Unit 17- Acids/Bases/Salts

General properties

• Taste sour• Turn litmus• React with active

metals

• React with bases

• Taste bitter• Turn litmus• Feel soapy or slippery

(react with fats to make soap)

• React with acids

blue to red red to blue

Acids have a

pH less

than 7

Acids React with Carbonates and Bicarbonates

HCl + NaHCO3

NaCl + H2O + CO2

Hydrochloric acid + sodium bicarbonate

salt + water + carbon dioxide

An old-time home remedy for relieving an upset

stomach

Effects of Acid Rain on Marble(marble is calcium carbonate)

George Washington:BEFORE acid rain

George Washington:AFTER acid rain

Bases Neutralize Acids

Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl.

2 HCl + Mg(OH)2

MgCl2 + 2 H2O

Magnesium salts can cause diarrhea (thus they are used as a laxative) and may also cause kidney stones.

Naming

• Acids are composed of hydrogen (H+) followed by an anion (negative ion).

• If the acid formula contains oxygen in the anion, such as in H2SO4, it is known as an oxyacid.

3 Rules To Naming Acids• If H + anion ending in –ide: Acid name is “hydro_____ic

acid” – Take the root from the anion name and fill in the blank.

• H + anion ending in –ate: Acid name is “_____ic acid” – Take the root from the anion name and fill in the blank.– “What I ATE was ICky”

• H + anion ending in –ite: Acid name is “_____ous acid”– Take the root from the anion name and fill in the blank.– “Don’t bITE; it’s infectiOUS”

Exceptions

• Sulfate (SO4 2-)

– Root is not sulf, but sulfur• Sulfuric acid

• Phosphate (PO4 3-)

– Root is not phosph, but phosphor• Phosphoric acid

Svante Arrhenius• He was a Swedish chemist (1859-

1927), and a Nobel prize winner in chemistry (1903)

• one of the first chemists to explain the chemical theory of the behavior of acids and bases

Arrhenius Definition - 1887• Acids produce hydrogen ions (H1+)

in aqueous solution (HCl → H1+ + Cl1-)

• Bases produce hydroxide ions (OH1-) when dissolved in water.

(NaOH → Na1+ + OH1-)

• Limited to aqueous solutions.• Only one kind of base (hydroxides)• NH3 (ammonia) could not be an

Arrhenius base: no OH1- produced.

Brønsted-Lowry - 1923• A broader definition than Arrhenius

• Acid is hydrogen-ion donor (H+ or proton); base is hydrogen-ion acceptor.

• Acids and bases always come in pairs.

• HCl is an acid.

– When it dissolves in water, it gives it’s proton to water.

HCl(g) + H2O(l) ↔ H3O+(aq) + Cl-(aq)

• Water is a base; makes hydronium ion.

Acids and bases come in pairs

• A “conjugate base” is the remainder of the original acid, after it donates it’s hydrogen ion

• A “conjugate acid” is the particle formed when the original base gains a hydrogen ion

• Thus, a conjugate acid-base pair is related by the loss or gain of a single hydrogen ion.

Definitions

• Acids – produce H+

• Bases - produce OH-

• Acids – donate H+

• Bases – accept H+

Arrehenius

Bronsted-Lowry

only in water

any solvent

When life goes either wayAmphiprotic substances

HCO3-

H2CO3 CO3-2

+ H+ - H+

Acting like a base

Acting like an acid

accepts H+ donates H+

Strong and Weak Acids/BasesStrong acids/bases – 100% dissociation into ions

HCl NaOHHNO3 KOHH2SO4

Weak acids/bases – partial dissociation, both ions and molecules

CH3COOH NH3

http://www.mhhe.com/physsci/chemistry/animations/chang_2e/acid_ionization.swf

Strong Acids

• Prechloric Acid HClO4

• Sulfuric Acid H2SO4

• Hydriodic Acid HI

• Hydrobromic Acid HBr

• Hydrochloric Acid HCl

• Nitric Acid HNO3

Strong Bases• Lithium hydroxide, LiOH• Calcium hydroxide,Ca(OH)2

• Sodium hydroxide, NaOH• Strontium hydroxide,Sr(OH)2

• Potassium hydroxide,KOH

•Barium Barium hydroxide,hydroxide,Ba(OH)Ba(OH)22

•Magnesium Magnesium hydroxide,hydroxide,Mg(OH)Mg(OH)22

Salt Hydrolysis• A salt is an ionic compound that:

–comes from the anion of an acid

–comes from the cation of a base

–is formed from a neutralization reaction

–some neutral; others acidic or basic

• “Salt hydrolysis” - a salt that reacts with water to produce an acid or base

Salt Hydrolysis• Hydrolyzing salts usually come from:

1. a strong acid + a weak base, or2. a weak acid + a strong base

• Strong refers to the degree of ionization

A strong Acid + a strong Base = Neutral Salt

• How do you know if it’s strong?– Refer the list on your notes

Salt Hydrolysis• To see if the resulting salt is

acidic or basic, check the “parent” acid and base that formed it. Practice on these:

HCl + NaOH H2SO4 + NH4OH

CH3COOH + KOH

NaCl, a neutral salt

(NH4)2SO4, acidic salt

CH3COOK, basic salt

Ionization constant of Water• Water ionizes, or falls apart into ions:

H2O ↔ H1+ + OH1-

• Called the “self ionization” of water

• Occurs to a very small extent:

[H1+ ] = [OH1-] = 1 x 10-7 M

• Since they are equal, a neutral solution results from water

Kw = [H1+ ] x [OH1-] = 1 x 10-14 M2

• Kw is called the “ion product constant” for water

Ion Product Constant• Kw is constant in every aqueous solution:

[H+] x [OH-] = 1 x 10-14 M2

• If [H+] > 10-7 then [OH-] < 10-7

• If [H+] < 10-7 then [OH-] > 10-7

• If we know one, other can be determined

• If [H+] > 10-7 , it is acidic and [OH-] < 10-7

• If [H+] < 10-7 , it is basic and [OH-] > 10-7

– Basic solutions also called “alkaline”

• If [H3O+] > [OH–‑] the solution is acidic.

• If [H3O+] < [OH–‑] the solution is basic.

• If [H3O+] = [OH–‑] the solution is neutral.

Acid/Base/Salts

Part 2

pH

• Expressing hydrogen ion concentration in numbers can be cumbersome.

• A widely used system for expressing [H3O+] is the pH scale.

Whether or not a solution is acidic, basic, or neutral depends on the balance of H+ and OH- ions:

• Neutral: [H+] = [OH-]• Acid: [H+] > [OH-]• Base: [H+] < [OH-]

pH

• pH is the negative base 10 logarithm of the hydronium ion concentration:

pH = - log [H3O+]

Measuring pH• Why measure pH?

Everyday solutions we use - everything from swimming pools, soil conditions for plants, medical diagnosis, soaps and shampoos, etc.

• Sometimes we can use indicators, other times we might need a pH meter

pH

• Remember, for pure water, [H3O+] is

1 x 10-7 M. So what’s the pH of pure water?

• Ex: What is the pH of a solution with a hydronium ion concentration of 1.0 x 10-10 M?

Note that as pH increases, [H3O+] decreases and [OH–‑] increases.

Note the relationship between [H3O+] and [OH–‑]. Remember, the product of these must always equal 1 x 10 -14 for aqueous solutions.

Note that as pH increases, [H3O+] decreases and [OH–‑] increases.Note the relationship between [H3O+] and [OH–‑]. Remember, the product of these must always equal 1 x 10 -14 for aqueous solutions.

pOH = - log [OH-]

pH + pOH = 14

Example

• Find the pH, the pOH = 5.3pH + pOH = 14

pH = 14 - 5.3

pH = 8.7

pH calculations

• Use the reverse of the equation to calculate the [H+] when pH is known.[H3O+] =10(-pH)

*** 2nd log on the calulator!

= 10-pH

pH calculations

• Use identical process for [OH-] when pOH is known.[OH-] = 10(-pOH)

Acid/Bases/Salts

Neutralization/Titrations

A. Neutralization• Chemical reaction between an acid and a

base.• Products are a salt (ionic compound) and

water.• Note in the reaction above, the acid:base mole ratio is 1:1.

• However, in the reaction between H2SO4 and NaOH:

• it takes 2 moles of base to neutralize 1 mole of acid.

• The reacting ratios of acid and base will be important in solving problems related to neutralization reactions.

A. Neutralization

ACID + BASE SALT + WATER

HCl + NaOH NaCl + H2O

HC2H3O2 + NaOH NaC2H3O2 + H2O

– Salts can be neutral, acidic, or basic.

– Neutralization does not mean pH = 7.

weak

strong strong

strong

neutral

basic

B. Titration

• Titration– A laboratory technique

that uses a neutralization reaction to determine the concentration of an unknown acid or base.

standard solution

unknown solution

•An indicator is used to show when neutralization has occurred. •changes color in response to changes in pH. •actually weak acid or base that change color in response to pH change. •In acidic solutions, indicators act as Bronsted-Lowry base. As the indicator molecules accept H+, they change color.•In basic solutions, indicators act as Bronsted-Lowry acid. As the indicator molecules donate H+, they change color.

• End Point – – point at which an indicator changes

color during a titration

• Equivalence point – Point at which equal amounts of

H3O+ and OH- have been added.– when mole ratio exactly equals

mole ratio required by reaction – Determined by…

• indicator color change

B. Titration

• dramatic change in pH

B. Titration

moles H3O+ = moles OH-

MV n = MV n

M: MolarityV: volumen: # of H+ ions in the acid

or OH- ions in the base

B. Titration

• 42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H2SO4. Find the molarity of H2SO4.

H3O+

M = ?V = 50.0 mLn = 2

OH-

M = 1.3MV = 42.5 mLn = 1

MV# = MV#M(50.0mL)(2)=(1.3M)(42.5mL)(1)

M = 0.55M H2SO4