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Introduction to Acids and Bases

IB Chemistry Power Points Topic 08 Acids and Bases. www.pedagogics.ca. Introduction to Acids and Bases. In aqueous solutions, a proportion of the water molecules dissociate; The ions formed are H + or positively charged hydrogen ions and negatively charged hydroxide ions (OH - ).

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Introduction to Acids and Bases

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  1. IB Chemistry Power Points Topic 08 Acids and Bases www.pedagogics.ca Introduction to Acids and Bases

  2. In aqueous solutions, a proportion of the water molecules dissociate; • The ions formed are H+ or positively charged hydrogen ions and negatively charged hydroxide ions (OH-) Technically 2 H2O(l)H3O+(aq) + OH-(aq) Kw = [H + ][OH − ] = 1 x 10-14

  3. Some chemical compounds contribute additional H+ to make the solution more acidic. Other compounds remove H+ ions. • A compound that increases H+ is called an acid • Examples: HCl, H2SO4, HNO3, CH3COOH

  4. A compound that removesH+ ions from an aqueous solution is called abase.This reaction is called a neutralization. Often this is done by adding OH-ions for example NaOH, KOH, Ca(OH)2. Soluble bases are called alkalis.

  5. Neutralization Reactions • hydroxides acid + base  water + salt HCl + NaOH H2O + NaCl(aq) • metal oxides • acid + base  water + salt • 2HCl + Cu2O  H2O + CuCl2(aq) • ammonia • acid + base  salt • HCl + NH3 NH4Cl (aq)

  6. Three theories of acids • Arrhenius (most common) • Bronsted-Lowry • Lewis

  7. Arrhenius (most common): an acid dissociates to yield H+ and a base dissociates to yield OH- Hydrochloric acidH++ Cl- Sodium hydroxide Na+ + OH-

  8. Bronsted-Lowry: an acid is a proton (H+)donor and a base is a proton acceptor

  9. amphiprotic

  10. Lewis: An acid is an electron pair acceptor and a base is an electron pair donor A dative covalent bond is formed

  11. Example of Lewis Acid Lewis Acid Lewis Base

  12. This is a common example that is not an obvious acid/base rxn Boron trifluorideacts as a Lewis Acid. The boron has only 6 electron in valence shell so the lone pair of electrons forms a dative bond and fills up the valence shell of the boron

  13. Indicators Acids and bases are substances with specific physical and chemical properties. We can determine if substances are acidic or basic by testing their reaction with indicators.

  14. Indicators are organic substances that change color in the presence of an acid or a base. Some common indicators in acid in base Litmus redblue Phenolphthalein colorless pink Methyl orange redyellow

  15. Reactions of acids • React with active metals (above copper in reactivity series) 2 HCl + Ca  CaCl2+ H2 • Reaction with carbonates H2SO4 + Na2CO3 Na2SO4 + CO2 + H2O • Reaction with bicarbonates HNO3 + NaHCO3 NaNO3 + CO2 + H2O

  16. Acid/base properties of Period 3 oxides (topic 3) • Metal oxides Na2O and MgO react with water to form hydroxides (basic solutions) Na2O + H2O  2 NaOH (aq) • Aluminum oxide is amphoteric (will react as a base with an acid or vice versa) Al2O3 + 6 HCl 2 AlCl3 + 3 H2O • Other period 3 oxides (non-metal S, P, Cl oxides) react with water to form acidic solutions SO3+ H2O H2SO4 (aq) see page 15 in study guide

  17. Acid/base properties of Period 3 chlorides (topic 13) • Chlorides across Period 3 become more acidic across the period NaCl (aq) is neutral MgCl2 (aq) is weakly acidic Chlorides of Al, Si, P, S and Cl2 react with water to produce HCl (aq) solutions see Study guide page 16

  18. Strong Acids vs Weak Acids The strength of an acid or base depends on how easily it dissociates in water. The dissociation of an acid or base is an equilibrium. HA(aq)H+(aq) + A-(aq) BOH(aq)B+(aq) + OH-(aq) Strong acids or bases dissociate (ionize) easily – the equilibrium favors the ionic products : kc >> 1

  19. Strong vs Weak When the strength of an acid or base is discussed, it is very important NOT to confuse “strength” with “concentration” A 5M acid solution contains 5 mol of acid per dm3 but its strength is determined by how much of that acid is ionized. Strong acids : HCl, H2SO4, HNO3 (mono vs diprotic) Strong bases : NaOH, KOH, Ba(OH) 2 Weak acids: CH3COOH, H2CO3, carbonic acid CO2(aq) Weak bases: NH3, ethylamine CH3CH2NH2

  20. Strong vs Weak How to tell Strong acids and bases are mostly ionized and therefore solutions are good electrolytes (high conductivity). The pH of the solution can also be measured.

  21. What is the pH scale? • pH is a measurement of hydrogen ion concentration • It tells you how acids or basic (or alkaline) something is • Ranges from 0 (most acidic) to 14 (most basic

  22. The scale is logarithmic. As you go up or down, the concentration is changed by a power of ten Example pH 3 is 100 times more concentrated than pH 5 neutral pH 10 is 100 times less concentrated than pH 8 How does scale work?

  23. Strong Acid example HCl HCl(aq) H+(aq) + Cl-(aq) • completely dissociated • pH of 0.1 M soln = 1 • strong electrolyte • reacts vigorously • note simplified “net ionic” equation

  24. Weak Acid example CH3COOH CH3COOH (aq) H+(aq) + CH3COO-(aq) • partially dissociated • pH of 0.1 M soln = 2.9 • weak electrolyte • reacts slowly

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