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Acids and Bases

Acids and Bases. Unit VII. I Electrolytes. An electrolyte is a compound, that when dissolved in water, conducts electricity How? Ions (charges) produced are free to move Movement of charge is conductivity Examples Acids Bases “Salts” Soluble Ionic compounds. II Properties. Acids

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Acids and Bases

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  1. Acids and Bases Unit VII

  2. I Electrolytes • An electrolyte is a compound, that when dissolved in water, conducts electricity • How? • Ions (charges) produced are free to move • Movement of charge is conductivity • Examples • Acids • Bases • “Salts” • Soluble Ionic compounds

  3. II Properties • Acids • Good conductors • Dissolve metals • Table J--Metals above “H2” dissolve in acid • Taste sour • Turns litmus paper red • Turns phenolphthalein clear

  4. II Properties • Bases • Good conductors • Dissolve fats • Feels slippery • Taste bitter • Turns litmus paper blue • Turns phenolphthalein pink

  5. III Definitions • Arrhenius • Acids • An Arrhenius acid contains H+ ions • When dissolved in water these H+ ions combine to form hydronium ion (H3O+) • Examples: HCl H2SO4 HC2H3O2 • Bases • An Arrhenius base contains OH- ions (hydroxide ion) bonded to NH4+ or a metal • Examples: NaOHCa(OH)2

  6. III Definitions B. Brönsted-Lowry • Acids • A Brönsted-Lowry acid loses or donates protons to its conjugate (substance that differs by an H+) • HCl + NH3→Cl- + NH4+ • HCl and Cl- are conjugate pairs; HCl is the acid and Cl- is its conjugate base • Bases • A Brönsted-Lowry base gains or accepts protons from its conjugate • HCl + NH3→Cl- + NH4+ • NH3 and NH4+ are conjugate pairs; NH3 is the base and NH4+ is its conjugate acid

  7. IV Nomenclature A. Naming Compounds • Binary Acids • A binary acid contains hydrogen and a nonmetal • To name a binary acid • Use “hydro-” • Add nonmetal root word • End with “ic acid” • Ex. HCl • Hydrochloric acid • Ex. H2O • Hydroxic acid

  8. IV Nomenclature • Ternary Acids • A ternary acid contains hydrogen and a polyatomic ion • To name a ternary acid • Determine the polyatomic that is present using Reference Table E • If the polyatomic ion ends in “ate” change the ending to “ic” • If the polyatomic ion ends in “ite” change the ending to “ous” • Ex. HClO3 • Chlorate becomes Chloric acid (no hydro is used) • Ex. HNO2 • Nitrite becomes Nitrous acid

  9. IV Nomenclature • Bases • To name a base, name as you would any compound • Write the first element • Write the polyatomic • Add a Roman numeral if needed • Ex. NaOH • Sodium hydroxide • Ex. Cu(OH) 2 • Copper II hydroxide

  10. IV Nomenclature • Writing Formulas - Acids • If binary • Write H+ and the other element present • Assign charges and criss-cross • Ex. Hydrochloric acid H+1 Cl-1 HCl

  11. IV Nomenclature • If ternary • Identify the polyatomic present using ending • Write H+ and the polyatomic ion • Assign charges and criss-cross • Ex. Chloric acid • chloric comes from chlorate ClO3-1 • H+1 ClO3-1 • HClO3

  12. V Reactions A Neutralization • Mixing of acid and base • HCl + NaOH → • Makes salt and water • Break (ionize) the acid and base • H+1 Cl-1 Na +1 OH-1 • Join H to OH (H2O) • Join metal to nonmetal (assign charges and crisscross) • HCl + NaOH → H2O + NaCl

  13. Lab technique for neutralization is called Titration • Occurs when moles of acid equals moles of base • For 1:1 acid–base reactions • Moles acid = Moles base • MAVA = MBVB

  14. Examples of titration problems Given the balanced equation: HCl + NaOH → H2O + NaCl How many milliliters of 3.0M NaOH are needed to neutralize 20 milliliters of 2.5M HCl? • MB= 3.0M NaOH • MA= 2.5M HCl • VA= 20 mLsHCl • Ratio is 1:1 so MAVA = MBVB can be used • 2.5M x 20 mLs = 3.0M x VB • 50 = 3VB • VB=16.7 mLsNaOH

  15. Examples of titration problems Given the balanced equation: H2SO4 + 2 NaOH → 2 H2O + Na2SO4 How many milliliters of 1.2 M NaOH are needed to neutralize 23 milliliters of 1.9 M H2SO4? • MB= 1.2 M NaOH • MA= 1.9 M H2SO4 • VA= 23 mLs H2SO4 • Ratio is NOT 1:1 so MAVA = MBVB CANNOT be used 23mL H2SO4 x 1L x 1.9 mole H2SO4 x 2 mole NaOH x 1L x 1000 ml 1000 mL 1 L 1 mole H2SO4 1.2 mole NaOH 1L VB= 72.8 mLs NaOH

  16. IV Reactions B Hydrolysis • Mixing of salt and water • Makes parent acid and base of the salt NaCl + HOH → HCl + NaOH • Reverse of neutralization

  17. VI Strength • Acid and base strength depend on number of ions in solution • More ions; stronger acid or base • Some acids ionize 100% (strongest acids) • HCl HBr HI • H2SO4 HNO3 HClO4 • Some bases ionize 100% (strongest bases) • LiOH NaOH KOH • RbOH CsOH NH4OH

  18. pH Every aqueous solution contains H + and OH- • Acids have more H+ than OH- • Bases have more OH- than H+ pH represents the amount of H + in a solution 1 7 14 Strong Weak Neutral Weak Strong acid acid base base Most H+ Equal H+ and OH- Least H+ Least OH- Most OH-

  19. pH Scale is logarithmic • Values change by factors of 10 • ex. pH = 3 vs. pH = 5 • Difference in pH • 2 units • 10 x 10 • pH 3 is 100 times stronger than pH 5 • pH 5 is 1/100th as strong as pH 3

  20. Acid Base Indicators • Compounds that change color over pH ranges • Table M Common Acid–Base Indicators • methyl orange 3.2–4.4 red to yellow • bromthymol blue 6.0–7.6 yellow to blue • phenolphthalein 8.2–10 colorless to pink • litmus 5.5–8.2 red to blue • bromcresol green 3.8–5.4 yellow to blue • thymol blue 8.0–9.6 yellow to blue

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