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Acids bases & salts

Acids bases & salts. Objectives. State the Bronsted-Lowry definition of acids and bases Identify the common physical and chemical properties of acids and bases Explain what dissociation constants indicate about an acid or base Use experimental data to calculate a dissociation constant.

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Acids bases & salts

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  1. Acids bases & salts

  2. Objectives • State the Bronsted-Lowry definition of acids and bases • Identify the common physical and chemical properties of acids and bases • Explain what dissociation constants indicate about an acid or base • Use experimental data to calculate a dissociation constant

  3. Properties of acids/bases • Taste – acid comes from latin meaning sour or tart/bases are bitter – soap • Touch most dilute acids feel like water although they sting on broken skin bases feel smooth, soothing and slippery except in your eyes (soap) • Reactions with metal acids react bases do not react

  4. Properties continued • Electrical conductivity water is poor – HCl is good, NaOH is good both are electrolytes • Indicators turn color – acid turns litmus paper from blue to red base turns from red to blue • Neutralization reaction between an acid and a base get salt and water (double replacement)

  5. Arrhenius definition • An acid is a substance that dissociates in water to produce hydrogen ions. • A base is a substance that dissociates in water to produce hydroxide ions • A salt is an ionic compound formed from any cation other that H+ and any anion other than OH- or O-2

  6. Arrhenius continued • Acids react with metals to produce H2 gas • Mg +2H+ -> Mg+2 + H2 • oxidation reduction reaction

  7. Bronsted – Lowry definitions • An acid is any substance that can donate H+ ions • A base is any substance that can accept H+ ions a Bronsted-Lowry acid is a proton donor and a base is a proton acceptor

  8. Hydronium Ion • H+ is strongly attracted to the electrons of surrounding water molecules • H+ + H2O -> H3O+ • More correct HCl + H2O -> H3O++ Cl- • In this case HCl is the Bronsted-Lowry acid and water is the base • We still describe a solution of HCl as acidic!

  9. NH3 + H2O -> NH4+ + OH-Ammonia is the H+ acceptor water is the H+ donor (acid) Amphoteric a substance that can act as either an acid or a base

  10. Conjugate Acid-Base Pairs • NH3 + H2O  NH4+ + OH- • In 1 direction water is the acid in the reverse reaction it is the base. • These cmpds become conjugate acids and conjugate bases when HCl loses an H+ ion to become its conjugate base Cl- when the conjugate base of water is the hydroxide ion OH- • When ammonia gains H+ to become its conjugate acid NH4+ and the conjugate acid of OH- is H2O

  11. Conjugate pairs • A pair of compounds that differ by only one H+ ion such as H2O and OH- or NH3 and NH4+ are called conjugate acid base pairs • NH3 + H2O  NH4+ + OH- • Base acid conj. Acid conj. base

  12. Determining the strengths of acids and bases • A strong acid HCl readily transfers hydrogen ions to water to form H3O+ • If you place 1M of HCl in 1 liter of water you would form 1 M H3O+ ions and 1 M Cl- ions

  13. weak acids • A weak acid does not readily transfer H+ ions • 1 mole of acetic acid in 1 liter of water only .4% of the acetic acid molecules would form H3O+ and C2H3O2-. Which means that 99.6% of the acetic acid molecules do not dissociate.

  14. To show a strong acid from a weak acid use arrows • HCl + H2O  H3O+ + Cl- • HC2H3O2 + H2O <-> H3O+ + C2H3O2-

  15. Strong and Weak Bases • The most widely used commercial base is CaO. When CaO is dissolved in water the O-2 ions react completely with H2O to form OH- ions. • O-2 + H2O  2OH- use a single arrow

  16. Strength of conjugate acid – base pairs • The stronger the acid the weaker its conjugate base. • The stronger the base the weaker its conjugate acid.

  17. The acid dissociation constant Ka • Weak acid HA • HA + H2O <--> H3O+ + A- • Keq = • For a 1 M solution of a typical weak acid may be only .007% of the water molecules react. Move the water to the left side of the equation • The higher the Ka the more the reaction goes to the right . The greater the Ka the stronger the acid • Weak acids have a Ka less than 1

  18. Diprotic acids • 2 step dissociation

  19. Base dissociation constant Kb • The base dissociation constant is a measure of the strength of a base • Do calculations of dissociation constants

  20. Objectives • Explain what most acidic hydrogen atoms have in common • Explain what most bases have in common • Describe how acids are named

  21. Naming and identifying acids and bases • 1st all H’s are not acidic CH4 • As a rule an acidic hydrogen already has a slight positive charge while is it part of a molecule. (It is in the positive side of a polar covalent bond) • Usually bonded with O, N, or a halogen

  22. 3 types of acids • Binary acids - H and 1 other element usually 6A or 7A • Strong HCl, HBr, and HI • Weak HF, H2S and H2Se • Oxy acids contain H, O and 1 other element • H2SO4, HNO3 and H3PO4 • Carboxylic Acid – organic acids COOH group • Acetic acid HC2H3O2 vinegar

  23. Bases • A Bronsted-Lowry base always contains an unshared pair of electrons NH3 attracts H+ • Anions: remember conjugate bases HCl Cl- • Weak Cl- Br- I- NO3- HSO4- CIO4- • Strong O-2 OH- PO4-3 and CO3-2 • Amines N has an unshared pr. Of electrons

  24. Naming Acids and Bases • If the name of an anion ends in ide the name of the acid that produces it includes the name of the anion, a hydo prefix • Hydrochloric acid – all binary acids • If the name of an anion ends in ate use and ic ending Nitric acid, carboxylic acid • Anion ends in ite SO3-2 sulfite ion H2SO3 sulfurous acid

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