Acids and Bases

Properties of Acids Aqueous solution have sour taste Change the color of acid / base indicators React with active metals to release H 2 gas 2HCl + Mg  MgCl 2 + H 2

Properties of Acids React with bases to produce salts and H 2 O (neutralization) Conduct electricity - electrolytes

Acid Nomenclature Binary Acid - contain H and 1 other element. Prefix “hydro” Root “name of 2nd element” Suffix “ic” Ex. Hydrochloric Acid HCl Hydroiodic Acid - HI

Oxyacid H, O and a 3rd element (mostly non-metal) Root - name the oxyion, replace suffix Suffix “ic” for “ate” “ous” for “ite” ions Carbonic Acid H 2 CO 3 Sulfuric Acid H 2 SO 4 2 H + + SO 4 2- sulfate Sulfurous Acid H 2 SO 3 2 H + + SO 3 2- sulfite

Common household substances that contain acids and bases. Vinegar is a dilute solution of acetic acid. Drain cleaners contain strong bases such as sodium hydroxide.

Common Industrial Acids H 2 SO 4 - sulfuric Acid - most produced chemical in the world. Metallurgy, fertilizer, paper, paint, dyes HNO 3 nitric acid- suffocating odor, stains, protein yellow - explosives, rubber, plastics

Common Industrial Acids H 3 PO 4 - phosphoric acid - fertilizer, flavors beverages, cleaning agent HCl - hydrochloric acid - found in stomach, cleaning agent, food processing, dilute forms called “muriatic acid” in stores.

Common Industrial Acids Acetic Acid- CH 3 COOH - foul smelling, vinegar - 4% - 6% acetic acid : plastics, food additives.

Bases Aqueous solutions taste bitter Change color of acid/ base indicators Dilute aqueous solutions feel slippery

Bases React with acids to produce salts and water (neutralization) Conduct electricity - electrolyte

Models of Acids and Bases Arrhenius Concept: Acids produce H + in solution, bases produce OH  ion. Brønsted-Lowry: Acids are H + donors, bases are proton acceptors. HCl + H 2 O  Cl  + H 3 O + acid base

Arrhenius Definition Acid – produces H + in water HCl  H + + Cl - H 2 SO 4  2H + + SO 4 2- Base – produces OH - in water NaOH  Na + + OH - NH 4 OH  NH OH -

Strong Acid ionizes completely in water - strong electrolyte. Ex. H 2 SO 4, HClO 4, HCl, HNO 3 HBr, HI

Weak Acid partially ionizes - weak electrolyte. HCN + H 2 O  H 3 O + + CN – Ex. H 3 PO 4, HF, CH 3 COOH, H 2 CO 3, H 2 S Answer questions on strong or weak (under the picture) !

Alkaline Arrhenius bases- increase concentration of OH- in aqueous solution. NaOH  Na + + OH – NH 3 + H 2 O  NH OH -

Strong Base strong electrolyte - completely dissociate. Ex. NaOH, KOH, LiOH

Weak Base Weak electrolyte- partially ionizes; Most of it stays in its molecular form Ex. NH 3 + H 2 O ↔ NH OH -

Acid/Base Strength Strong Acids HNO 3 HClO 4 HCl HBr HI H 2 SO 4 Strong Bases NaOH KOH LiOH Ba(OH) 2 Ca(OH) 2 (slightly soluble) Sr(OH) 2 (slightly soluble)

Bronsted - Lowry Acids and Bases Based on whether a substance is a proton acceptor or donor in non-aqueous solutions. HCl + NH 3  NH Cl - HCl donated a proton(H +) to NH 3. Proton donor - acid The NH 3 accepted a proton from HCl proton acceptor- base

Bronsted - Lowry Acids and Bases Bronsted-Lowry is a way to study proton transfer!! Acid – H+ donor Base – H+ acceptor For Example: HCl + NH 3  NH Cl - H 2 SO 4 + H 2 O  2 H 3 O + + SO 4 2-

Conjugate Acids/Bases Conjugate Acid – formed when BL base gains a proton Conjugate Base – formed when BL acid looses a proton HCl + NH 3  NH Cl - H 2 SO 4 + 2H 2 O  2 H 3 O + + SO 4 2-

Conjugate Acids/Bases HCl + NH 3 <  NH Cl - H 2 SO 4 +2H 2 O <  2 H 3 O + + SO 4 2- AcidBaseConj. Acid Conj. Base

Conjugates Strength The stronger the acid, the weaker its conjugate base; the stronger the base, the weaker its conjugate acid. Proton transfers favor the production of weaker acids and weaker base. Therefore, CH 3 COOH + Water  H 3 O + + CH 3 COO - (weak acid) (weak base)(stronger acid) (stronger base) Reactants are favored!!

Amphoteric can react as either an acid or base HCl + Water  H 3 O + + Cl - proton acceptor(water) Water + NH 3  NH OH - Water(proton donor)

Monoprotic Acids Ionization – when ions are formed from solute molecules by the action of the solvent. Monoprotic donates 1 hydrogen For example: H 2 O (l) + HCl (s)  H 3 O + (aq) + Cl - (aq) Hydronium ion = H +

Polyprotic Acids/Bases Some acids have more than one ionizable hydrogen and are called polyprotic: diprotic (2 H + ), triprotic (3 H + ). For example: 2 H 2 O (l) + H 2 SO 4 (s)  2 H 3 O + (aq) + SO 4 2- (aq) Two moles of hydronium ions

Polyprotic Acids Ionization is in several distinct steps: e.g., H 2 CO 3 : carbonic acid H 2 CO 3 + H 2 O  H 3 O + + HCO 3 - HCO H 2 O  H 3 O + + CO 3 2- Transfer of 2 nd (or 3 rd ) proton are more difficult than 1 st.

Lewis Acids and Bases based on bonding structures - includes acids that do not contain H. Broadest definition of acid/base Includes all BL acids and bases

Lewis Acids and Bases Lewis Acid: electron pair acceptor Lewis Base: electron pair donor

Lewis Acid an atom, ion or molecule that accepts an electron pair to form a covalent bond. BF 3 + F -  BF 4 (Lewis acid) (Lewis Base)

Lewis Base an atom, ion or molecule that donates an electron pair to form a covalent bond.

Self-ionization of water a Very weak electrolyte. Autoionization of water: 2 H 2 O (l)  H 3 O + (aq) + OH - (aq) hydronium hydroxide When protons (H+) are produced in water, they bind to the lone pair e- of water to produce H 3 O +

Acid/Base Equilibria At 25 o C pure water, has a pH = 7 [H 3 O + ] = [OH - ] = 1 x K w = ionization constant for water K w = 1.0 X o C Note: your book presents the autoionization of water on the reaction: H 2 O  H + + OH - ; [H + ] is analogous to [H 3 O + ] !

Acids and Bases H 2 O = HOH = H + + OH - AcidsBases HClNaOH HNO 3 KOH HFNH 4 OH

The pH Scale pH   log[H + ] pH in water ranges from 0 to 14. K w = 1.00  10  14 = [H + ] [OH  ] pH + pOH = As pH rises, pOH falls (sum = 14.00).

pH Scale pH – related to the concentration of H + ions in solutions. The more H + ions, the lower the pH.

The pH scale and pH values of some common substances.

HIn --> H+ + In - For phenolphthalein: pH 0 to 8.2 = colorless; then pink, then pH 10 = red Add H+ then shift to left Add OH- then shift to right

pH Practice 151x x x x x x x10 -4 A1x A21x10 -2 AorB[H 3 O + ]pHAorBpH[H 3 O + ] pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

pH Practice B1x B141x B1x B121x B1x B101x B1x B81x10 -8 N1x N7 A1x A61x10 -6 A1x A41x10 -4 A1x A21x10 -2 AorB[H 3 O + ]pHAorBpH[H 3 O + ] pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

More pH Practice 4.60x x x x x A41.00x10 -4 AorBpH[H 3 O + ] pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

More pH Practice B x10 -8 A x10 -2 B x A x10 -5 B x A41.00x10 -4 ANBpH[H 3 O + ] pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

More pH Practice 4.6bananas 7.8eggs 2.0stomach acid 8.5seawater 10.5milk of mag. 3.1apples N4.0x blood AorB[H 3 O + ]pHItem pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

More pH Practice A2.5x bananas B1.6x eggs A1.0x stomach acid B3.2x seawater B3.2x milk of mag. A7.9x apples N4.0x blood AorB[H 3 O + ]pHItem pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

Kw and pOH Practice 4.0x x bananas 6.25x x eggs 1.0x x stomach acid 3.1x x seawater 3.1x x milk of mag. 1.27x x apples 2.5x x blood [OH - ][H 3 O + ]pHItem pOH = -log [OH - ]; Kw = [H 3 O + ][OH - ]; Kw = [H 3 O + ][OH - ] = 1 x

Titrations Titration – experimental technique that provides a sensitive means of determining the chemically equivalent amounts of acid and base. Titration equation: How to titrate !! Vacid x (Molarity Acid x # Equivalent acid) = Vbase x (Molarity base x # Eq base)

Titrations Equivalence point – point in reaction when equal # moles of acid and base have reacted. Neutralization reaction equation: HCl + NaOH  NaCl + H 2 O 1 mol acidbasesaltwater

Titrations We can now look at titrations more quantitatively. Possible combinations we will consider: strong acid (SA) - strong base (SB) weak acid (WA) - strong base (WB) strong acid (SA) - weak base (WB) Consider each case before/at/after equivalence point (EP).

Titrations A Titration Curve shows pH at various points in a titration experiment(before, at, and after equivalence point (EP)). Can generate by: experimentally measuring pH during a titration experiment calculating pH at various points for an acid-base reaction

Titrations The Titration curve of acids with bases (or visa versa) has four different environments to consider: 1.Before rxn (before titration begins) 2.Before the EP (at the very beginning of the titration) 3.At the EP (moles of acids = moles of base) 4.After the EP

SA/SB Titrations KOH (aq) + HBr (aq)  KBr (aq) + H 2 O (l) Net reaction: OH - (aq) + H + (aq)  H 2 Equivalence Point (EP): What compounds are present at EP? H 2 O, K +, and Br - K + and Br - are conjugates of strong base and acid; too weak to react with EP, [OH - ] = [H 3 O + ]; pH=7

SA/SB Titrations The Titration Curve for a SA/SB will look like this. But: Acid/Base “neutralization” reactions do not always result in “neutral” solutions!

SA/SB Titrations Before & After Equivalence Point (EP) In this case, we need to consider what species are present and at what concentrations. At each stage – we need the [H 3 O + ] or [OH - ] concentration! Example: KOH + HBr Have 20.0 ml 2.0 M HBr; titrate with 1.0 M KOH

SA/SB Titrations 20.0 ml 2.0 M HBr; titrate 1.0 M KOH Before rxn: 2.0 M HBr = 2.0 M H 3 O + pH = ~ L (2.0 M) = mol HBr present M A V A = M B V B (2.0 M)(0.02 L) = (1.0 M)(vol KOH) EP will occur at 0.04 L or 40 ml KOH

SA/SB Titrations 20.0 ml 2.0 M HBr; titrate 1.0 M KOH Before Equivalence Point (30.0 ml KOH): 30.0 ml KOH = L = mol KOH added mol HBr have reacted = mol HBr in 50.0 ml  0.20 M H 3 O + pH = +0.70

SA/SB Titrations 20.0 ml 2.0 M HBr; titrate 1.0 M Equivalence Point (40.0 ml KOH): 40.0 ml KOH = L = mol KOH added mol HBr have reacted = mol HBr in 60.0 ml  [OH - ] = [H 3 O + ]; pH = 7

SA/SB Titrations 20.0 ml 2.0 M HBr; titrate 1.0 M KOH After Equivalence Point (50.0 ml KOH): 50.0 ml KOH = L = mol KOH added all HBr reacted; mol KOH leftover! mol KOH in 70.0 ml  [OH - ] = M OH - pOH = 0.84; pH = 13.15

WA/SB Titrations F - + H 2 O  HF + OH - This equilibrium will make the solution basic. The EP of a WA/SB reaction is always basic.

SA/WB Titrations NH H 2 O  H 3 O + + NH 3 This equilibrium will make the solution acidic. The EP of a SA/WB reaction is always acidic.

Titration Summary Strong Acid/ Strong Base

Titration Summary Weak Acid/ Strong Base

Titration Summary Strong Acid/ Weak Base

Titration Summary

Click on Titration curve Then, you pick the correct Curve in the question

Polyprotic Acid Titrations When polyprotic acids are titrated with strong bases, there are multiple equivalence points. The titration curve of a polyprotic acid shows an equivalence point for the each acid H + :

Acid-Base Properties of Salts