Acids/Bases/Salts

Properties

Common Acids

Lactic sour milk

Acetic vinegar

Phosphoric tart taste in soda

Citric citrus fruits

Malic apples

Tartaric grapes

Formic ant bites

Common Base

Ammonia window cleaner

Sodium hydroxide lye (drain and oven cleaners)

Sodium Bicarbonate Baking soda

antacid

Milk of magnesia antacid

Properties

electrolytes

turn litmus red

sour taste

react with metals to form H2 gas

slippery feel

turn litmus blue

bitter taste

vinegar, soda, apples, citrus fruits

ammonia, lye, antacid, baking soda

electrolytes

pH less than 7 pH more than 7

Acid nomenclature –naming acids

Binary acids – is an acid that contains only two different elements.

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

Oxyacids- is an acid that is a compound of hydrogen, oxygen and third element (usually a non-metal)

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.

Acid Naming Example

Example: HCl • Cl is the anion, its name is chloride• Name of acid is: hydrochloric acid

Example: HF• F is the anion, its name is fluoride• Name of acid is: hydrofluoric acid

3 Rules To Naming Acids

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”

Acid Naming Example

Example: HNO3

NO3 1- is the anion, its name is

nitrate Name of acid is: nitric acid

Example: H2CO3

CO3 2- is the anion, its name is carbonate Name of acid is: carbonic acid

Exceptions

Sulfate (SO4 2-) Root is not sulf, but sulfur

• Sulfuric acid

Phosphate (PO4 3-) Root is not phosph, but phosphor

• Phosphoric acid

3 Rules To Naming Acids

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”

Acid Naming Example

Example: HNO2

NO2 1- is the anion, its name is nitrite Name of acid is: nitrous acid

Example: HClO2

ClO2 1- is the anion, its name is chlorite Name of acid is: chlorous acid

Strength of Acids

Strong Acid

- is one that ionizes completely in aqueous solution.

- completely dissociates.

- a strong acid is a strong electrolyte.

- increases with increasing polarity and decreasing bond energy.

Strength of Acids

Weak acids

- are weak electrolytes

- it contains hydronium ions, anions and dissolved acid molecules

- Organic acids (acidic carboxyl group) -COOH

Aqueous solutions for base

Most bases are ionic compounds containing metal cations and the hydroxide anion.

When a base completely dissociates in water to yield aqueous OH- ions, the solution is called alkaline.

NaOH (s) water Na+(aq) + OH-

(aq)

Aqueous solutions for base

Not all bases are ionic compounds.

Ammonia is one example because it produces hydroxide ions when it reacts with water molecules.

NH3 (g) + H20 (l) NH4+ (aq) + OH-

(aq)

Strength of bases

The strength of the base depends on the extent to which it dissociates to its ions.

Strong bases are strong electrolytes

Bases that are not very soluble do not produce a large number of hydroxide ions when added to water

Base Naming Example

NaOH name of base: sodium hydroxide

Mg(OH)2

name of base: magnesium hydroxide Fe(OH)2

name of base: iron (II) hydroxide

pH and pH scale

0

7INCREASING

ACIDITY NEUTRALINCREASING

BASICITY

14

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 and pOH Calculations

pH is the negative base 10 logarithm of the hydrogen ion concentration:

pH = - log10 [H+]

pH is the expression of the acidity or alkalinity of a solution in terms of its hydronium ion concentration.

pH = - log [H3O+]The sum of the pH and the pOH always equals 14.

pH + pOH = 14

Example

Calculate the pH, if the

[H3O] = 2.4 X 10-6 M

pH = - log [H3O+]

= - log(2.4 X 10-6)

= -(-5.6)

= 5.6 Find the pH, the pOH = 5.3

pH + pOH = 14

pH = 14 - 5.3

pH = 8.7

pH calculations

Use the reverse of the equation to calculate the [H3O+] when pH is known.

[H3O+] = antilog (-pH)

= 10-pH

Use an identical equation to calculate pOH. pOH = - log [OH-]

Calculate the [H3O+], if the pH is 4.71. [H3O+] = antilog (-pH)

= antilog (-4.71) = 1.95 X 10-5 M

Acid- Base Theories

Type Acid Base

Arrhenius(Traditional)

H+ or H30 +

producer

OH –

producer

Bronsted-Lowry proton donor proton acceptor

Lewis electron –pair acceptor

electron- pairdonor

Arrhenius Acids and BasesArrhenius acid – is a chemical compound

that increases the concentration of hydrogen ions H+, in an aqueous solution.

Arrhenius base – is a substance that increases the concentration of hydroxide ions OH-, in aqueous solution.

Acid-Bases Theories

Acid-Bases TheoriesBronsted-Lowry acid

-Is a molecule or an ion that is a proton donor (H+)

- Donates proton to water

HCl + H2O H3O+ + Cl-

HCl + NH3 NH4+ + Cl –

In both reaction HCl is a Bronsted-Lowry acid

Bronsted-Lowry Acid

Water could be a Bronsted-Lowry acid also.

H2O (l) + NH3 (g) NH4+

(aq)+ OH-(aq)

The water donates a hydrogen ion (proton) to the ammonia molecule.

Bronsted-Lowry Base

A molecule or ion that is a proton acceptor (H+)

H2O (l) + NH3 (g) NH4+

(aq)+ OH-(aq)

Ammonium (NH4) is a Bronsted-Lowry Base because it accepts the proton from the water.

Bronsted-Lowry Acid-Base Reaction

Protons are transferred from one reactant (the acid) to another (base).

H2O (l) + NH3 (g) NH4+

(aq)+ OH-(aq

Bronsted-Lowry Acid-Base Monoprotic Acids – an acid that can donate

only one proton (HCl, HNO3)

Polyprotic Acids – an acid that can donate more than one proton per molecule.

- Diprotic – can donate two protons per molecule (Sulfuric acid, H2SO4)

- Triprotic – can donate three protons per molecule(Phosphoric acid, H3PO4)

Acid –Base Reaction

The Bronsted-Lowry definition s of acids and bases provide the basis for studying proton (H+) transfer reaction.

Suppose a Bronsted-Lowry acid gives up a proton, the remaining ion or molecule can re-accept that proton and can act as a base- a conjugate base.

Acid –Base Reaction

Conjugate base- the ion or molecule that remains after a Bronsted-Lowry acid has given up a proton is the conjugate base of that acid.

Conjugate acid- the ion or molecule that is formed when a Bronsted-Lowry base gains a proton is the conjugate acid of that base.

Acid –Base Reaction

HF + H2O F- + H3O+

Acid BaseConjugate Base

Conjugate Acid

The species remaining after a Brønsted-Lowry acid gives up its proton is the conjugate base of that acid: Take off one H from the acid.

The species remaining after a Brønsted-Lowry base accepts its proton is the conjugate acid of that base: Take off one H. Add an H to the base.

Acid –Base ReactionHCO3

- (aq) + H2O (l) H2CO3 (aq) + OH-

(aq)

base acid conjugate acid

conjugate base

(proton acceptor)

HF (aq) + H2O (l) F- (aq) + H3O+

(aq)

acid baseconjugate base

conjugate acid

(proton donor)

Strength of conjugate acids and bases

The stronger an acid is, the weaker its conjugate base; the stronger a base is, the weaker its conjugate acid.

The weaker an acid is, the stronger its conjugate base; the weaker a base is, the stronger its conjugate acid.

Lewis Acids and Base

A Lewis acid is an atom, ion, or a molecule that accepts an electron pair to form a covalent bond.

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

Lewis Acids and Base

H+ (aq) + : NH3

(aq) [H-NH3

+ ] (aq)

H+ (aq) + : NH3

(aq) [NH4

+ ] (aq)

A bare proton (H+) is a Lewis acid in reactions in which it forms a covalent bond.

Lewis Acids and Base . .

BF3 (aq) + : F : - (aq) BF4 – (aq)

. .

An anion is a Lewis Base in a reaction in which it forms a covalent bond by donating an electron pair.

Indicators Chemical dyes that change color as pH

changes. Different indicators change colors at

different pH levels choose an indicator that will show a color

change at the pH that you are interested in. Indicators can be on a strip of paper

called pH or litmus paper Other indicators can be added to the

solution directly. Some indicators change color more than once

and can be added to solutions so that we can see what is happening over time.

Acid/Bases/Salts

Neutralization/Titrations

Neutralization

Chemical reaction between an acid and a base.

Products are a salt (ionic compound) and water.

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

Neutralization

HCl + NaOH

H2SO4 + KOH

HNO3 + Ca(OH)2

NaCl + H2O

K2SO4 + H2O

Ca(NO3)2 + H2O

-1 +1

-2 +1

-1+2

Hydrocholoric acid Sodium

hydroxide

Sodium chloride

Water

Sulfuric acid Potassium hydroxide

Potassium sulfate Water

Nitric acid Calcium hydroxide

Calcium nitrate

water

2 2

2 2

Titration

Titration Analytical method in

which a standard solution is used to determine the concentration of an unknown solution.

standard solution

unknown solution

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

Titration

• dramatic change in pH

Titration

moles H3O+ = moles OH-

MV n = MV nM: MolarityV: volumen: # of H+ ions in the acid

or OH- ions in the base

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