Lesson Notes By Weeks and Term v3 - Senior Secondary 1
Acids, Bases and Salts
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Acids, Bases and Salts
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Subject: Chemistry
Class: Senior Secondary 1
Term: 3rd Term
Week: 2
Theme: Chemistry And Environment
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define acids, basesand salts; identify acids and bases; describe the natureof proton in an aqueoussolution; explain neutralizationreactions; explain how an acid-base in dicator works; use p H as a scaleand discuss the importance of the p Hvalue; identify and preparesalts (normal, acidic,basic); state properties of salts; state the rules of solubility of salts in water.
acid and Potassium hydroxide: H$_2$SO$_4$(aq) + 2KOH(aq) → K$_2$SO$_4$(aq) + 2H$_2$O(l)
3. Nitric acid and Calcium hydroxide: 2HNO$_3$(aq) + Ca(OH)$_2$(aq) → Ca(NO$_3$)$_2$(aq) + 2H$_2$O(l)
Importance of Neutralization: Agriculture: Reducing soil acidity by adding lime (calcium hydroxide/oxide).
Medicine: Treating indigestion with antacids (e.g., magnesium hydroxide, aluminium hydroxide) to neutralize excess stomach acid.
Industrial waste treatment: Neutralizing acidic or basic effluents before discharge.
Household: Using a mild acid (like vinegar) to neutralize a strong base spill, or vice versa. 2.5 How an Acid-Base Indicator Works Definition: An acid-base indicator is a weak organic acid or base that changes colour depending on the pH of the solution.
Mechanism: Indicators exist in two forms, an acidic form and a basic form, each having a distinct colour. These forms are in equilibrium.
For a weak acid indicator (HIn): HIn (acidic colour) ⇌ H$^+$ + In$^-$ (basic colour) In an acidic solution (high H$^+$ concentration), the equilibrium shifts to the left, favouring the acidic coloured form (HIn). In a basic solution (low H$^+$ concentration), H$^+$ ions are removed (by OH$^-$ ions), causing the equilibrium to shift to the right, favouring the basic coloured form (In$^-$). Common Indicators and their Colour Changes: Litmus: Red in acid, Blue in base, Purple in neutral.
Methyl Orange: Red in acid (pH 4.4).
Phenolphthalein: Colourless in acid (pH 10.0).
Universal Indicator: A mixture of several indicators that show a range of colours over a wide pH range (e.g., red for strong acid, orange for weak acid, green for neutral, blue for weak base, violet for strong base). 2.6 pH as a Scale and its Importance pH Scale: A numerical scale (ranging from 0 to 14) used to express the acidity or alkalinity of an aqueous solution. It measures the concentration of hydrogen ions (H$^+$) or hydronium ions (H$_3$O$^+$) in a solution.
Definition: pH = -log$_{10}$[H$^+$], where [H$^+$] is the molar concentration of hydrogen ions.
Interpretation: pH 7: Basic or Alkaline solution (lower H$^+$ concentration, higher OH$^-$ concentration).
Relationship between pH and pOH: pH + pOH = 14 (at 25°C).
Tools for Measuring pH: pH paper/Universal Indicator: Gives an approximate pH value by colour matching. pH meter: An electronic device that provides a precise numerical pH reading.
Importance of pH Value (Nigerian Context):
1. Agriculture: Soil pH significantly affects nutrient availability and crop growth. For example, cocoa, rice, and oil palm thrive in slightly acidic to neutral soils (pH 5.5-7.0), while cassava and yam tolerate slightly acidic soils. Farmers use pH testing to determine if liming (to increase pH) or adding organic matter (to decrease pH) is needed.
2. Biology and Health: Human blood pH is maintained within a narrow range (7.35-7.45). Deviations can lead to serious health issues. The pH of stomach acid (1-3) is crucial for digestion.
3. Food and Beverage Industry: pH control is essential for brewing burukutu or palm wine, fermenting garri or ogi, and in the production of soft drinks and bottled water to ensure taste, stability, and safety.
4. Water Treatment: The pH of drinking water is monitored and adjusted (usually to slightly alkaline, pH 7-8.5) to prevent pipe corrosion and ensure the effectiveness of disinfectants.
5. Environmental Monitoring: pH is a key indicator of water quality in rivers and lakes. Acid rain, a concern in some industrial areas, lowers the pH of water bodies, affecting aquatic life like fish in Nigerian rivers.
6. Industrial Processes: Many chemical reactions, including those in dyeing textiles (e.g., adire production), soap making, and pharmaceutical manufacturing, require specific pH conditions for optimal results. 2.7 Identification and Preparation of Salts (Normal, Acidic, Basic)
Types of Salts:
1. Normal Salts: Formed when all the replaceable hydrogen ions of an acid are completely replaced by metal ions or ammonium ions. These salts do not contain replaceable H$^+$ or OH$^-$ ions.
Examples: NaCl, Na$_2$SO$_4$, KNO$_3$, (NH$_4$)$_2$CO$_3$.
2. Acidic Salts: Formed when only part of the replaceable hydrogen ions of a polyprotic acid (an acid with more than one replaceable hydrogen) are replaced by metal ions or ammonium ions. These salts Efflorescence, Deliquescence, Hygroscopy: Some hydrated salts lose water of crystallization to the atmosphere (efflorescence). Some salts absorb moisture from the atmosphere and dissolve in it (deliquescence, e.g., CaCl$_2$). Some absorb moisture without dissolving (hygroscopy, e.g., silica gel). 2.9 Rules of Solubility of Salts in Water These rules are crucial for predicting whether a salt will dissolve or form a precipitate.
1. Nitrates (NO$_3$$^-$): All nitrates are soluble. (e.g., NaNO$_3$, Ca(NO$_3$)$_2$)
2. Chlorides (Cl$^-$): All chlorides are soluble, EXCEPT silver chloride (AgCl), lead(II) chloride (PbCl$_2$), and mercury(I) chloride (Hg$_2$Cl$_2$). (
Note: PbCl$_2$ is sparingly soluble in cold water but soluble in hot water).
3. Sulphates (SO$_4$$^{2-}$): All sulphates are soluble, EXCEPT barium sulphate (BaSO$_4$), lead(II) sulphate (PbSO$_4$), calcium sulphate (CaSO$_4$), and strontium sulphate (SrSO$_4$).
4. Group 1 Metal Salts (Na$^+$, K$^+$) and Ammonium Salts (NH$_4$$^+$): All salts of Group 1 metals and all ammonium salts are soluble. (This rule generally overrides other rules for these specific ions).
5. Carbonates (CO$_3$$^{2-}$): Most carbonates are insoluble, EXCEPT those of Group 1 metals and ammonium carbonate. (e.g., CaCO$_3$, FeCO$_3$ are insoluble, but Na$_2$CO$_3$, (NH$_4$)$_2$CO$_3$ are soluble).
6. Hydroxides (OH$^-$): Most hydroxides are insoluble, EXCEPT those of Group 1 metals. Calcium hydroxide (Ca(OH)$_2$), barium hydroxide (Ba(OH)$_2$), and strontium hydroxide (Sr(OH)$_2$) are sparingly soluble.
7. Sulphides (S$^{2-}$): Most sulphides are insoluble, EXCEPT those of Group 1 metals and ammonium sulphide. --- 2.1 Definition of Acids, Bases, and Salts Acids: Arrhenius Definition: A substance that produces hydrogen ions (H$^+$) or hydronium ions (H$_3$O$^+$) when dissolved in water.
Example: HCl(aq) → H$^+$ (aq) + Cl$^-$(aq)
Common examples: Hydrochloric acid (HCl), Sulfuric acid (H$_2$SO$_4$), Nitric acid (HNO$_3$), Ethanoic acid (CH$_3$COOH – found in vinegar), Citric acid (found in oranges, lemons, limes), Lactic acid (found in fermented pap or ogi). Brønsted-Lowry Definition: A proton (H$^+$) donor.
Example: HCl + H$_2$O → H$_3$O$^+$ + Cl$^-$. Here, HCl donates a proton to H$_2$
O. Bases: Arrhenius Definition: A substance that produces hydroxide ions (OH$^-$) when dissolved in water.
Example: NaOH(aq) → Na$^+$ (aq) + OH$^-$(aq)
Common examples: Sodium hydroxide (NaOH), Potassium hydroxide (KOH), Calcium hydroxide (Ca(OH)$_2$), Ammonia (NH$_3$). In Nigeria, ash extract (potash) and caustic soda are common examples used in soap making. Brønsted-Lowry Definition: A proton (H$^+$) acceptor.
Example: NH$_3$ + H$_2$O ⇌ NH$_4$$^+$ + OH$^-$. Here, NH$_3$ accepts a proton from H$_2$
O. Alkalis: Soluble bases are called alkalis. All alkalis are bases, but not all bases are alkalis.
Salts: Ionic compounds formed when the hydrogen ions of an acid are wholly or partially replaced by a metal ion or an ammonium ion (NH$_4$$^+$). They are typically formed from the reaction between an acid and a base (neutralization reaction).
Example: NaCl (Table salt), CuSO$_4$ (Copper(II) sulphate), KNO$_3$ (Potassium nitrate). 2.2 Identification of Acids and Bases Properties of Acids:
1. Taste: Sour (e.g., citrus fruits). (Students should NOT taste laboratory acids).
2. Litmus Test: Turns blue litmus paper red.
3. Reaction with metals: Reacts with active metals (e.g., Zinc, Magnesium) to produce hydrogen gas.
Example: Zn(s) + H$_2$SO$_4$(aq) → ZnSO$_4$(aq) + H$_2$(g)
4. Reaction with carbonates/bicarbonates: Reacts with carbonates and bicarbonates to produce carbon dioxide gas, water, and a salt.
Example: CaCO$_3$(s) + 2HCl(aq) → CaCl$_2$(aq) + H$_2$O(l) + CO$_2$(g) (This is why sour substances like vinegar react with eggshells or limestone).
5. Corrosive nature: Strong acids are corrosive (cause burns).
6. Conductivity: Conduct electricity in aqueous solutions.
Properties of Bases:
1. Taste: Bitter (e.g., soap). (Students should NOT taste laboratory bases).
2. Feel: Soapy or slippery feel.
3. Litmus Test: Turns red litmus paper blue.
4. Corrosive nature: Strong bases are corrosive.
5. Conductivity: Conduct electricity in aqueous solutions (if soluble).
6. Neutralization: Reacts with acids to form salt and water. 2.3 Nature of Proton in an Aqueous Solution When an acid (e.g., HCl) dissolves in water, it releases hydrogen ions (H$^+$).
However, these H$^+$ ions are extremely small and reactive. They do not exist freely in the aqueous solution. Instead, they immediately combine with water molecules (H$_2$O) to form hydronium ions (H$_3$O$^+$).
This reaction can be represented as: H$^+$ (aq) + H$_2$O(l) → H$_3$O$^+$ (aq) Therefore, the acidic properties of solutions are due to the presence of hydronium ions, not free protons. 2.4 Neutralization Reactions Definition: A chemical reaction between an acid and a base (or alkali) to form a salt and water. The products are generally neutral.
General Equation: Acid + Base → Salt + Water Ionic Equation (Net): H$^+$ (aq) + OH$^-$(aq) → H$_2$O(l) This equation shows the essence of neutralization: the combination of hydrogen ions (from the acid) and hydroxide ions (from the base) to form water. The other ions (spectator ions) remain in solution.
Examples:
1. Hydrochloric acid and Sodium hydroxide: HCl(aq) + NaOH(aq) → NaCl(aq) + H$_2$O(l) (Sodium chloride is table salt, commonly used in Nigerian homes for cooking).
2. Sulfuric acid and Potassium hydroxide: H$_2$SO$_4$(aq) + 2KOH(aq) → K$_2$SO$_4$(aq) + 2H$_2$O(l)
3. Nitric acid and Calcium hydroxide: 2HNO$_3$(aq) + Ca(OH)$_2$(aq) → Ca(NO$_3$)$_2$(aq) + 2H$_2$O(l)
Importance of Neutralization: Agriculture: Reducing soil acidity by adding lime (calcium hydroxide/oxide).
Medicine: Treating indigestion with antacids (e.g., magnesium hydroxide, aluminium hydroxide) to neutralize excess stomach acid.
Industrial waste treatment: Neutralizing acidic or basic effluents before discharge.
Household: Using a mild acid (like vinegar) to neutralize a strong base spill, or vice versa. 2.5 How an Acid-Base Indicator Works Definition: An acid-base require specific pH conditions for optimal results. 2.7 Identification and Preparation of Salts (Normal, Acidic, Basic)
Types of Salts:
1. Normal Salts: Formed when all the replaceable hydrogen ions of an acid are completely replaced by metal ions or ammonium ions. These salts do not contain replaceable H$^+$ or OH$^-$ ions.
Examples: NaCl, Na$_2$SO$_4$, KNO$_3$, (NH$_4$)$_2$CO$_3$.
2. Acidic Salts: Formed when only part of the replaceable hydrogen ions of a polyprotic acid (an acid with more than one replaceable hydrogen) are replaced by metal ions or ammonium ions. These salts contain replaceable hydrogen ions.
Examples: NaHSO$_4$ (Sodium hydrogen sulphate), NaHCO$_3$ (Sodium hydrogen carbonate, baking soda).
3. Basic Salts: Formed when only part of the replaceable hydroxide ions of a polyacidic base (a base with more than one replaceable hydroxide) are replaced by acidic radicals. These salts contain replaceable hydroxide ions.
Examples: Zn(OH)Cl (Zinc hydroxychloride), Mg(OH)NO$_3$ (Magnesium hydroxynitrate).
Methods of Preparing Salts: The choice of method depends on whether the salt is soluble or insoluble in water, and the reactivity of the metal involved.
A. Preparation of Soluble Salts:
1. Neutralization (Acid + Soluble Base/Alkali): Used to prepare soluble salts from acids and soluble bases (alkalis).
Example (Preparation of NaCl): HCl(aq) + NaOH(aq) → NaCl(aq) + H$_2$O(l)
Procedure: Titrate a known volume of acid with the alkali using an indicator to determine the endpoint. Repeat without the indicator, then evaporate the solution to crystallize the salt.
2. Reaction of Acid with an Active Metal: Used for soluble salts of metals above hydrogen in the reactivity series. Example (Preparation of MgSO$_4$): Mg(s) + H$_2$SO$_4$(aq) → MgSO$_4$(aq) + H$_2$(g)
Procedure: Add excess magnesium powder to dilute sulfuric acid until effervescence stops. Filter off excess magnesium. Evaporate the filtrate to crystallize magnesium sulfate.
3. Reaction of Acid with an Insoluble Metal Oxide/Hydroxide: Used for soluble salts of less reactive metals. Example (Preparation of CuSO$_4$): CuO(s) + H$_2$SO$_4$(aq) → CuSO$_4$(aq) + H$_2$O(l)
Procedure: Heat dilute sulfuric acid and add excess copper(II) oxide powder while stirring until no more dissolves. Filter off the excess copper(II) oxide. Evaporate the filtrate to crystallize copper(II) sulphate.
4. Reaction of Acid with a Carbonate: Used for soluble salts of metals whose carbonates are insoluble. Example (Preparation of CaCl$_2$): CaCO$_3$(s) + 2HCl(aq) → CaCl$_2$(aq) + H$_2$O(l) + CO$_2$(g)
Procedure: Add excess calcium carbonate (e.g., powdered eggshells, limestone) to dilute hydrochloric acid until effervescence stops. Filter off excess carbonate. Evaporate the filtrate to crystallize calcium chloride.
B. Preparation of Insoluble Salts:
1. Precipitation (Double Decomposition): Mixing two solutions containing soluble salts whose ions combine to form an insoluble salt precipitate. Example (Preparation of BaSO$_4$): BaCl$_2$(aq) + Na$_2$SO$_4$(aq) → BaSO$_4$(s) + 2NaCl(aq)
Procedure: Mix solutions of barium chloride and sodium sulphate. A white precipitate of barium sulphate will form. Filter, wash the precipitate with distilled water, and dry it. 2.8 Properties of Salts
1. State: Most salts are crystalline solids at room temperature.
2. Melting and Boiling Points: Generally have high melting and boiling points due to strong ionic bonds.
3. Solubility: Many salts are soluble in water, but some are insoluble (refer to solubility rules below).
4. Conductivity: Do not conduct electricity in solid state.
However, they conduct electricity in molten state or when dissolved in water (as electrolytes), because their ions become mobile.
5. Hydrolysis: Some salts react with water to produce solutions that are acidic or basic (e.g., NH$_4$Cl solutions are acidic, Na$_2$CO$_3$ solutions are basic, NaCl solutions are neutral).
6. Hydration: Some salts crystallize with water molecules trapped within their crystal lattice, forming hydrated salts (e.g., CuSO$_4$.5H$_2$O, Na$_2$CO$_3$.10H$_2$O).
7. Efflorescence, Deliquescence, Hygroscopy: Some hydrated salts lose water of crystallization to the atmosphere (efflorescence). Some salts absorb moisture from the atmosphere and dissolve in it (deliquescence, e.g., CaCl$_2$). Some absorb moisture without dissolving (hygroscopy, e.g., silica gel). 2.9 Rules of Solubility of Salts in Water These rules are crucial for predicting whether a salt will dissolve or form a precipitate.
1. Nitrates (NO$_3$$^-$): All nitrates are soluble. (e.g., NaNO$_3$, Ca(NO$_3$)$_2$)
2. Chlorides (Cl$^-$): All chlorides are soluble, EXCEPT silver chloride (AgCl), lead(II) chloride (PbCl$_2$), and mercury(I) chloride (Hg$_2$Cl$_2$). (
Note: PbCl$_2$ is sparingly
Soil Management in Agriculture (Integration with Biology/Agricultural Science): Application: Many Nigerian soils, especially in rainforest zones, tend to be acidic due to high rainfall and decomposition of organic matter. Crops like cocoa, oil palm, maize, and cassava have optimal pH ranges. If the soil is too acidic (e.g., pH 4-5), nutrients become locked up, and toxic elements become available, leading to poor yields.
Local Context: Farmers use substances like powdered limestone (calcium carbonate, CaCO$_3$) or slaked lime (calcium hydroxide, Ca(OH)$_2$) to neutralize acidic soils, a process called liming. This increases the pH, making essential nutrients more available for crop uptake, thus improving food security and farmer income. Students can investigate local soil samples and discuss remediation strategies. Food Fermentation and Preservation (Integration with Food Science/Home Economics): Application: Traditional Nigerian food processing often involves fermentation, where microorganisms produce acids (e.g., lactic acid in ogi/pap, acetic acid in palm wine). The change in pH during fermentation is crucial for flavour development, texture, and preservation. For example, the low pH of garri (from cassava fermentation) inhibits the growth of spoilage bacteria.
Local Context: Students can analyze the pH changes during the fermentation of local foods like ogi (pap), garri, or souring of milk. Understanding pH helps in controlling the fermentation process to achieve desired product quality and shelf life, which is vital for small-scale food entrepreneurs in Nigeria. Traditional Soap Making (Integration with Vocational Skills/History): Application: In many rural communities in Nigeria, traditional soap is made using ash extract (potash), which is a crude source of potassium hydroxide (a strong base). This base reacts with oils (fats) in a process called saponification to produce soap and glycerol.
Local Context: Students can learn about the chemical basis of traditional soap making. The teacher can demonstrate how ash water (an alkaline solution) is traditionally prepared and used. This connects the concept of bases and neutralization to a historical and vocational skill relevant to local economies. ---