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Acids and Bases in all Different Places

Acids and Bases in all Different Places. I. Properties of Acids. A. Molecular substances which ionize when added to water to form hydronium (H 3 O +1 ) ions  all acids are electrolytes B . React with active metals to form H 2(g ) 1. _ Mg (s) + _ HCl (aq) 

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Acids and Bases in all Different Places

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  1. Acids and Bases in all Different Places

  2. I. Properties of Acids • A. Molecular substances which ionize when added to water to form hydronium (H3O+1) ions  all acids are electrolytes • B. React with active metals to form H2(g) • 1. _ Mg(s) + _ HCl(aq) • 2. _ Zn(s) + _ HCl(aq) • 3. _ Cu(s) + _ HCl(aq)

  3. I. Properties of Acids • C. Acids affect the colors of indicators • Universal Indicator • Phenolphthalein • D. Acids neutralizebases • E. Dilute acids taste sour think citric acid and ascorbic acid (Vitamin C) • **SAFETY TIP: Acids release tremendous amounts of heat when you dilute them • (esp. H2SO4)  ALWAYS ADD ACID TO WATER

  4. II. Naming Acids • treated as an ionic compound with H+1 (hydrogen ion) as cation • negative ion can be nonmetal (binary acid) or polyatomic anion (ternary acid) • A) Binary acids – acids that contain a negative ion ending in “-ide” • 1) Formula Name • use prefix: hydro- • use root of anion’s name • use suffix: -ic • a) HCl b) HBr c) HF • hydrochloric acid hydrobromicacid hydrofluoric acid

  5. II. Naming Acids • 2) Name  Formula • follow above rules in reverse • be sure to balance charges • a) hydroiodic acid b) hydrosulfuric acid • HI H2S

  6. II. Naming Acids • B) Ternary acids • DO NOT BEGIN WITH “hydro-“!!!!!!! • use name of polyatomic ion and switch its ending: • NOTE: sulfur stays “sulfur-” + ending, phosphorus stays “phosphor-” + ending

  7. II. Naming Acids • 1) Formula Name • a) H2CO3 • CO3 carbonATE  carbonic acid • b) H2SO4 • SO4 sulfATE  sulfuric acid • 2) Name Formula • a) acetic acid • acetIC  acetATE  HC2H3O2 • b) phosphoric acid • phosphorIC  phosphATE  H3PO4 • c) nitric acid • nitrIC nitrATE  HNO3

  8. ** some acids are stronger than others: Rank the following acids from weakest to strongest: sulfuric acid, carbonic acid, hydrochloric acid, hydrofluoric acid, acetic acid H2CO3 , HC2H3O2 , HF, H2SO4 , HCl

  9. III. Bases - ionic substance which dissociates to form hydroxide (OH-1) ions in water * examples: lye (NaOH) , lime (Ca(OH)2) , milk of magnesia (Mg(OH)2) • Naming Review. Name (or give the formula for) the following bases: • 1. NaOH • sodium hydroxide • 2. Mg(OH)2 • magnesium hydroxide • 3. aluminum hydroxide • Al(OH)3 • 4. ammonium hydroxide • NH4OH

  10. IV. Properties of Bases - often referred to as caustic or alkaline substances • A. Bases are electrolytes - dissociate in water to form OH-. • B. Bases affect the colors of indicators. • Universal IndicatorPURPLE • PhenolphthaleinMAGENTA • C. Bases neutralize acids. • D. Water solutions are bitter and slippery. • E. Emulsify fats and oilsthis is why they are useful in soap

  11. V. Salt – any ionic compound that does not contain hydroxide (OH-1) • * all are good electrolytes

  12. formed by a neutralization reactionAcid + Base  Salt + Water • 1) _____ HCl(aq) + _____ NaOH(aq) • 2) _____ H2SO4(aq) + ____ KOH(aq) • 3) ____ HBr(aq) + _____ Ca(OH)2(aq) • 4) _____ HC2H3O2(aq) + _____ NaOH(aq)

  13. Acid, Base, Salt, or Neither: • 1. NaCl 2. KCl 3. KOH 4. SO2 5. NH4C2H3O2 • Salt salt baseneithersalt • 1. KBr 2. H2SO4 3. HgCl2 4. Al(OH)3 5. HCl • Salt acid salt base acid • 6. KOH 7. CaO 8. K3PO4 9. CO2 10. NH4OH • Base saltsalt neither base

  14. VI. pH – a mathematical way of measuring how acidic a solution is

  15. It’s a logarithmic scale; that means each step is worth 10 • lemon juice is 10 times more acidic than vinegar • battery acid is 10 times more acidic than lemon juice • How many times more acidic is battery acid than vinegar?

  16. Red Orange Green Blue Purple pH: 3 5 7 9 11 Color scale for Universal Indicator:

  17. Which of the solutions above is the most acidic? • Battery acid • 2) Which of the solutions above is the most basic? • Lye • 3) Look at the solutions that your teacher is testing with universal indicator. • Label each as acidic, basic, or neutral • Estimate the pH based on the color • Rank the substances from most acidic to least acidic

  18. RANK:

  19. VII. Buffer - a solution which is able to resist major changes in pH • example: HC2H3O2(aq) H+1(aq) + C2H3O2-1(aq) • common-ion effect - by adding a salt with the negative ion (NaC2H3O2, KC2H3O2), we increase the concentration of that ion, therefore: • add H+1: • the acid will react with the acetate ion to produce molecular acetic acid, thus “neutralizing” it and keeping the pH the same • add OH-1: • the base will react with the molecular acetic acid to produceacetate ions, thus “neutralizing” it and keeping the pH the same

  20. BLOODY BUFFERS!! • biological example: carbonic acid/bicarbonate in blood  Hold your Breath!!! • There is a balance between the ions which acts as a buffer, keeping the pH of the blood right around 7.4. The hemoglobin molecule in red blood cells can only withstand pH extremes of 7.2-7.6

  21. VIII. Acid-Base Indicators - chemicals specifically designed to show specific colors in acids and different colors in bases

  22. IX. Acid-Base Neutralization H+1 + OH-1 H2O • if you have 35 molecules of acid, 35 molecules of base will neutralize it • equivalence point - when an equivalent amount of OH-1 ions has been added to H+1 ions  it’s “neutralized”

  23. X. Acid-Base Titration - lab procedure used to determine the concentration of an unknown acid or base solution. • standard solution – solution whose concentration is known • unknown solution – solution whose concentration you are trying to determine • MaVa = MbVb

  24. Titration Problems • 1) If you begin a titration with 20.0 mL of unknown HCl and titrate it to the equivalence point using 35.6 mL of 0.600 M standard NaOH, what is the concentration of HCl? • Ma(20.0 mL) = (0.600M)(35.6 mL) Ma= 1.07 M • 2) If you titrate 65.0 mL of an unknown NH3 solution to the equivalence point with 31.2 mL of a 1.50 M HCl solution, what is the concentration of the ammonia? • (1.50M)(31.2mL) = Mb(65.0mL) Mb= 0.720 M

  25. Titration Problems • 1) Ma = ??? Va = 50.0 mL Mb= 1.50 M Vb = 71.3 mL • Ma(50.0mL) = (1.50M)(71.3mL)Ma= 2.14 M • 2) What is the concentration of an unknown NaOH solution if you titrate 100.0 mL of it to the equivalence point with 43.5 mL of 6.0 M HCl? • (6.0M)(43.5mL) = Mb(100.0mL) Mb= 2.6 M

  26. Titration Problems • 3) What is the concentration of a vinegar (HC2H3O2) solution if you titrate exactly 20 drops of it to the equivalence point with 26 drops of 0.600M NaOH? • Ma(2θdr) = (0.600M)(26dr) Ma = 0.78 M

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