WORLD OF MOLECULES

(1)

Development of Complex Curricula for Molecular Bionics and Infobionics Programs within a consortial* framework**

Consortium leader

PÁZMÁNY PÉTER CATHOLIC UNIVERSITY

Consortium members

SEMMELWEIS UNIVERSITY, DIALOG CAMPUS PUBLISHER

The Project has been realised with the support of the European Union and has been co-financed by the European Social Fund ***

**Molekuláris bionika és Infobionika Szakok tananyagának komplex fejlesztése konzorciumi keretben

***A projekt az Európai Unió támogatásával, az Európai Szociális Alap társfinanszírozásával valósul meg.

PÁZMÁNY PÉTER CATHOLIC UNIVERSITY SEMMELWEIS

UNIVERSITY

(2)

WORLD OF MOLECULES

CHEMICAL EQUILIBRIA, ACID-BASE THEORIES

(Molekulák világa)

(Kémiai egyensúlyok, sav-bázis elméletek)

KRISTÓF IVÁN

semmelweis-egyetem.hu

(3)

World of Molecules: Chemical equilibria, acid-base theories

Previously – Chemical compounds, stoichiometry

1. Compounds

2. Chemical composition

3. Ambiguity of the chemical formula 4. Stoichiometry

5. Main groups of chemical compounds 6. Grouping of inorganic compounds 7. Salts

8. Properties of water

(4)

• organic, which has – C – C – bonds

• inorganic, which lacks – C – C – bonds

• acids

• the reaction result of water and oxydes of non-metallic compounds

• bases

• the reaction result of water and oxydes of metallic compounds

• salts

• produced from the reaction of acids and bases

• metal complexes

• a central, metallic atom or ion is bonded to the surrounding

molecules/ligands, which usually donate dative bond to the metal

Previously - Grouping of chemical compounds

(5)

Previously - Different phases of crystalline water

(6)

Previously - Network of Hydrogen bonds with 500 molecules

(7)

1. Chemical equilibria

• equilibria in gases

• acid-base equilibria

2. Acid-base theories

• Arrhenius theory

• Brønsted-Lowry theory

• Lewis theory

• Pearson theory (HSAB)

3. Superacids and superbases

Chemical equilibria

• regarding reversible reactions in gaseous or liquid phase

• e.g. H

₂

+ I

₂

2 HI

• the equilibrium is dynamic

• reaching the equilibrium means the rate of the two opposing reactions become equal

• the ratio between reactants and products remain fixed (concentrations are constants)

• description of equilibrium (for simple reactions): law of mass action (by Guldberg and Waage, 1867)

(9)

Law of mass action (1867)

• can be derived from the equality of the forward and backward reaction rates

• the product of the equilibrium concentrations of the reaction products raised to a power according to their stoichiometic constant divided by the product of the equilibrium

concentrations of the reactants raised to their respective stoichiometic power is constant (K) in equilibrium.

• (under the same conditions: pressure, Volume, Temperature )

• e.g. for the above reaction (H₂ + I₂ ֖ 2 HI)

[ ] [ ] [ ]

^constant

2

⋅ =

= H I

K HI

(10)

• H

₂

+ I

₂

2 HI • 2 HI H

₂

+ I

₂

Chemical equilibria

[ ] [ ] [ ]

₂ ₂

2

I H

K HI

= ⋅

G

[ ] [ ]

[ ]

^HI ^I ^K

KH H 1G

2 2

2 ⋅ =

=

•the relationship of the equilibrium constant to the direction of the reaction is evident

•to be more precise instead of the concentration, the more accurate activity value of each molecule should be substituted into the law of mass action

•the activity can be approximated with the concentration in ideal gases or solutions

(11)

Le Châtelier - Braun's principle

• If a chemical system at equilibrium experiences a change in its status, then the equilibrium shifts to

counteract the imposed change and a new equilibrium is established.

• status:

• concentration,

• temperature,

• volume,

• pressure,

• ...

(12)

Le Châtelier - Braun's principle - example

• PCl₅ ֖ PCl₃+ Cl₂

• direction of concentration changes

• increasing the concentration of PCl₃

• the equilibrium will favor the porduction of reactants

• since the equilibrium constant remains constant, increasing the PCl₃ concentration, decreases the concentration of Cl₂, and increases the concentration of PCl₅ reactant.

[ ] [ ] [

³ 5

]

²

PCl Cl

K PCl

= ⋅

G

(13)

Le Châtelier - Braun's principle - example

• decreasing the concentration of PCl₃

• the equilibrium will favor the porduction of the products

• decreasing the PCl₃ concentration, increases the concentration of Cl₂, and decreases the concentration of PCl₅ reactant.

[ ] [ ] [

³ 5

]

²

PCl Cl

K PCl

= ⋅

G

(14)

Le Châtelier - Braun's principle - example

• increasing the concentration of PCl₅

• the equilibrium will favor the porduction of products

• basically, we introduce fresh amounts of the reactant, thus this will increase the concentration of both products Cl₂ and PCl₃.

[ ] [ ] [

³ 5

]

²

PCl Cl

K PCl

= ⋅

G

(15)

Le Châtelier - Braun's principle - example

• decreasing the concentration of PCl₅

• the equilibrium will favor the porduction of the reactant

• basically, we remove reactant, therefore the products will favor the reverse direction.

[ ] [ ] [

³ 5

]

²

PCl Cl

K PCl

= ⋅

G

(16)

• Haber-Bosch process

• N

₂

+ 3H

₂

2 NH

₃

• hard to manage industrially

• requires

• 150-250 bar pressure

• 300-550 °C

Chemical equilibria

[ ] [ ] [ ]

₂ ₂ ³

2 3

H N

K NH

= ⋅ G

(17)

Le Châtelier - Braun's principle - example

• 1N₂ + 3H₂ ֖ 2 NH₃

• K_p is the equilibrium constant with partial pressures

• p’ is the new partial pressure of the compound

• Q is the non equilibrium reaction quotient

• increasing the pressure of the system

• the compression will enforce fewer molecules to form

• therefore the reaction will favor the products to form

[ ] [ ] [ ]

2 3

3

2 2

K NH

N H

= ⋅

G

K Q

p K p

Q p p

p

H N

NH NH

NH

G

→

⋅ =

=

= 4

1 8

2 4

,

2 ₂

2 '

2 2

3 3

3

2 2

3

H N

NH

p p p

K p

= ⋅ G

(18)

Le Châtelier - Braun's principle - example

• increasing the pressure of reversible gas reactions in equilibrium will cause the reaction to remove

molecules from the system

• in case of reactions where the total number of molecules change the equilibrium will shift

• e.g. 1N₂ + 3H₂ ֖ 2 NH₃ where 1+3=4 molecules react to result in 2 molecules

• in case of reaction where there is no molecule number change such an effect cannot be observed

• e.g. 1H₂ + 1I₂ ֖ 2HI, where 2 molecules react to produce 2

(19)

Le Châtelier - Braun's principle - example

• increasing the temperature of reversible gas reactions in equilibrium will favor

• endothermic reactions, where heat is required for the reaction to take place

• decreasing the temperature of an equilibrium

• will favor exothermic reaction, where heat removal is required for the reaction to take place

e.g. a dimerization, 2 NO₂ ֖ N₂O₄

2 NO₂ → N₂O₄ is an exothermic process 2 NO₂ ← N₂O₄ is an endothermic process NO₂ is brown, N₂O₄is colorless

(20)

Le Châtelier - Braun's principle - example

the equilibrium reaction of

2 NO₂ ֖ N₂O₄ left: hot, more NO₂ right: cold, more N₂O₄

(21)

• in the case of homogenous equilibrium both the

reactants and the products are in the same phase (gas, liquid)

• in the case of inhomogenous equilibrium some of the reactants or products are in a different phase

• e.g. C(s)+CO₂(g) ֖ 2CO(g)

• CaCO₃(s) ֖ CaO(s) + CO₂(g)

• thus the equilibrium constants are modified accordingly

Homogenous and heterogenous equilibria

CO2

KG = p

2

2 CO CO

K p

= p G

(22)

• in case of a general acid, the dissociation reaction is

• HA ֖ H⁺ + A^-

• the equilibrium constant can be written

• it is called the acid dissociation constant (K_a)

• for water

• 2H₂O ֖ H₃O⁺ + OH^-

Acid dissociation equilibrium

[ ] [ ]

[ ]

^a

^{( )}

^a

a pK K

HA A

K = H ⁺ ⋅ ⁻ , = −log

[ ] [ ]

[ ]

2

[

3

] [ ]

¹⁴ ²

2

3 ⁺ ⁻ 10⁻ M

−

+ ⋅ → = ⋅ =

= K H O OH

O H

OH O

K_a H _w

=14 pKw

(23)

name pK_a

benzoic acid 4.204

formic acid 3.751

hydrofluoric acid 3.17

hydrocyanic acid 9.21

hydrogen peroxide 11.7

nitric acid -1.64

acetic acid 4.76

nitrous acid 3.37

Acid dissociation constant of some compounds

(24)

• in case of a general base, the dissociation reaction is

• BOH ֖ B⁺ + OH^-

• the equilibrium constant can be written

• it is called the base dissociation constant (K_b)

• connection between K_b and K_a:

• B + H₂O ֖ BH⁺ + OH^-

Base dissociation equilibrium

[ ] [ ]

[ ]

^b

^{( )}

^b

b pK K

BOH OH

K = B⁺ ⋅ ⁻ , = −log

[ ] [ ]

[ ] [ ] [ ]

a w b

w

b K

K K OH

O H K

B and OH

K BH ⋅ = ⋅ ⇒ =

= ⁺ ⁻ ₃ ⁺ ⁻

=14

=

+ _b _w

a pK pK

pK

(25)

name pK_b

ammonia 4.775

methylamine 3.355

thiethylamine 3.28

aniline 4.6

pyrrolidine 2.6

pyridine 8.8

Base dissociation constant of some compounds

(26)

• organic, which has – C – C – bonds

• inorganic, which lacks – C – C – bonds

• acids

• the reaction result of water and oxydes of non-metallic compounds

• bases

• the reaction result of water and oxydes of metallic compounds

• salts

• produced from the reaction of acids and bases

• metal complexes

• a central, metallic atom or ion is bonded to the surrounding

molecules/ligands, which usually donate dative bond to the metal

Grouping of chemical compounds

(27)

• acid: dissociates to produce proton

• HA ֖ H⁺ + A^-

• increases the number of protons in water

• base: dissociates to produce hydroxide ions

• BOH ֖ B⁺ + OH^-

• increases the number of hydroxide ions in water

• aqueous acids and bases (only aqeous is considered!)

• neutralize according: H⁺_(aq) + OH^-_(aq) ֖ H₂O

• acid + base → salt + water

• e.g. 2 NaOH + H SO → 2 H O + Na SO

Arrhenius – Ostwald acid base theory

(28)

• Brønsted acid (HA): proton donor

• HA ֖ H⁺ + A^-

• Brønsted base (B): proton acceptor

• B + H⁺֖ BH⁺

HA + B A

^-

+ BH

⁺

• the resulting ions can also be classified

• A^- : conjugate base,

• in the reverse direction this can accept a proton

• BH⁺ : conjugate acid,

• in the reverrse direction this can donate a proton

Brønsted – Lowry acid base theory

(29)

• according to this theory

• the autoprotolysis/self dissociation of water

• H

₂

O + H

₂

O H

₃

O

⁺

+ OH

^-

• can be classified both

• H₂O: Brønsted acid

• H₂O: Brønsted base

• furthermore

• H₃O⁺ : conjugate acid

• OH^- : conjugate base

Brønsted – Lowry acid base theory

property amphoteric

⎭⎬

⎫

(30)

• examples (water acts as the counterpart)

• Brønsted acids – proton donors

• HCl + H₂O ֖ Cl^- + H₃O⁺

• CH₃COOH + H₂O ֖ CH₃COO^- + H₃O⁺

• Brønsted bases – proton acceptors

• H₂O + NH₃ ֖ OH^- + NH₄⁺

• these acids and bases can always be defined by the reaction

• acid + base conjugate base + conjugate acid

Brønsted – Lowry acid base theory

(31)

• Lewis acid (A)

• electron pair acceptor

• Lewis base (:B)

• electron pair donor

• through an acid – base reaction a dative covalent bond is created

A + :B A-B

• the product is called a Lewis adduct

• e.g.: (CH

₃

)

₃

B + :NH

₃

(CH

₃

)

₃

B-NH

₃

Lewis acid base theory

(32)

• examples ( Lewis acid (A), Lewis base (:B))

• H

⁺

+ :NH

₃

→ NH

₄⁺

• BF

₃

+ F

⁻

→ BF

₄⁻

• PCl

₅

+ Cl

⁻

→ PCl

₆⁻

• SF

₄

+ F

⁻

→ SF

₅⁻

• CO + BF

₃

→ COBF

₃

• H

₂

O

• molecules with non-bonding electron pairs

• two non-bonding electrons on HOMO orbitals

Lewis acid base theory

(33)

• Oxygen theory of acids and bases,

used in the geochemistry and electrochemistry of molten salts

• Acid: oxide ion acceptor (O

^2-

) (some form)

• Base: oxide ion donor

• e.g.:

• MgO + CO₂ → MgCO₃

• CaO + SiO₂ → CaSiO₃

• NO₃⁻ + S₂O₇²⁻ → NO₂⁺ + 2 SO₄²⁻

Lux-Flood acid base theory

(34)

• Lewis base = Brønsted base

• the electron pair can be donated to the proton

• conjugate base = Lewis base

• an electron pair remains after donating a proton

• BUT: Lewis base cannot be protonated easily

• acid-base strength do no compare in the two theories

• Brønsted –Lowry theory is useful in non-aqueous solvents (proton is still required)

• Lewis theory is useful in complex/coordination chemistry

Brønsted – Lewis compatibility

(35)

• Hard and Soft Acids and Bases

• hard acids and bases

• small atomic and ionic radius

• high oxidation state

• low polarizability

• high electronegativity

• low energy HOMOs (in case of bases)

• high energy LUMOs (in case of acids)

• hard acids: H⁺, Na⁺, K⁺, ...

• hard bases: OH^-, F^-, Cl^-, ...

HSAB theory (Pearson, 1963)

(36)

• soft acids and bases

• large atomic and ionic radius

• low or zero oxidation state

• high polarizability

• low electronegativity

• soft bases have higher energy HOMOs than hard bases

• soft acids have lower energy LUMOs than hard acids

• soft acids: Hg²⁺, Pt²⁺, Ag⁺, ...

• soft bases: H^-, SCN^-, I^-, ...

HSAB theory (Pearson, 1963)

(37)

HSAB theory - examples

(38)

HSAB theory – hardness values of acids and bases (in eV)

(39)

• this theory is the extension of the Lewis acid - base theory

• the theory predicts that the interaction between hard acids bases (ionogenic), and soft acids and bases

(covalent character) are more stable

• useful for reaction mechanism descriptions

• softness of a base (B) is described by determining the equilibrium constant for the following

• BH + CH₃Hg⁺ ֖ H⁺ + CH₃HgB

• where CH Hg⁺is a soft acid and H⁺is a strong acid

HSAB theory (Pearson, 1963)

(40)

• special acids that have acidity greater than pure sulfuric acid (pK

_a

~ -3)

• used to create and maintain organic cations for certain reactions

• usually prepared by the combination of a strong Lewis acid and a strong Brønsted acid

• e.g. SO₃ + HF → FSO₃H (fluorsulfuric acid)

• trifluoromethanesulfonic acid, CF

₃

SO

₃

H, pKa~ -15

• Fluoroantimonic acid, HSbF

₆

, pKa~ -25

Superacids

(41)

György Oláh

• used magic acid (which has a pK_a~ -20) to protonate

hydrocarbons,

• magic acid: fluorsulfonic acid and antimony pentafluoride (HSFO₃·SbF₅)

• protonated hydrocarbons, i.e.

organic cations are useful in many synthetic reactions

• CH₄ + H⁺ → CH₅⁺

Superacids – György Oláh, Nobel prize in chemistry, 1994

(42)

• a compound that has a high affinity to protons in non- aqueous solutions

• destroyed by water, carbon dioxide, and oxygen

• used in organic synthesis for deprotonation

• e.g. lithium diisopropylamide, LiN(C₃H₇)₂, pK_b~ -16 (in tetrahydrofuran),

• e.g. proton sponge: 1,8-Bis(dimethylamino)-naphthalene, C₁₄H₁₈N₂, pK_b~ 1.9,

• classes: organic, organometallic, inorganic

Superbases

(43)

WORLD OF MOLECULES

WORLD OF MOLECULES

Previously – Chemical compounds, stoichiometry

1. Compounds

2. Chemical composition

3. Ambiguity of the chemical formula 4. Stoichiometry

5. Main groups of chemical compounds 6. Grouping of inorganic compounds 7. Salts

8. Properties of water

• organic, which has – C – C – bonds

• inorganic, which lacks – C – C – bonds

Previously - Grouping of chemical compounds

Previously - Different phases of crystalline water

1. Chemical equilibria

2. Acid-base theories

3. Superacids and superbases

Table of Contents

Chemical equilibria

• regarding reversible reactions in gaseous or liquid phase

• e.g. H

+ I

2 HI

• the equilibrium is dynamic

Law of mass action (1867)

[ ] [ ] [ ]

• H

+ I

2 HI • 2 HI H

+ I

Chemical equilibria

[ ] [ ] [ ]

[ ] [ ]

[ ]

Le Châtelier - Braun's principle

• If a chemical system at equilibrium experiences a change in its status, then the equilibrium shifts to

counteract the imposed change and a new equilibrium is established.

• status:

Le Châtelier - Braun's principle - example

[ ] [ ] [

]

PCl Cl

K PCl

= ⋅

G

Le Châtelier - Braun's principle - example

[ ] [ ] [

]

PCl Cl

K PCl

= ⋅

G

Le Châtelier - Braun's principle - example

[ ] [ ] [

]

PCl Cl

K PCl

= ⋅

G

Le Châtelier - Braun's principle - example

[ ] [ ] [

]

PCl Cl

K PCl

= ⋅

G

• Haber-Bosch process

• N

+ 3H

2 NH

• hard to manage industrially

• requires

Chemical equilibria

[ ] [ ] [ ]

Le Châtelier - Braun's principle - example

[ ] [ ] [ ]

Le Châtelier - Braun's principle - example

• increasing the pressure of reversible gas reactions in equilibrium will cause the reaction to remove

molecules from the system

Le Châtelier - Braun's principle - example

Le Châtelier - Braun's principle - example

• in the case of homogenous equilibrium both the

^{( )}

^{( )}