Hein's Chem Flashcards

Acids/Bases HL/SL

Acid

Acid

any compound that forms H+ ions in solution

Base (Alkali)

produces hydroxide Ions (OH+)

Neutralization reaction

the reaction of an acid and a base to form a neutral solution of water and a salt

What allows the Brønsted-Lowry

Brønsted-Lowry Acid Definition

proton (H+) donor

Proton representation in aqueous solution

• proton (H+)

• Hydronium (H30+) formed when a water molecule and proton coordinately bond

Brønsted-Lowry Base Definition

proton (H+) acceptor

Monoprotic Acid

An acid that can only donate one proton

• example: hydrochloric acid

Diprotic Acid

an acid that can donate two protons

• example: sulfuric acid

Triprotic Acid

an acid that can donate three protons

• example: phosphoric acid

Conjugate Base

a base that forms when an acid loses a proton

Conjugate Acid

an acid that forms when a base gains a proton

Conjugate acid-base pair

consists of two substances related to each other by the donating and accepting of a single hydrogen ion (proton)

Amphiprotic

A species that can either accept (act as a Brønsted-Lowry base) or donate a proton (act as a Brønsted-Lowry acid).

Acid Properties

• taste sour

• pH < 7.0

• litmus is red

• phenolphthalein is colorless

• methyl orange is red

Base Properties

• taste bitter

• ph > 7.0

• litmus is blue

• phenolphthalein is pink

• methyl orange is yellow

Standard Enthalphy Change of Neutralization

the energy change associated with the formation of 1 mol of water from the reaction between a strong acid and a strong base under standard conditions

Why are neutralization reactions always exothermic (ΔH<0)?

Energy released when H2O is formed

Enthalpy Change of Neutralization of strong vs weak acids

almost identical as neutralization drives the reaction to completion

Enthalpy change in a strong acid and base reaction

Acid and Base fully dissociate so only the enthalpy change of formation of water from hydrogen and hydroxide ions is needed (Exothermic)

Enthalpy change in a reaction involving a weak acid or base

Ionization of a weak acid or base is endothermic so neutralization of a strong base-weak acid will be slightly less exothermic than a strong base-strong acid

The weaker the acid the more endothermic the dissociation reaction becomes, lowering the enthalpy change of neutralization

Enthalpy Change Negative

exothermic

Enthalpy Change Positive

endothermic

Alkali

A soluble base

Common Examples of Bases

• metal hydroxides

• metal oxides

• ammonium hydroxide (Weak)

acid + metal

salt + hydrogen

acid + base

salt + water

acid + metal carbonate/metal hydrogencarbonate

salt + carbon dioxide + water

pH scale

a simple and effective way of representing the concentration of hydrogen ions [H+] in a solution

pH definition

measure of hydrogen ion concentration

Calculating pH from [H+(aq)]

-log[H+(aq)]

Find [H+] from pH

10^-pH

Acids on the pH scale

below 7

Bases on the pH scale

above 7

if [H+] increases what happens to the pH?

pH decreases

Change of 1 pH unit

10 fold change in [H+]

Calculating the pH of strong acids and strong bases

concentration of hydrogen ions [H+] is the same as the concentration of the acid

Ex: 0.1 mol dm^-3 of hydrochloric acid equates to [H+] = 0.1 mol dm^-3

Ionization of Water (Considering strong acids and weak bases)

ionization of water is endothermic

ion product constant (Kw) of water

1.0 x 10^-14

[H2O] is constant (Kw)

the dissociation constant or ionization constant of water

Pure water: 1.0 x 10^-14 so [H+]=[OH-]= 1.0 x 10^-7

Find pH from [OH-]

pH = 14 - (-log[OH-]) or

1. [H+] = (10^-14)/[OH-]

2. pH = -log[H+]

What determines the strength of an Acid or Base?

the degree to which it ionizes or dissociates in water

Strong Acid Definition

an effective proton donor that is assumed to completely dissociate in water and their equations are assumed to go to completion (→) with the concentration of each of the two ions produced being the same as the initial concentration

Strong Acid Examples

• HCl (hydrochloric acid)

• HNO3 (nitric acid)

• H2SO4 (sulfuric acid)

Weak Acid Definition

a poor proton donor that dissociates only partially in water where at equilibrium a majority of the weak acid molecules remain unreacted

Weak Acid Examples

• CH3COOH (Ethanoic acid)

• H2CO3 (Carbonic acid)

• H3PO4 (phosphoric acid)

Difference of Conjugate bases of Weak and Strong acids

Conjugate base of a weak acid has a higher affinity for a proton than the conjugate base of a strong acid

Strong acids have ___ conjugate bases

weak

weak acids have ____ conjugate bases

strong

Strong and Weak Acids and bases are distinct from

concentrated and dilute

Strong Base Definition

A proton acceptor that completely dissociates in water

Strong Base Examples

Group 1 metal Hydroxides

• LiOH (Lithium Hydroxide)

• NaOH (Sodium Hydroxide)

• KOH (Potassium Hydroxide)

• Ba(OH)2 (Barium Hydroxide

Can metal hydroxides act as a Brønsted-Lowry base

No, it doesn't have the capacity to accept a proton but in a solution the hydroxide ion acts as the Brønsted-Lowry base, accepting the proton

Weak Base Definition

A proton acceptor that does NOT completely dissociate in water

Weak Base Examples

• NH3 (Ammonia)

• C2H5NH2 (Ethylamine) and other amines

Calculating pH of Monoprotic Acids

[H+] = [Acid]

Calculating pH of Diprotic acids

[H+] = 2[Acid]

Experimental determination of the strength of acids and bases

Dissociated acids and bases in water create ions, which are charge carrier, the stronger the acid or base is the more ions are produced increasing the conductivity and increases the reading on the ammeter

Acid deposition Definition

the process by which acid-forming pollutants are deposited on the earth's surface

Acid Deposition effects

• deforestation

• leaching of minerals from soils leading to elevated acid levels in lakes and rives

• uptake of toxic minerals by plants

• reduction of pH of lake and river systems

• uptake of toxic metals by shellfish and other marine animals which can damage the fishing industry

• Corrosion of marble, limestone, and metal buildings, bridges and vehicles

pH of regular rain

5.6 due to the presence of dissolved Carbon Dioxide which forms weak carbonic acid

pH of acid rain

less than 5.6

What causes acid rain?

• Natural: volcanic eruptions and the decomposition of vegetation

• Unnatural: burning of fossil fuels

Main types of acid rain

Nitric acid (HNO3) and sulfuric acid (H2SO4)

Pre-combustion methods to reduce sulfur emissions

• 80-90% of all inorganic sulfur removed

• mineral benefaction involves crushing coal, followed by floatation

Post-combustion method to remove sulfur dioxide, nitrogen oxide, heavy metals and dioxins

Calcium oxide and lime will react with sulfur dioxide and remove it from flue gases

Lewis Acid Definition

an electron pair acceptor

Lewis Base Definition

an electron pair donor

Are all Brønsted-Lowry acids and bases Lewis acids and bases?

Yes, since Brønsted-Lowry acids donate protons (acting as a Lewis acid by accepting an electron pair) and Brønsted-Lowry bases accept protons (acting as a Lewis base by donating an electron pair)

Are all Lewis acids and bases Brønsted-Lowry acids and bases?

No, not all Lewis acids/bases are Brønsted-Lowry acids/bases as they don't donate or accept an H+ ion

Forming Compound Bonds

The Lewis base donates a pair of electrons which to a hybridized 2sp2 orbital (H+ doesn't have to be involved)

Transition Metals Lewis Acids/Base

Transition metals act as Lewis Acid with the Ligand acting as the Lewis Base

Nucleophiles

• electron rich with at least one lone pair of electrons

• Lewis Base

Electrophile

• an electron-deficient species that can accept a lone pair from a nucleophile

• Lewis Acid

Ka

acid dissociation constant

pKa

= -logKa

Acid Strength relationship with Ka

the higher Ka, the stronger the acid (lower pH)

Acid strength relationship with pKa

the higher pKa, the weaker the acid (higher pH)

Where can Ka and Kb values be found in the data booklet

section 21

pKb

= -logKb

Base Strength relationship with Kb

higher the Kb the stronger the base (higher pH)

Base strength relationship with pKb

higher pKa the weaker the base (lower pH)

Calculate the pH of a weak acid

Calculate the pH of a weak base

Ka calculation for weak acid and strong base reaction (Excess HA)

1. Find the moles of the HA and OH-

2. Find number of HA moles left and A- moles formed

3. Calculate [HA] leftover and [A-] formed

4. Use Ka to find [H+] using equation to the left

5. pH is the -log[H+]

Note with different molar ratios it has to be multiplied or divided

Ka calculation for weak acid and strong base reaction (Excess OH-)

1. Find the moles of the AH and OH-

2. Find number of OH- moles left and A- moles formed

3. Calculate [OH-] left over

4. Calculate pOH is the -log[OH-]

5. pH = 14 - pOH

Note with different molar ratios it has to be multiplied or divided

Ka calculation for weak acid and strong base reaction (No excess since mol HA = OH-)

pH = pKa of weak acid

pH and pOH equations

pH + pOH = 14

Relationship between Ka, Kw, and Kb

1. Ka x Kb = Kw

2. Ka = Kw/Kb

3. Kb = Kw/Ka

The stronger the acid:

• the larger the Ka

• the weaker the conjugate base

• the smaller the Kb of the conjugate base

The stronger the base:

• the larger the Kb

• the weaker the conjugate acid

• the smaller the Ka of the conjugate acid

Temperature affect on ionic product constant (Kw)

change in the temperature of the system will result in a change in the position of equilibrium resulting in the forward reaction being favored, increasing the concentration of hydrogen and hydroxide ions.

• represents an increase in the magnitude of Kw and decrease in pH

• distinction between the neutrality and pH of the solution as the pH decreases with an increase in [H+] but [OH-] increases equally so the solution remains neutral

Temperature affect on ionic product constant (Kw) data booklet

section 23

Salt Hydrolysis

salt completely separate into their ions when dissolved in an aqueous solution

Salt hydrolysis of strong acid and strong base

neutral with pH equal to or close to 7

Salt hydrolysis of weak acid and strong base

are basic with a pH greater than 7

Salt hydrolysis of strong acid and weak base

are acidic with a pH less than 7

Salt hydrolysis of weak acid and weak base

tend to be neutral with the pH dependent on the relative values of pKa and pKb

Buffer Solution

a solution consisting of a weak acid and its conjugate base (stored in salt) or a weak base and its conjugate acid (stored in salt), which can resists a change in pH when a small amounts of strong base or strong acid is added

How buffer solutions work (addition of strong acid)

Strong acids fully dissociate in water which increases the [H+], which are then mopped up by the conjugate base keeping the pH constant