9. Acids And Bases Flashcards
Arrhenius acid base theory
-Svante Arrhenius (1884)
-ACID: substance that produces H3O+ ions in aqueous solution
-BASE: substance that produces OH- ions in aqueous solution
-still valid today for aqueous solution
-when acid dissolves in water, it reacts with water to produce H3O+ (i.e. HCl)-
-when bases (usually metal hydroxides, KOH, NaOH, Mg(OH)2) dissolve in water, it merely separate into ions
-other bases are not hydroxides; they produce OH- by reacting with water molecules
-i.e. NH3 + H2O <—-> NH4+ + OH-
Strong acid
-An acid that ionizes completely in aqueous solution
-strong electrolyte
HCl hydrochloric acid
HBr hydrobromic acid
HI hydroiodic acid
HNO3 nitric acid
H2SO4 sulfuric acid
HClO4 perchloric acid
Strong base
-ionizes completely in aqueous solution
-strong electrolyte
-LiOH, NaOH, KOH, Ba(OH)2
Weak acids
-an acid that is only partially ionized in aqueous solution
-produces much smaller amount of H3O+
-weak electrolyte
Glacial acetic acid
-pure acetic acid
-melting point 17°C, freezes on a moderately cold day
Lye
NaOH
-manufacture of glass and soap
Milk of magnesia
-suspension of about 8% Mg(OH)2 in water
-laxative
-treat wastewater in metal-processing plants
-flame retardant in plastic
Weak base
-base that is only partially ionized in aqueous solution
-weak electrolyte
Strength of an acid is NOT related to its concentration
Bronsted-Lowry theory
-Johannes Bronsted and Thomas Lowry (1923)
-ACID: proton donor
-BASE: proton acceptor
-Acid-base reaction is a proton transfer reaction
Conjugate acid-base pair
-a pair of molecules or ions that are related to one another by the gain or loss of a proton
Conjugate base
-a substance formed when an acid donates a proton to another molecule or ion
Conjugate acid
-substance formed when a base accepts a proton
Additional info for Bronsted-Lowry theory:
- Acid can be positively charged, neutral, or negatively charged.
(H3O, H2CO3, H2PO4-) - Base can be negatively charged or neutral (PO4^(3-), NH3L)
- Acids may be monoprotic, diprotic, triprotic, depending on number of protons each may give up
- Some substance can act as either an acid or base (amphiprotic).
-i.e. HCO3- can either be acid (becomes CO3^2-) or base (becomes H2CO3)
i.e. H2O (either H3O+ or OH-) - A substance cannot be Bronsted-Lowry acid unless it contains a hydrogen atom that can be given up.
-hydrogen to be given up must be bonded to a strongly electronegative atom, such as oxygen or halogen, to be acidic
-i.e. Acetic acid CH3COOH 4 hydrogen atom but monoprotic only - Inverse: strength of acid and conjugate base
-STRONGER acid, WEAKER conjugate base
How to determine if the position of equilibrium lies toward left or right in acid-base reaction
- Identify two acids and two bases in equilibrium
- Determine which acid and base are stronger
-stronger acid gives weaker conjugate base vice versa - The stronger acid and base react to give the weaker acid and weaker base
-Position of equilibrium lies on the SIDE OF WEAKER ACID AND BASE
Acid ionization constant (Ka)
aka
Acid dissociation constant
-equilibrium constant for the ionization of an acid in aqueous solution to H3O+ and its conjugate base
-can be used to quantify just how strong any weak acid is; weaker acid smaller Ka
pKa = -logKa
Properties of Acids and Bases
1. Neutralization
-acids and bases react with each other in a process called NEUTRALIZATION
-when a strong corrosive acid reacts with strong corrosive base, the product is a neutral solution
Properties of Acids and Bases
2. Reaction with metals
-strong acids react with ACTIVE METALS to produce hydrogen gas,H2 and a salt
-redox reaction
i.e. Mg + 2HCl ➡️ MgCl2 + H2
Active metals
-Alkali metals
-Alkaline earth metals
Properties of Acids and Bases
3. Reaction with Metal Hydroxides
-acids react with metal hydroxide to give salt and water
-both acid and metal hydroxide are ionized in aqueous solution
HCl + KOH ➡️ H2O + KCl
Properties of Acids and Bases
4. Reaction with metal oxides
-strong acid reacts with metal oxides to give water and salt
2H3O+ + CaO ➡️ 3H2O + Ca2+
Properties of Acids and Bases
5. Reactions with Carbonates and Bicarbonates
-strong acid and carbonate, carbon dioxide bubbles are rapidly given off
1st reaction: carbonate ion reacts with H3O to give carbonic acid
2nd reaction: carbonic acid immediately decomposes into carbon dioxide and water
Overall rxn
2 H3O+ + CO3^2- ➡️ CO2 + 3 H2O
-strong acid also react with bicarbonates to give carbon dioxide and water
• Any acid stronger than carbonic acid will react with carbonate or bicarbonate ion to give CO2 gas
Production of CO2
-earliest method: thru fermentation using yeast pero mabagal
-bakers use reaction of NaHCO3 (sodium bicarbonate or baking soda)
and weak acid
-BAKING POWDER contains weak acid (Na/KH2PO4) and Na/KHCO3
•do not react when dry
•when mixed with water in dough, react rapidly to produce CO2
Properties of Acids and Bases
5. Reaction with Ammonia and Amines
-any acid stronger than NH4+ is strong enough to react with NH3 to form salt
HCl + NH3 ➡️ NH4+ + Cl
NH4Cl is a salt
-base strength of amine is similar to NH3. So same rule, acid stronger than NH4+ will react with NH3 to form salt
HCl + CH3NH2 ➡️ CH3NH3+ + Cl-
Amines
-compounds similar to ammonia but one or more hydrogen atoms are replaced by carbon groups
i.e. Methylamine CH3NH2
Kw
aka
water constant
-ion product of water
-important because it applies not only to pure water but also to any water solution
Kw = 1 x 10^(-14)
Kw = k[H2O]^2 = [H3O+][OH-]
[H3O+] and [OH-] in aqueous solutions
⬆️ [H3O+] ⬆️ Acidic
⬆️ [OH-] ⬆️ Basic
[H3O+] 🟰 [OH-] Neutral
pH and pOH
-hydronium ion concentrations are numbers with negative exponents, it is more convenient to express in terms of pH
pH = -log[H3O+]
-acidic if pH < 7
-basic if pH > 7
-neutral if pH = 7
pOH = -log[OH-]
14 = pH + pOH
Measuring pH of an aqueous solution
- Use of pH paper (contains pH indicators that change color at certain pH)
-methyl orange (indicator): turns red when pH is 3.2 or lower (super acidic);
turns yellow when pH is 4.4 or higher - pH meter- dip the electrode of pH meter into the solution and read pH on a dial
Titration
-analytical procedure whereby we react a known volume of a solution of known concentration with known volume of a solution of unknown concentration
Equivalence point
-point at which there is an equal amount of acid and base in a neutralization reaction
End point
-pH indicator changes color
-mauuna muna si equivalence point bago end point since need yung unang excess ng hydronium ion/ hydroxide ion para madetect ng indicator yung acidity/ basicity
Phenolphthalein
-colorless in acid solution
-pink in basic solution
Buffer
-solution that resists change in pH when limited amounts of an acid or base are added to it
-aqueous solution containing weak acid and its conjugate base
-pH buffer is an acid/ base “shock absorber”
•kaya name-maintain ung pH level ng blood kahit kumain ng acidic foods
Buffer pH
-if buffer solution by mixing equimolar concentrations of any weak acid and its conjugate base, the pH of the solution will equal the pKa of the weak a id
Buffer capacity
-the extent to which a buffer solution can prevent a significant change in pH of a solution upon addition of a strong acid or a strong base
-the closer the pH of the buffer to pKa of the weak acid, the greater the buffer capacity
•effective buffer pH = pKa
- the greater the concentration of the weak acid and its conjugate base, the greater the buffer capacity
•buffer capacity of 1 mol each of CH3COOH and CH3COONa is greater than 0.1 mol of each;
•addition of 0.2 mol HCl in former results to pH 4.57 while addition of 0.2 mol HCl in latter results to pH 1.0
Blood buffers
-average blood pH=7.4; pH<6.8 or pH>7.8 result to death
-body uses 3 buffer system
•carbonate (most important), phosphate, proteins
Henderson-Hasselbach equation
-mathematical relationship between pH, pKa, and concentration of weak acid and its conjugate base
pH = pKa - log([acid]/[base])
Acidosis
-when pH of blood goes lower than 7.35-7.45 range
-caused by difficulty in breathing (hypoventilation); starvation; heavy exercise (excessive amount of lactic acid)
Alkalosis
-when pH of blood goes higher than 7.35-7.45 range
-results from rapid, heavy breathing (hyperventilation)
-sprinter’s trick: hyperventilating right before start; can absorb more lactic acid before the blood pH drops to the point where performance is impaired
Buffer with strange names (TRIS, HEPES, etc)
-limitation of previous buffers: too much pH change when solution is diluted or temperature changes
-N.E. Good
•developed buffers that consist of ZWITTERIONS (molecules with both positive and negative charges)
Oxoacids
Acids that contain hydrogen, oxygen, and another element (central element)
