Chemistry Unit 3 Notes PDF

Unit 3: Equilibrium, Acids and Redox Reaction Chemical Equilibrium Systems Open and Closed Systems - A System: any part of the universe that is being studied, e.g., an ocean or contents of a test tube - Open system: One that interacts with its environment - both energy and matter can...

Unit 3: Equilibrium, Acids and Redox Reaction Chemical Equilibrium Systems Open and Closed Systems - A System: any part of the universe that is being studied, e.g., an ocean or contents of a test tube - Open system: One that interacts with its environment - both energy and matter can move in and out of the system. Energy and matter are constantly moving between the system and the environment - Equilibrium is not achieved - Closed System: Energy may still be able to flow in and out between the system and the environment, but matter cannot enter or leave the system - Equilibrium is achieved in a closed system when the rates of the forward and reverse reactions are equal Chemical Equilibrium - Reactions can go in two directions, forward and reverse - All reactions are considered to be reversible under suitable conditions - When both reactions continue but there are no more changes in concentration, then the system is at equilibrium Physical Changes - Usually reversible - ​H2O(s) ⇌ H2O(l) ⇌ H2O(g) Chemical Changes - A change in which one or more new substances are formed - Occur in both open and closed systems - Not usually reversible - Many chemical reactions are reversible in a closed system Reversible Reactions - A chemical reaction in which the products can react together to re-form the reactions is called a reversible reaction - A⇌B+C - The forward and reverse reactions can occur at the same time - Can be explained by referring to energy profile diagrams - At a point, the forward reaction rate will equal the reverse reaction rate - When this occurs, the reaction has reached equilibrium - A reversible chemical reaction is at chemical equilibrium when the rate of its forward reaction equals the rate of its reverse reaction and the concentrations of its products and reactants remain unchanged - Even though the rates of the forward and reverse are equal, the concentrations of components on both sides may not be equal - An equilibrium position may be shown - Note the emphasis of the arrows direction - It depends on which side is favoured Equilibrium, a DYNAMIC State - Dynamic: the forward and reverse reactions are still taking place. But overall, there is no change in the equilibrium state - Dynamic equilibrium is reached by reversible physical or chemical reactions taking place in a closed system Extent of Reaction - Different reactions proceed to different extents, i.e, the ratio of reactant to products is different for different equilibrium systems - The extent of reaction describes how far the reaction has proceeded in the forward direction (how much product is formed) when equilibrium is achieved. - The extent of a reaction does not give any information about how fast a reaction will proceed. It only indicates how much product is formed once the system is at equilibrium Factors That Affect Equilibrium Le Chatelier’s Principle - The french chemist Henri Le Chatelier studied how the equilibrium position shifts as a result of changing conditions - Le Chatelier’s principle: if stress is applied to a system in equilibrium, the system changes in a way that relieves the stress - Items considered to be stress on the equilibrium: - Concentration - Adding more reactant produces more product and removing the product as it forms will provide more product - Adding more products produces more reactants and removing reactants as it forms will create more reactants - Pressure: - changes in pressure will only affect gaseous equilibria - Increasing the pressure will usually favour the direction that has fewer molecules - N2(g) + 3H2(g) ⇌ 2NH3(g) - For every 2 molecules of ammonia made, four molecules of reactant are used up - this equilibrium shifts to the right with an increase in pressure - Changing pressure by adding an inert gas - Pressure can be increased in a container, without changing the volume, by adding an inert gas - The equilibrium is not affected as the concentration of the reactants or products remains the same - Dilution - equilibria in solution - Dilution lowers the concentration of particles - The equilibrium will shift in the direction to increase the number of particles - Temperature: increasing the temperature causes the equilibrium position to shift in the direction that absorbs heat - If heat is one of the products (just like a chemical), it is part of the equilibrium  - So cooling an exothermic reaction will produce more product, and heating it would shift the reaction to the reactant side of the equilibrium Effect of a Catalyst on Equilibrium - A catalyst lowers the activation energy of both the forward and reverse reactions by the smae amount - Does not change the relative concentrations of the reactants or products - Does not change the position of equilibrium - Lowering the activation energy of a reaction by the addition of a catalyst does not change the position of equilibrium of a system, it only affects how quickly equilibrium is attained Equilibrium Constants - If the concentrations of reactants and products are known, then reaction conditions can be changed to maximise product formation Equilibrium Constants: Keq - Chemists generally express the position of equilibrium in terms of numerical values - These values relate to the amounts (molarity) of reactants and products at equilibrium - This is called the equilibrium constant, Keq - There are types of equilibrium constants - KC: Concentration - KP: Pressures  - Kb: Weak Bases - Ka: Weak Acids - KSp: Solubility Product The Expression for the Equilibrium Law - Consider this reaction: aA + bB ⇌ cC + dD - The equilibrium constant is the ratio of product concentration to the reactant concentration at equilibrium, with each concentration raised to a power (which is the balancing coefficient) - Thus, the “equilibrium constant expression” has a general form: - NOTE: Pure solids and liquids are left out since the concentrations will not change (for example: H2O(l)) Determining Equilibrium Constant - To determine the equilibrium constant, we have to look at the balanced chemical equation - 2SO2(g) + O2(g) ⇌ 2SO3(g) - After equilibrium is reached, the concentrations of the reactants and product should remain constant; so a ratio of their concentrations should also be constant Value of the Equilibrium constant - Kc helps determine the product yield Reaction Quotient, QC - Can be used to calculate the ratio of products to reactants at any time during a reaction - If QC is greater than KC, the system shifts to the left (The reaction will proceed in the reverse direction) - If QC is less than KC, the system shifts to the right (The reaction will proceed in the forward direction) - If QC equals KC, the system is at equilibrium Effect of Temperature on Equilibrium Constant - The value of KC depends only upon the temperature - It is essential to specify the temperature at which an equilibrium constant has been measured Exothermic Reactions: For a reaction that releases heat, increasing the temperature shifts the equilibrium position toward the reactants (the left side), resulting in a decrease in K. Conversely, decreasing the temperature shifts the equilibrium position toward the products (the right side), resulting in an increase in K. Endothermic Reactions: For a reaction that absorbs heat, increasing the temperature shifts the equilibrium position toward the products (the right side), resulting in an increase in K. Decreasing the temperature shifts the equilibrium position toward the reactants (the left side), resulting in a decrease in K. - Only a change in temperature will change the value of KC for a given reaction Rice Tables R → Balanced Chemical Equation I → Initial concentration C → Change in concentration E → Concentration at equilibrium Equilibrium in Acid and Base Chemistry Bronsted-Lowry Model - Acid: Donates protons (H+ ions ) to bases - Bases: Accepts protons from acids - Acid-Base Reaction: Exchange of protons from an acid to a base - Bronsted-lowry are not restricted to aqueous solutions as Arrhenius are (broader definition compared to Arrhenius) Acids are Proton Donors - All acids contain hydrogen - When acids dissolve in water, they donate hydrogen ions (H+) - For example, when HCL is added to water, the hydrogen atom in HCL bonds to one of the two lone pairs of oxygen in a water molecule. The HCl bond breaks and a hydronium ion and a chloride ion form - HCl(g) + H 2O(l) → H 3O+(aq) + Cl-(aq) Monoprotic Acid - An acid that can donate one hydrogen ion (proton) per molecule - Hydrogen chloride (HCL), hydrogen bromide (HBR), nitric acid (HNO3) and ethanoic acid (CH3COOH) Polyprotic Acids - An acid that can donate more than one proton per molecule - Diprotic acids e.g. sulphuric acid (H2SO4) - Triprotic acids e.g. phosphoric acid (H3PO4) H+ Transfer - acid/base reactions involve ‘proton transfer’ - The reaction does not need to occur in water - H+ leaves one compound and is transferred to another - NH3 + H2O ⇌ NH4+ + OH- Conjugate Acid-Base Pairs - The acid on one side becomes the base on the other side and vice versa - Formulas differ by only a hydrogen ion (H+) - Conjugate bases contains one less H+ in its formula - Conjugate acids contain one more H+ in its formula - Cl- (base) is formed when HCL (acid) loses a proton (donates it to H3O+) - Cl- is the conjugate base of HCl - HCl is the conjugate acid of Cl- - H3O+ (acid) is formed when H2O (base) gains a proton (accepts it from HCL) - H3O+ is the conjugate acid of H2O - H2O is the conjugate base of H3O+ Strong and Weak Acids and Bases - The strength of an acid or base is determined by the level of dissociation (not concentration) - A strong acid completely ionises in water - A strong base completely ionises in water - A weak acid does not completely ionise in water - A weak base does not completely ionise in water Strong Acids - Strong acids ionise completely - The single reaction arrow in each equation below indicates that the dissociation reaction is complete Weak Acids - Weak acids only partially ionise - An equilibrium is eventually established - E.g. ethanoic acid is a weak acid and when dissolved in water only about 5% of the molecules ionise - The partial dissociation is shown in an equation using reversible (double) arrows Strong vs Weak Bases - Strong bases completely ionise - e.g. NaOH(s) → Na+(aq) + OH-(aq) - Weak bases only partially ionise - e.g. NH3(g) + H2O(l) ⇌ NH4+(aq) + OH-(aq) Concentration vs Strength - Concentration refers to the number of particles of solute in a particular volume of solution (commonly mol/litre) - You can therefore have a concentrated solution of a strong or weak acid or base, or a dilute solution of a strong or weak acid or base Electrical Conductivity - Electrolyte - a substance that conducts electricity when dissolved in a solution - Acids and bases are electrolytes because they dissociate into positive and negative ions - The ions can migrate towards oppositely charged terminals of a potential difference which is what an electric current is in solution - Given an equal concentration of an acid and/or base the strength, of the acid and/or base, can be compared by the electrical conductivity of the solution pH Scale Water as a Weak Electrolyte - In water, a small fraction of H2O molecules react to form H3O+ and OH- ions - This process is called the self-ionisation of water and can be represented by the following equilibrium expression  - In pure water, the concentration of hydrogen ions is equal to the concentration of hydroxide ions - At 25oC, the concentration of each ion is 10-7 mol L-1 - I.e. the extent of self-ionisation is very low which is why water is a weak electrolyte - For every H3O+ ion, there are 560 million H2O molecules Ionic Product of Water (Kw) - Also known as the ionisation constant - Kw = [H3O+][OH-] - = 10-7 x 10-7 - = 1x10-14 at 25oC - Because Kw is a constant, at specific temperatures, the concentration of hydrogen and hydroxide ions can be determined in aqueous solutions pH - The pH of a solution is defined as: - pH = -log10[H3O+] - This expression can be rearranged to give: - [H3O+] = 10-pH mol L-1 Effect of Temperature on pH - The value of Kw increases as temperature increases - NB – the pH of pure water changes with temperature. These samples are all neutral. Calculating pOH - At 25oC the pH of pure water is 7 and the pOH is 7 as the concentrations of hydrogen and hydroxide ions are equal: - Ph + pOH = 14 - Hydroxide ion concentration can be calculated using the following formulas: - pOH = -log10[OH-] - [OH-] = 10-pOH mol L-1 Amphoteric/Amphiprotic Substances - Can act as either an acid or base - Must have an ‘H’ in the formula (to donate) - Must have a free e- pair (to accept a H+) - Water is amphoteric - It can donate or accept an H+ - It depends on what it is combined with Buffers - Buffer solutions are conjugate in nature and resist a change in pH when a small amount of an acid or base is added - An acid must be present to react with any OH- added, and a base must be present to react with any H3O+ added - When an acid or a base is added to water, the pH changes drastically  - In a buffer solution, the pH is maintained; pH does not change when acids or bases are added How Buffers Work - Buffers work because: - They resist changes in pH from the addition of an acid or a base - In the body, they absorb H3O+ or OH- from foods and cellular processes to maintain pH - They are important in the proper functioning of cells and blood - They maintain a pH close to 7.4 in blood - A change in the pH of the blood affects the uptake of oxygen and cellular processes - In the buffer with acetic acid (HC2H3O2) and sodium acetate (NaC2H3O2), - The salt produces acetate ions and sodium ions - NaC2H3O2(aq) → C2H3O2−(aq) + Na+(aq) - The salt is added to provide a higher concentration of the conjugate base C2H3O2− than from the weak acid alone Components of a Buffer - Contains a combination of acid-base conjugate pairs, a weak acid and a salt of its conjugate base, such as - HC2H3O2(aq) and C2H3O2−(aq) - Has equal concentrations of a weak acid and its salt Function of a Weak Acid in a Buffer - If a small amount of base is added to this same buffer solution, it is neutralised by the acetic acid, HC2H3O2, which shifts the equilibrium in the direction of the products, acetate ion and water - HC2H3O2(aq) + OH-(aq) → C2H3O2− (aq) + H2O(l) - Added OH⁻ ions react with H⁺ ions in the solution, forming water (H₂O). - This reaction reduces the concentration of H⁺, which would normally increase the pH. - To counter this, acetic acid (HC₂H₃O₂) dissociates to release more H⁺ ions, replenishing the H⁺ ions that were removed. - The equilibrium shifts to the right, as acetic acid dissociates to form additional H⁺ and acetate ions (C₂H₃O₂⁻). Function of Conjugate Base in a Buffer - When a small amount of acid is added, the additional H3O+ combines with the acetate ion, C2H3O2−, causing the equilibrium to shift in the direction of the reactants, acetic acid and water - The acetate ions act like a sponge, absorbing added H3O+ and stabilising the pH by converting the H3O+ into acetic acid, a weak acid that doesn't fully dissociate, thus the pH does not decrease - HC2H3O2(aq) + H2O(l) ← C2H3O2− (aq) + H3O+(aq) Working Buffers - The buffer described here consists of about equal concentrations of acetic acid (HC2H3O2) and its conjugate base, acetate ion (C2H3O2−) - Adding H3O+ to the buffer reacts with the salt, C2H3O2−, whereas adding OH- neutralises the acid HC2H3O2  - The pH of the solution is maintained as long as the added amounts of acid or base are small compared to the concentrations of the buffer components Dissociation Constants Acid Dissociation Constant - The acid dissociation constant (Ka) is the ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form - The dissociated form includes both the H3O+ and the anion - A strong acid, such as HCL, completely dissociates in water - As a result, [H3O+] is high in an aqueous solution of strong acid - By contrast, weak acids remain largely undissociated - In an aqueous solution of ethanoic acid, less than 1% of the molecules are ionised - You can use a balanced equation to write the equilibrium-constant expression for a reaction - The acid dissociation constant expression shown below is for ethanoic acid - The acid dissociation constant (Ka) reflects the fraction of an acid that is ionised  - Thus, dissociation constants are sometimes called ionisation constants - If the degree of dissociation or ionisation of the acid is small, the value of the dissociation constant will be small - Weak acids have small Ka values - If the degree of ionisation of an acid is more complete, the value of Ka will be larger - The stronger an acid is, the larger its Ka value will be - Nitrous acid (HNO2) has a Ka of 4.4 × 10-4, but ethanoic acid (CH3COOH) has a Ka of 1.8 × 10-5. - This means that nitrous acid is more ionised in solution than ethanoic acid - Nitrous acid is a stronger weak acid than ethanoic acid - Some acids have more than one dissociation constant because they have more than one ionisable hydrogen - Oxalic acid is a diprotic acid - It loses two hydrogens, one at a time. - Therefore, it has two dissociation constants. - Oxalic acid is found naturally in certain herbs and vegetables. - With each ionisation, the Ka decreases (first to second to third) Calculating Dissociation Constants - To calculate the acid dissociation constant (Ka) of a weak acid, you need to know the initial molar concentration of the acid and [H+] (or alternatively, the pH) of the solution at equilibrium - You can use these data to find the equilibrium concentrations of the acid and the ions - These values are then substituted into the expression for Ka - You can find the Ka of an acid in water by substituting the equilibrium concentrations of the acid, [HA], the anion from the dissociation of the acid, [A-], and the hydrogen ion, [H+] into the equation below Base Dissociation Constant - The base dissociation constant (Kb) is the ratio of the concentration of the conjugate acid times the concentration of the hydroxide ion to the concentration of the base - - Just as there are strong acids and weak acids, there are strong bases and weak bases  - A strong base dissociates completely into metal ions and hydroxide ions in aqueous solution - A weak base reacts with water to form the conjugate acid of the base and hydroxide ions (e.g. ammonia) - - For a weak base, the amount of dissociation is relatively small - When equilibrium is established, only about 1% of the ammonia is present as NH4+ - This ion is the conjugate acid of NH3 - The concentrations of NH4+ and OH- are low and equal - The base dissociation constant expression for the dissociation of ammonia in water is as follows: - - The magnitude of Kb indicates the ability of a weak base to compete with the very strong base OH- for hydrogen ions - Because bases such as ammonia are weak relative to the hydroxide ion, the Kb for such a base is usually small - The smaller the value of Kb the weaker the base Relationship Between Ka and Kb - If you multiply Ka of an acid by Kb of its conjugate base, you get: - Kw = Ka × Kb = 1.0 × 10-14 at 25oC pKa and pKb - Values for Ka and Kb span a wide range - acid/base dissociation constants are often expressed as pKa/pKb where: - - The lower the value of pKa/pKb the greater the ionisation in water Acid-base Indicators Weak Acids - Most indicators are weak acids (HIn) - In solution, the acid is in equilibrium with its conjugate base (In-) - - The acid and its conjugate base are different colours - At low pH, the equilibrium position is on the left, and the solution shows colour A - At high pH, the equilibrium position is on the right, and the solution shows colour B Weak Bases - Some indicators are weak bases - - At high pH, the equilibrium position is on the left, and the solution shows colour A - At low pH, the equilibrium position is on the right, and the solution shows colour B pH Range and pKa - Indicators change colour when the concentration of the acid HIn, is equal to the concentration of its conjugate base, In- - HIn(aq) ⇌ In- (aq) + H+(aq) - - Colour change occurs when, [In-] = [HIn], therefore: -  - Which leads to: - - This means that the colour of the indicator changes when pKa = pH - Different indicators change colour over different pH ranges - Most indicators change colour over a range of pKa +- 1 Indicators for titration - Titration is a method for determining the concentration of an unknown acid or base by neutralising it with a known base or acid - The equivalence point in a titration is when the reactants have reacted in the molar ratio of the balanced chemical equation - The endpoint is the point in a titration when the indicator changes colour - It is important to choose an indicator that changes colour close to the equivalence point Titration Setup - Typically, the titrant (the solution of known concentration) is added through a burette to a known volume of the analyte (the solution of unknown concentration) in a conical volumetric flask until the reaction is complete. - Knowing the volume of titrant added (aliquot) via a pipette allows us to determine the concentration of the unknown analyte. - Often, an indicator is used to signal the end of the reaction, the endpoint. - Endpoint refers to the point at which the indicator changes colour in an acid-base titration Titration Curves - A titration curve is the plot of the pH of the analyte solution versus the volume of the titrant added as the titration progresses - - The half-equivalence point occurs at one-half the volume of titrant needed to reach the equivalence point - The half equivalence point is found in the buffer region - At the half equivalence point the pH = the pKa of the acid  - The pH at the half-equivalence point is not half the pH at the equivalence point A graph of pH against the volume of reactant added - Titration of a strong acid with a strong base - - Titration of a weak acid with a strong base - - Titration of a weak acid with a weak base - The pH change is gradual and does not show a steep increase. This behaviour makes it impossible to use an indicator to find the equivalence point - Data test steps: - Neutralisation Reactions - General equation: - Acid + Base → Salt + Water - - Volumetric Analysis Standard Solutions - Standard Solutions are solutions of an accurately known concentration; prepared by dissolving an accurately measured mass of a primary standard in an accurately measured volume of water. - Volumetric Analysis: Quantitative analysis using measurement of solution volumes, usually involving titration. It involves reacting solutions of unknown concentration with a solution of accurately known concentration. (standard solution) - Primary Standard: Pure substance widely used in the laboratory to prepare solutions of accurately known concentration. - Readily available in pure form - Have a known chemical formula - Be easy to store without reacting or deteriorating on the atmosphere - Doesn’t absorb water eg KOH, NaOH - Doesn’t absorb CO2 eg KOH, NaOH - Lose water to the atmosphere eg hydrated salts - Have a high molar mass - Eliminates the effect of errors in weighing  - Be Inexpensive Preparing a Standard Solution 1. Use the chemical formula to determine the molar mass (M) 2. Use the molar mass (M) and the mass (m) to determine the number of moles (n) 3. Use the number of moles (n) and the volume of the flask to determine the concentration of the standard solution