UNIT 3: ACIDS, BASES AND SALTS
Main Ideas from the Transcript
Page 3
Learning Objectives
Characteristics of acids and bases
Concepts: Arrhenius, Bronsted Lowry, Lewis
Chemical reactions of acids and bases
pH value, types of indicators
Strength of Acids and Bases
Dilute and Concentrated Acids and Bases
Oxides and types
Types and preparation of salts
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Unit Overview
Statement of inquiry: Properties of acids and bases determine their use and function
Key concept: System
Related concepts: Function, evidence
Global context: Identities and relationships
Inquiry questions: Factual, Conceptual, Debatable
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Physical Properties of Acids and Bases
Acids are corrosive, have pH below 7, taste sour
Bases are corrosive, have pH above 7, taste bitter
Bases are slippery, both are good conductors of electricity
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Arrhenius Concept
Definition of acids and bases by Svante Arrhenius
Acids produce H+ ions, bases produce OH- ions in water
Page 14-16
Indicators
Used to indicate acidity, alkalinity, or neutrality
Change color based on the type of substance
Natural indicators like red cabbage, artificial indicators like litmus
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Universal Indicator
Shows the strength of acids or bases
Range of colors to indicate the strength
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Alkali and Base
Alkalis are bases soluble in water
Not all bases are alkalis
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Strength of Acids and Bases
Differences in strength among acids and bases
Use of indicators to measure strength
Page 28-30
Neutralization Reaction
Acid + Base → Water + Salt
Ionic equation for dissociation
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Salts
Formation of salts from acid-base reactions
Salt characteristics based on acid-base strengths
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Different Concepts of Acids and Bases
Bronsted Lowry, Arrhenius, Lewis concepts
Note
This note provides an overview of the key concepts discussed in the transcript related to acids, bases, salts, and indicators. It covers learning objectives, unit overview, physical properties, concepts like Arrhenius, indicators, universal indicators, alkali and base differences, neutralization reactions, and different concepts of acids and bases.
Arrhenius Concept of Acids and Bases
Definition of Acids: Compounds containing hydrogen that produce H+ ions when dissolved in water.
Strong Acid: Completely dissociates in water to give H3O+ ions.
Weak Acid: Partially dissociates in water, producing few H3O+ ions.
Limitations of Arrhenius Concept
H+ and OH- ions are associated with water molecules.
Water presence is necessary for acidic or basic properties.
Cannot explain properties in non-aqueous solvents.
Fails to explain certain salts' acidic character.
Compounds like ammonia behave as bases but are not recognized under this theory.
Bronsted-Lowry Concept
Acid: Proton donor containing H in its formula.
Base: Proton acceptor with a lone pair of electrons.
Acid-base reaction involves proton transfer.
Acids donate protons, bases accept protons.
Implications of Bronsted-Lowry Concept
Acids and bases can play both roles.
Acids donate H+ ions, bases accept H+ ions.
In reactions, acids and bases form without the other conjugated species.
Advantages and Limitations of Bronsted-Lowry Concept
Advantages: Includes ionic species, does not require aqueous medium, explains basic nature of ions without OH-.
Limitations: Cannot explain compounds without tendency to lose or gain H+ ions, or those with OH- ions.
Lewis Definition of Acids and Bases
Lewis Acids: Electron pair acceptors.
Lewis Bases: Electron pair donors.
Bronsted acids are not Lewis acids.
Leads to a general description of reaction patterns without a strength scale.
Reactions of Acids with Metals
Metals react with acids to produce salt and hydrogen gas.
Salt produced depends on the metal and type of acid.
Detection of hydrogen gas production by burning with a pop sound.
Reactions of Metal Oxides with Acids
Metal oxides react with acids to form salts and water.
Example: MgO + 2HCl → MgCl2 + H2O.
Reactions of Metal Hydroxides with Acids
Metal hydroxide + acid → metal salt + water.
Example: Potassium hydroxide + sulfuric acid → potassium sulfate + water.
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Alkali metal hydroxides react with acids to produce soluble chloride and nitrate
Example: Mg(OH)2(s) + 2HCl -> MgCl2 + 2H2O
Reaction with H2SO4 is similar to the oxide reaction
Example: Mg(OH)2(s) + H2SO4 -> MgSO4 + 2H2O
Strontium hydroxide and barium hydroxide reactions stop due to passive action
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All carbonates react with acids to form a salt, carbon dioxide, and water
Example of neutralization reaction
Different acids produce different salts: chlorides, nitrates, sulfates
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Describe reactions of hydrochloric acid and sulfuric acid with metal hydroxides, oxides, and carbonates
Metal hydroxides dissolve in water to form clear solutions, react with acids to form salt and water
Metal oxides react with acids to form salt and water
Metal carbonates react with acids to form salt, water, and carbon dioxide
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Bases release hydroxide ions in water (Arrhenius concept)
Bases accept protons (Bronsted-Lowry concept)
Example: Sodium hydroxide (Arrhenius), ammonia (Bronsted-Lowry)
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Acids and bases neutralize each other
General formula for acid-base reaction: Acid + Base -> Water + Salt
Salt refers to any ionic compound from the reaction
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Alkalis heated with ammonium salts produce ammonia gas
General equation: Alkali + Ammonium Salt -> Ammonia + Water + Salt
Example: Calcium hydroxide + Ammonium chloride -> Calcium Chloride + Water + Ammonia
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Soil pH affects mineral nutrient availability for plants
Alkaline soils limit phosphorous, iron, and zinc
Acidic soils limit calcium and magnesium
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Reasons to lower soil pH: increase availability of minerals like phosphorous, iron, and zinc
Acidic soils make it difficult for weeds to thrive
Fruiting plants benefit from lower soil pH
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Ways to lower soil pH: increase soil nitrogen, add compost or manure
Coffee grounds are not a quick fix for lowering soil pH
Coffee grounds have a pH close to neutral and add nitrogen
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Reasons to raise soil pH: cool-season vegetables perform better in slightly higher pH
Dolomite Lime is a common amendment to raise soil pH
Dolomite Lime should not be used in soils with excess magnesium
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Basic oxides are oxides that exhibit basic properties and react with water to form a base or with an acid to form a salt.
Examples include sodium oxide reacting with water to produce sodium hydroxide.
Basic oxides are mostly oxides of metals, particularly alkali and alkaline earth metals.
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Basic oxides react with acids to produce salts and water only.
Metals are insoluble in water and solid at room temperature.
Only sodium oxide and potassium oxide dissolve in water.
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Acidic oxides are oxides of non-metals like CO2, often gases at room temperature, and dissolve in water.
They do not react with acids but react with alkalis to form salt and water.
Example: Silicon dioxide reacts with hot concentrated sodium hydroxide to form sodium silicate.
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Acidic oxides react with bases and alkalis to produce salts.
Example: carbon dioxide reacts with sodium hydroxide to form sodium carbonate and water.
Page 73
Examples of acidic oxides and the acids they produce.
Example: sulfur trioxide produces sulfuric acid, sulfur dioxide produces sulfurous acid, etc.
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Acidic oxides reactions with water and alkali.
Example reactions of SiO2, P4O6, and P4O10 with water and alkali.
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Amphoteric oxides can act as both an acid and a base depending on the reaction.
Examples include ZnO, Al2O3, PbO, SnO.
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Amphoteric oxides react with acids and alkalis to produce salts.
Example: aluminum oxide reacts with hydrochloric acid to form aluminum chloride and water.
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Amphoteric oxides behave as an acid or a base depending on the reaction.
Aluminum oxide is an example of an amphoteric oxide.
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Types of oxides include acidic, basic, amphoteric, and neutral oxides.
Neutral oxides are non-metallic oxides with low solubility in water and no effect on litmus.
Page 79
Types of oxides: Acidic, Basic, Amphoteric, and Neutral oxides.
Examples of each type of oxide and their reactions with acids and alkalis.
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Flow chart to determine the type of oxide based on solubility in acid and alkali.
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Salts are ionic compounds produced when an acid and a base react.
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Two types of salts: Soluble salts and Insoluble salts.
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Method 1 for preparing salts by reacting acids with metals.
Suitable for metals above hydrogen in the reactivity series.
Example: Preparation of zinc sulphate by reacting zinc with sulfuric acid.
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Procedure for the preparation of zinc sulphate by reacting zinc with sulfuric acid and subsequent steps for crystallization.
Page 86
Limitations of the method of preparing salts by reacting acids with metals.
Not suitable for metals close to or below hydrogen in the reactivity series.
Page 87
Process for making salts from metals, including determining solubility, filtering, and reacting with acids.
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Method 2 for preparing salts: Salts from insoluble bases
Salts of metals low in reactivity series can be made by reacting insoluble bases with acids.
Example: Preparation of copper sulphate: CuO(s) + H2SO4 (aq) ---à CuSO4 (aq) + H2O (g)
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Procedure to form copper sulphate from copper oxide
Step 1: Add copper oxide to sulphuric acid, warm gently.
Step 2: Solution turns blue as copper sulphate forms.
Step 3: Filter solution to remove excess copper oxide.
Step 4: Evaporate water from filtrate and let crystallize.
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Steps to prepare copper(II) sulphate from copper oxide and sulphuric acid
Heat copper oxide and sulphuric acid.
Copper oxide reacts to form copper sulphate crystals.
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Making salts from soluble bases (alkalis)
Mix solutions of soluble compounds to form insoluble salts.
Use solubility rules to determine if a salt is soluble or insoluble.
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Soluble or insoluble salts
Some salts like copper sulphate are soluble, while others are insoluble.
Precipitates are formed when mixing solutions of soluble compounds.
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Solubility rules for compounds
Lists soluble and insoluble compounds based on groups and elements.
Helps in determining which compounds will form precipitates.
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Making an insoluble salt like lead chloride
Identify ions in the insoluble salt.
Use solubility rules to choose soluble compounds with these ions.
Mix solutions, filter off precipitate, wash, and dry the solid.
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Preparation of an insoluble salt
Mixing two solutions of soluble compounds to form a precipitate.
Filter, wash, and dry the precipitate to obtain the insoluble salt.
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Precipitation reaction explanation
Lead nitrate + sodium chloride yields lead chloride precipitate.
Ions in solution move freely until precipitate forms due to stronger attraction.
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Formation of lead chloride precipitate
Lead ions attract chloride ions, forming an insoluble salt lattice.
Sodium and nitrate ions remain in solution as spectator ions.
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Credits and bibliography
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