower in the series are generally preferred for discharge.
Example (Cathode): In a solution containing Na+ and Cu2+, Cu2+ (below H+) will be discharged in preference to Na+ (above H+).
Example (Anode): In a dilute solution containing SO42− and OH−, OH− will be discharged in preference to SO42−.
2. Concentration of Ions: For some ions, particularly the halide ions (Cl−, Br−, I−) and hydroxide ions (OH−), a high concentration of a less easily discharged ion can override its position in the electrochemical series.
Example: In the electrolysis of dilute sodium chloride solution, OH− ions (from water) are discharged at the anode in preference to Cl− ions, producing oxygen gas. 4OH−(aq) → 2H2O(l) + O2(g) + 4e− However, in the electrolysis of concentrated sodium chloride solution (brine), Cl− ions are in much higher concentration than OH− ions, so Cl− ions are preferentially discharged, producing chlorine gas. 2Cl−(aq) → Cl2(g) + 2e− This effect is primarily observed for anions at the anode.
3. Nature of the Electrodes (Active vs.
Inert Electrodes): Inert Electrodes (e.g., Platinum, Graphite): These electrodes do not react chemically with the electrolyte or the products of electrolysis. They merely provide a surface for electron transfer. The discharge products are solely determined by the electrochemical series and concentration. Active Electrodes (e.g., Copper, Silver, Nickel): These electrodes can participate in the reaction by dissolving (oxidizing) at the anode.
Example: In the electroplating of copper using a copper anode and copper(II) sulphate solution: At the anode (positive): The copper anode itself oxidizes and dissolves, replenishing Cu2+ ions in the solution, instead of OH− or SO42− from the electrolyte. Cu(s) → Cu2+(aq) + 2e− At the cathode (negative): Cu2+ ions from the solution are deposited. Cu2+(aq) + 2e− → Cu(s) This effect is significant at the anode. Summary of Discharge Rules for Aqueous Solutions (containing water): At the Cathode (Negative Electrode): If H+ (from an acid) and metal ions are present: H+ is discharged if the metal is above hydrogen in the series (K, Na, Ca, Mg, Al, Zn, Fe, Pb). The metal itself is discharged if it is below hydrogen (Cu, Ag). If only metal ions from water (H2O) are present: If the metal is above hydrogen, H2 gas is produced. If the metal is below hydrogen, the metal is deposited.
At the Anode (Positive Electrode): If the solution contains halides (Cl−, Br−, I−) and OH− (from water), and using inert electrodes: In dilute solutions, OH− is discharged, producing O2 gas. In concentrated solutions, halides are discharged, producing halogen gas (Cl2, Br2, I2). If the solution contains SO42− or NO3− and OH− (from water), OH− is always discharged (producing O2 gas) because SO42− and NO3− are very difficult to oxidize. If an active anode is used, the anode itself may dissolve (oxidize) preferentially. This section provides in-depth explanations of the core concepts related to Ionic Theory, ensuring teachers have sufficient content knowledge to deliver the lesson comprehensively. 2.
1. Electrovalent (Ionic) Compounds vs. Covalent Compounds Electrovalent (Ionic)
Compounds: Definition: Formed by the complete transfer of one or more electrons from a metal atom to a non-metal atom, resulting in the formation of oppositely charged ions which are held together by strong electrostatic forces of attraction.
Formation: Occurs between elements with a large difference in electronegativity (typically a metal and a non-metal). The metal atom loses electrons to form a positively charged cation, and the non-metal atom gains electrons to form a negatively charged anion.
Example: Formation of Sodium Chloride (NaCl) Na (2,8,1) → Na+ (2,8) + e− Cl (2,8,7) + e− → Cl− (2,8,8) Na+ + Cl− → NaCl Properties: Physical State: Crystalline solids at room temperature (due to strong electrostatic forces forming a lattice structure).
Melting and Boiling Points: High melting and boiling points (requires significant energy to overcome the strong electrostatic forces).
Solubility: Generally soluble in polar solvents like water (water molecules can surround and separate the ions).
Electrical Conductivity: Do not conduct electricity in the solid state (ions are fixed in the lattice). They conduct electricity when molten or in aqueous solution (ions become mobile and free to move).
Examples: Sodium chloride (common salt), Magnesium oxide, Potassium iodide, Calcium chloride.
Covalent Compounds: Definition: Formed by the sharing of one or more pairs of electrons between two non-metal atoms, allowing each atom to achieve a stable electron configuration (duplet or octet).
Formation: Occurs between elements with similar electronegativities (typically two non-metals). The shared electrons are attracted by both nuclei, holding the atoms together.
Example: Formation of Water (H2O) Oxygen shares electrons with two hydrogen atoms.
Example: Formation of Methane (CH4) Carbon shares electrons with four hydrogen atoms.
Properties: Physical State: Can be gases, liquids, or soft solids at room temperature (due to weak intermolecular forces).
Melting and Boiling Points: Generally low melting and boiling points (less energy required to overcome weak intermolecular forces).
Solubility: Soluble in non-polar solvents (like benzene, carbon tetrachloride); many simple covalent compounds are insoluble or sparingly soluble in water, unless they can form hydrogen bonds or react with water.
Electrical Conductivity: Do not conduct electricity in any state (solid, liquid, or gas) because they do not contain free ions or mobile electrons.
Examples: Water, methane, oxygen, carbon dioxide, ethanol, sugar (sucrose). 2.
2. Electrolytes vs.
Non-electrolytes Electrolytes: Definition: Substances that conduct electricity when molten or dissolved in a solvent (usually water), due to the presence of mobile ions. The conduction is accompanied by chemical decomposition.
Mechanism: When dissolved or melted, electrolytes dissociate or ionize into free, mobile positive ions (cations) and negative ions (anions). These ions migrate towards oppositely charged electrodes, carrying electric current.
Types of Electrolytes: All ionic compounds are electrolytes. Some covalent compounds (e.g., acids like HCl, H2SO4) also act as electrolytes because they react with water to produce ions (e.g., HCl + H2O → H3O+ + Cl−).
Examples: Acids (HCl, H2SO4, HNO3), Bases (NaOH, KOH, Ca(OH)2), Salts (NaCl, CuSO4, KNO3).
Non-electrolytes: Definition: Substances that do not conduct electricity in any state (solid, molten, or dissolved) because they do not produce free, mobile ions.
Mechanism: Non-electrolytes are typically covalent compounds that do not ionize or dissociate when dissolved. They remain as neutral molecules. *
Examples: Sugar (glucose, sucrose), ethanol, petrol, kerosene, distilled water (pure water is a very poor conductor due to minimal self-ionization). 2.
3. Movement of Ions in Solution When an electrolyte dissolves in water, its constituent ions separate (dissociation for ionic compounds) or form through reaction with water (ionization for certain covalent compounds like acids). For example, when NaCl dissolves in water: NaCl(s) + H2O(l) → Na+(aq) + Cl−(aq) If two electrodes connected to a power source are dipped into this solution, the positive electrode (anode) attracts the negative ions (anions), and the negative electrode (cathode) attracts the positive ions (cations). This directional movement of charged ions very poor conductor due to minimal self-ionization). 2.
3. Movement of Ions in Solution When an electrolyte dissolves in water, its constituent ions separate (dissociation for ionic compounds) or form through reaction with water (ionization for certain covalent compounds like acids). For example, when NaCl dissolves in water: NaCl(s) + H2O(l) → Na+(aq) + Cl−(aq) If two electrodes connected to a power source are dipped into this solution, the positive electrode (anode) attracts the negative ions (anions), and the negative electrode (cathode) attracts the positive ions (cations). This directional movement of charged ions constitutes the flow of electric current through the solution. At the electrodes, the ions lose or gain electrons, undergoing redox reactions (oxidation at anode, reduction at cathode), leading to their discharge. 2.
4. Strong and Weak Electrolytes The distinction between strong and weak electrolytes lies in their degree of ionization in solution.
Strong Electrolytes: Definition: Substances that ionize almost completely (or 100%) when dissolved in water. They produce a high concentration of mobile ions.
Conductivity: Exhibit high electrical conductivity.
Examples: Strong Acids: Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3).
Strong Bases: Sodium hydroxide (NaOH), Potassium hydroxide (KOH), Calcium hydroxide (Ca(OH)2).
Most Salts: Sodium chloride (NaCl), Potassium nitrate (KNO3), Copper(II) sulphate (CuSO4).
Weak Electrolytes: Definition: Substances that ionize only partially (a small percentage) when dissolved in water. They exist predominantly as un-ionized molecules in solution, producing a low concentration of mobile ions.
Conductivity: Exhibit low electrical conductivity.
Examples: Weak Acids: Ethanoic acid (acetic acid, CH3COOH), Carbonic acid (H2CO3), Citric acid.
Weak Bases: Aqueous ammonia (NH3(aq) or NH4OH), Magnesium hydroxide (Mg(OH)2).
Few Salts: Mercury(II) chloride (HgCl2) (though some teachers may categorize as non-electrolyte depending on context). 2.
5. Electrochemical Series (Reactivity Series of Ions) The electrochemical series, also known as the reactivity series of ions, is an arrangement of elements (specifically, their ions) in order of their decreasing tendency to accept electrons (get reduced) or discharge at an electrode during electrolysis. It is fundamentally derived from standard electrode potentials. For Cations (Positive Ions) at the Cathode (Reduction): Order of Discharge: Ions lower down in the series (more positive electrode potential) are more easily discharged (reduced) than those higher up.
Series (common cations): K+ > Na+ > Ca2+ > Mg2+ > Al3+ > Zn2+ > Fe2+ > Pb2+ > H+ > Cu2+ > Ag+ Explanation: Ions above H+ (K+ to Pb2+): These are ions of very reactive metals. They are difficult to reduce (prefer to remain as ions) and require more energy to be discharged. If H+ ions are present (from water), H+ will be discharged in preference to these metal ions. H+ (from acids) / H2O (from water): In aqueous solutions, if metal ions are above hydrogen in the series, hydrogen gas will be produced at the cathode. Ions below H+ (Cu2+ to Ag+): These are ions of less reactive metals. They are relatively easy to reduce and will be discharged preferentially to H+ ions from water.
General Rule for Cations: The cation that is lower in the electrochemical series (or less electropositive) is preferentially discharged at the cathode. For Anions (Negative Ions) at the Anode (Oxidation): Order of Discharge: Ions higher up in the series (more negative electrode potential, or less easily oxidized) are more difficult to discharge. Ions lower down (more easily oxidized) are preferentially discharged. Series (common anions, decreasing ease of discharge / increasing standard electrode potential): SO42− > NO3− > Cl− > Br− > I− > OH− Explanation: Sulphate (SO42−) and Nitrate (NO3−): These ions are very stable and difficult to oxidize. If OH− ions are present (from water), OH− will be discharged in preference to these complex ions. Halides (Cl−, Br−, I−) and Hydroxide (OH−): In dilute solutions, OH− is preferentially discharged over halides. In concentrated solutions, the higher concentration of halide ions can override the series, leading to their preferential discharge over OH−.
General Rule for Anions: The anion that is lower in the electrochemical series (or less electronegative, more easily oxidized) is preferentially discharged at the anode. *2.
6. This section outlines practical activities for both teachers and students to facilitate understanding and engagement. 3.
1. Teacher Activities: Introduction (10 minutes): Begin by asking students to recall different types of chemical bonds (ionic and covalent) and provide examples. Introduce the concept of electrical conductivity in substances and pose the question: "Why do some substances conduct electricity while others don't?" Briefly link this to the topic of Ionic Theory and its relevance to everyday life (e.g., batteries, rust). Concept Explanation & Demonstration (30 minutes): Differentiating Electrovalent/Covalent Compounds: Use models or diagrams to illustrate electron transfer (ionic) and electron sharing (covalent). Present a table comparing the key properties (state, M.P./B.P., solubility, conductivity) of ionic and covalent compounds. Provide diverse examples relevant to Nigeria (e.g., NaCl, water, sugar, kerosene). Electrolytes vs. Non-electrolytes (Practical Demonstration): Set up a simple conductivity tester (battery, bulb, connecting wires with carbon electrodes/crocodile clips).
Test various substances: Electrolytes:* Dilute HCl, NaOH solution, NaCl solution, CuSO4 solution (ensure these are available and safe to use).
Non-electrolytes:* Distilled water, sugar solution, ethanol, kerosene. Guide students to observe the bulb's brightness and record results. Emphasize that for non-electrolytes, the bulb does not light up. Strong vs.
Weak Electrolytes: Using the same conductivity tester, compare the brightness of the bulb for: Strong electrolytes:* Dilute HCl, NaOH, NaCl (bright bulb).
Weak electrolytes:* Ethanoic acid, aqueous ammonia (dim bulb, or no light for very weak ones depending on setup). Explain the difference in terms of complete vs. partial ionization.
Movement of Ions: Draw a simple electrolysis cell on the board (beaker, electrodes, power supply, electrolyte). Explain how ions dissociate/ionize and then migrate towards the oppositely charged electrodes. Use arrows to show movement. Introducing the Electrochemical Series (25 minutes): Present the electrochemical series for both cations and anions on the board. Explain what the series represents (tendency for discharge/reduction for cations, oxidation for anions). Use mnemonics or catchy phrases to help students remember the order (e.g., "Please Stop Calling Me A Zealous Indian, Try Learning How Copper Saves Gold" for metals relative to hydrogen, or simple memorization of anion series). Relate the position of ions to their ease of discharge. Provide hypothetical examples of solutions containing multiple ions and ask students to predict discharge. Factors Affecting Preferential Discharge (25 minutes): Systematically explain the three main factors: Position in Series: Reiterate its importance with examples.
Concentration: Use the classic example of dilute vs. concentrated NaCl electrolysis at the anode to illustrate how concentration can override the series for halides.
Nature of Electrode: Explain active vs. inert electrodes, giving the example of copper refining/plating. Draw diagrams to illustrate these scenarios. 3.
2. Student Activities: Classification Task (10 minutes): In groups, students classify a given list of compounds as electrovalent or covalent, and as electrolytes or non-electrolytes. Observation and Data Recording (15 minutes): During the teacher's conductivity demonstration, students observe, record their observations, and infer properties. Students can draw a diagram of the conductivity setup and label components.
Electrolysis Prediction (15 minutes): Given hypothetical solutions (e.g., aqueous CuSO4, molten NaCl, dilute H2SO4), students, in pairs, predict the ions present, their movement, and products at each electrode, justifying their answers using the electrochemical series.
Discussion and Questioning (Ongoing): Active participation in Q&A sessions, asking clarifying questions. Group discussions to analyze scenarios and arrive at conclusions.
Note-Taking: Students are expected to take comprehensive notes throughout the lesson.
Corrosion Prevention and Electroplating: Application: Ionic theory explains the process of corrosion (electrochemical oxidation of metals). It also underpins electroplating, a vital industrial process in Nigeria for protecting metals from corrosion and enhancing their appearance. For example, zinc plating (galvanization) protects iron roofing sheets and buckets from rusting, while chromium plating is used on vehicle parts (e.g., bumpers, rims) and household fixtures for aesthetic appeal and durability. The movement of ions (e.g., Zn2+, Cr3+) in the plating solution towards the cathode is central to this process.
Local Context: Many small and medium-scale industries in Nigeria engage in electroplating for furniture, jewelry, and automotive parts. Understanding ion discharge helps in optimizing these processes. Industrial Production of Chemicals (Chlor-Alkali Process): Application: The production of chlorine gas (Cl2), sodium hydroxide (NaOH), and hydrogen gas (H2) from the electrolysis of concentrated sodium chloride solution (brine) is a major industrial process. Chlorine is used for water purification, PVC production, and bleaching. Sodium hydroxide is crucial in soap making, textiles, and aluminium production. Hydrogen is used in ammonia synthesis.
Local Context: Industries in Nigeria that use these chemicals, such as water treatment plants (chlorine for disinfection) and soap factories (NaOH), rely on products derived from this process, which is a direct application of preferential ion discharge based on concentration.
Medical Applications: Rehydration and Intravenous Fluids: Application: Electrolytes (ions like Na+, K+, Cl−, HCO3−) are crucial for maintaining fluid balance, nerve function, and muscle contraction in the human body. Dehydration, common in Nigeria due to illnesses like cholera or strenuous activities, leads to electrolyte imbalance. Oral rehydration salts (ORS) and intravenous fluids (IV drips) work by providing a balanced solution of these vital ions, which are absorbed and move within the body's fluids according to electrochemical principles.
Local Context: ORS is a life-saving commodity readily available in Nigerian pharmacies and healthcare centres, particularly critical for children suffering from diarrhea. Medical personnel frequently administer IV fluids, emphasizing the importance of electrolyte balance.