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ACIDS AND BASES

ACIDS AND BASES. Chapter 15. Bronsted Lowry Acids and Bases. An acid is a proton donor and a base is a proton acceptor. The loss of a proton is called as deprotonation: HCl(aq)+ H 2 O(l) H 3 O + (aq) +Cl - (aq) A strong acid is fully ionized in water.

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ACIDS AND BASES

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  1. ACIDS AND BASES Chapter 15

  2. Bronsted Lowry Acids and Bases • An acid is a proton donor and a base is a proton acceptor. • The loss of a proton is called as deprotonation: HCl(aq)+ H2O(l)H3O+ (aq) +Cl- (aq) • A strong acid is fully ionized in water. • A weak acid is only partially ionized in water. • A strong acid reacts completely with water to produce hydronium ions. • A weak acid reacts incompletely with water to produce hydronium ions.

  3. A strong acid is fully deprotonated in water. • A weak acid is only partially deprotonated in water. • Common strong acids in water are: • HBr, HCl, HI, HNO3, HClO4 ,HClO3, H2SO4

  4. The process of accepting a proton is called as protonation. • A strong base reacts completely with water to produce hydroxide ion. • A weak base reacts incompletely with water to produce hydroxide ion. • A strong base is completely protonated in water. • A weak base is partially protonated in water.

  5. Arrhenius,Bronsted and Lewis definition of acid /base • An Arrhenius acid or base is defined according to the ability of a compound to produce hydronium ions or hydroxide ions in water. • A Bronsted acid or base is defined according to the ability of a species to donate or accept a proton.Water need not be involved. • A Lewis acid or base is defined according to the ability of a species to donate or accept a pair of electrons and to form a coordinate covalent bond. A proton need not be involved. • AN ACID IS A PROTON DONOR AND BASE IS A PROTON ACCEPTOR.

  6. Conjugate acids and bases • After losing a proton, the acid species becomes the conjugate base. A base and its protonated partner also form a conjugated acid-base pair. • CH3COOH(aq) + H2O(l)↔H3O+ (aq) +CH3COO- (aq) • The water molecule is acting as a Bronsted base and acetic acid is a Bronsted acid. Because the acetate ion a base is formed from acetic acid by proton loss , it is called the conjugate base of acetic acid. • Acid donates H+ conjugate base

  7. When sodium acetate is dissolved in water, an acetate ion can accept a proton from a water molecule and be converted into an acetic acid molecule. • H2O(l)+CH3COO- (aq)↔CH3COOH(aq)+OH- (aq) • In the above reaction water is a Bronsted acid and the acetate ion is a Bronsted base. Because a CH3COOH molecule is an acid formed by attaching a proton to an acetate ion, it is the conjugate acid of the base CH3COO- • Base-accepts H+conjugate acid • The conjugate base of an acid is formed when the acid has donated a proton. The conjugate acid of a base is formed when the base has accepted a proton.

  8. Proton exchange between water molecules • Is water an acid or a base? • Water molecule accepts a proton from an acid molecule to form H3O+ ion. So water is a base. However a water molecule can donate a proton to a base and become an OH- ion. So water is also an acid. A molecule that can act as both a proton donor and a proton acceptor is said to be amphiprotic.

  9. Protons migrate between water molecules even in the absence of another acid or base: • 2H2O(l)↔H3O+ (aq) + OH- (aq) • Because H3O+ is an acid and OH- is a base, the reverse reaction H3O+(aq) +OH-(aq)2H2O(l) also occurs. The transfer of protons is very rapid and the equilibrium • 2H2O(l)↔H3O+ (aq) +OH- (aq) is always present in water and aqueous solutions.This type of reaction where one molecule transfers a proton to another molecule of the same kind is called autoprotolysis.

  10. Kc=[H3O+][OH-]/[H2O]² • In solutions we consider water to be very pure so the molar concentration of water can be treated as a constant and combined with Kc. The resulting expression is called autoprotolysis constant of water and is written as Kw=Kc[H2O]² =[H3O+][OH-] • The molarities of H3O+ and OH- in pure water at 25⁰C are known by experiment to be 1.0x10-7 mol/L • Kw=(1.0x10-7 ) x(1.0x10-7)= 1.0x10-14 • The concentration of H3O+ and OH- are very low in pure water which explains why pure water is a poor conductor of electricity.

  11. Class Practice • What are the molarities of H3O+ and OH- ions in 0.020M HCl (aq) at 25⁰C? • What are the molarities of H3O+ and OH- ions in 0.0030M Ba(OH)2(aq) at 25⁰C?

  12. Homework • Page 694 • 15.13,15.14,15.15,15.16

  13. Weak and Strong Acids and Bases • STRONG ACIDS Acids that are essentially 100% ionized in aqueous solutions ex: HCl, HNO3, HClO4 produce the maximum concentration of H+ [acid] = [H+] • WEAK ACIDS Acids that are partially ionized ( usually less than 5%) in equilibrium. HF + H2O(l) ↔H3O+(aq) + F-(aq) The forward and the reverse reaction are occurring simultaneously most found as HF. • STRONG BASES Those compounds that completely ionize in water to produce OH- ions NaOH(s) ↔ Na+(aq) + OH-(aq) Concentration of base = concentration of hydroxide ions • WEAK BASES NH3(aq) + H2O(l) ↔NH4+(aq) + OH-(aq) equilibrium lies far to the left (mostly reactants present)

  14. Types of acids • Monoprotic - a solution that produces one mole of H+ ions per mole of acidHCl , HNO3 • Diprotic - a solution that produces two moles of H+ ions per mole of acidH2SO4 • Triprotic - a solution that produces three moles of H+ ions per mole of acidH3PO4 • Polyprotic - two ore more H+ per mole of acid

  15. Proton transfer equilibria • This is one of the fastest reaction in aqueous solution. Because it is so fast we can be confident that conjugate acids and bases are always in equilibrium with each other in water. • CH3COOH(aq) +H2O(l)-H3O+(aq)+ CH3COO- (aq) • And H3O+(aq)+ CH3COO- (aq) CH3COOH(aq) +H2O(l) • Can be combined into • CH3COOH(aq) +H2O(l)↔H3O+(aq) +CH3COO- (aq) acid base conjugate acid conjugate base This equilibrium is a proton transfer equilibria

  16. Class Practice • Identify the (a) Bronsted acid and base in the reaction • (a) HNO3(aq)+HPO42- (aq)- NO3- (aq) + H2PO4- (aq) • (b) the conjugate base and acid formed

  17. pH/pOH • The pH scale is defined as the negative log of the concentration of H+: pH = -log[H+] • The pOH scale is defined as the negative log of the concentration of OH-, [OH-]: pOH = -log[OH-]

  18. pH of solutions of weak acids and bases • The acidity or basicity of a substance is defined most typically by the pH value, defined as below: • pH=-log[H+] • At equilibrium, the concentration of H+ is 10-7, so we can calculate the pH of water at equilibrium as: • pH = -log[H+]= -log[10-7] = 7 • Solutions with a pH of seven (7) are said to be neutral, while those with pH values below seven (7) are defined as acidic and those above pH of seven (7) as being basic.

  19. gives us another way to measure the acidity of a solution. It is just the opposite of pH. A high pOH means the solution is acidic while a low pOH means the solution is basic. • pOH = -log[OH-] • pH + pOH = 14.00

  20. Weak acids produce a lower concentration of H₃O+ ions in aqueous solution than do strong acids of the same initial concentration. 0.01 M HCl (aq) has a pH of 2; and 0.01M of CH3COOH has a much lower concentration of H3O+ ions and its pH is 3. To find the H3O+ molarity in a solution of a weak acid, we have to take into account the equilibrium between the acid HA and its conjugate base A- : HA(aq) + H2O(l)↔H3O+(aq) + A- (aq) • Ka= [H3O+][A-]/[HA] • Percentage deprotonated=[H3O+]/[HA]initial x100%

  21. Home work • Page 15.43, 15.47,15.49,15.51

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