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Acids and Bases

Acids and Bases. Definitions:. Arrhenius- Acid- substance that dissociates in water to produce hydrogen ions - H + Examples: HC l , HNO 3 , H 2 SO 4 , etc Base- substance that dissociates in water to produce hydroxide ions- OH - Examples: NaOH , KOH, Mg(OH) 2

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Acids and Bases

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  1. Acids and Bases

  2. Definitions: • Arrhenius- Acid- substance that dissociates in water to produce hydrogen ions - H+ Examples: HCl, HNO3, H2SO4, etc Base- substance that dissociates in water to produce hydroxide ions- OH- Examples: NaOH, KOH, Mg(OH)2 According to this definition, H must be present in an acid, and OH in a base

  3. HNO3 + H2O H3O+ + NO3- H3O+ is the Hydronium ion You may see HNO3 H+ + NO3- • HCl + KOH KCl + H2O Neutralization reaction Acid + Base a Salt + water

  4. 2. Brønsted-Lowry- Acid- substance that can donate H+ ions Base- substance that can accept H+ ions Substances do not need to be in water, and the base does not need to have OH. This expands the Arrhenius definition so that other substances can be considered as acids and bases. H+ is simply a proton, so the definition of an acid can now be a proton donor and a base is a proton acceptor.

  5. NH3 + H2O NH4++ OH- What is the acid? H2O – it donated the H+ What is the base? NH3 – it accepted the H+ Conjugate Acid- Base Pairs: Many times a reaction occurs in a forward and reverse direction, so what is an acid in the forward direction becomes the conjugate base in the reverse direction, and the base becomes the conjugate acid in the reverse reaction

  6. HCl + H2O H3O+ + Cl - acid base conjugate conjugate acid base Monoprotic acid- an acid that will donate 1 H+ HNO3 Diprotic acid- an acid that donates 2 H+ H2SO4 Triprotic acid- an acid that donates 3 H+ H3PO4 Amphoteric- a substance that can act as both an acid or a base- Water

  7. 3. Lewis Acids and Bases- Acid- substance that can accept a pair of electrons to form a covalent bond Base- substance that can donate a pair of electrons to form a covalent bond This now allows other substances to be considered an acid or a base

  8. Properties of Acids and Bases • Acids: -sour or tart taste -can burn skin if stronger, may burn if you get it in a cut -react strongly with most metals, usually to produce hydrogen gas -form weak or strong electrolytes -change blue litmus paper to red

  9. Bases: -bitter taste -feels smooth, slippery -not reactive with most metals -weak or strong electrolytes -turns litmus paper from red to blue

  10. Strengths of Acids and Bases • Strong acid- easily dissociates H+ to water - becomes strong electrolyte • Weak acid- does not easily dissociate H+ ions - weak electrolyte Common strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4 Weak acids:HC2H3O2, HCN, HNO2, HF, HClO, HCO3-

  11. Strong Bases: substances with a strong affinity for H+, those with OH- Weak Bases- ions that react only partially in water to form OH- ions Common strong bases: CaO, NaOH, KOH, Ca(OH)2 Weak bases- NH3, H2NNH2, CO32-, PO43- Strong bases are strong electrolytes, weak bases are weak electrolytes

  12. Self-Ionization of Water • When you have pure water, due to the motion of the water molecules and the high polarity of water, there is a very small amount of H3O+ and OH- that exist H2O + H2O H3O+ + OH-

  13. In pure water, the [H+] = [OH-], where [X] is equal to the concentration of X in solution, usually measured in M • When the concentrations are equal, the substance is known as a neutral solution • [H+] =[OH-] = 1.0 x 10-7 M • So, [H+] x [OH-] = 1.0 x 10-14M2 = Kw = ion product constant of water • Since these concentration values are multiplied to give a constant, if [H+] increases, [OH-] must decrease

  14. Because Kw is a constant, if you know the concentration of H+ or OH-, you can calculate the other. • Example: If [H+] = 1.0 x 10-2, then (1.0 x 10-2) [OH-] = 1.0 x 10-14 [OH-] = 1.0 x 10-14 1.0 x 10-2 = 1.0 x 10-12 When [H+] > [OH-], acidic solution When [H+] < [OH-], basic solution or alkaline solution

  15. Because these concentrations are very small, another scale, the pH scale is used to describe [H+], and pOHis used to describe [OH-] • pH = -log[H+] • In a neutral solution, [H+] = 1.0 x 10-7, so pH = - log (1.0 x 10-7) = -(0 + -7) = 7 • pH < 7 = acidic solution • pH = 7 = neutral solution • pH > 7 = basic solution

  16. pOH = - log [OH-] • pH + pOH = 14 • To calculate pH, use – log [H+] • To calculate pOH, can use 14 – pH or – log [OH-] • To calculate the concentration from a pH value, use the antilog or 10x button

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