Reactions in Chemistry

Combination reactions can be exothermic or endothermic, but are more often exothermic, and negative entropy. Decomposition reactions are therefore mostly endothermic, and positive entropy.

Furthermore, for a change in entropy to be positive, are reactions with more moles of gas in the product side (decomposition reactions).

Heating a reaction shifts it to the endothermic side, while cooling it shifts to the exothermic side. Increasing the pressure on an equilibrium mixture (decreasing the volume) will shift the reaction to the side containing the fewest number of gas molecules, and increasing the volume (decreasing the pressure) will shift it to the side containing the most moles of gas.

Fusion reactions are exothermic as long as the starting material are bigger than Fe-56, with products less than Fe-56. They are endothermic if products are bigger than Fe-56. For fission reactions, are vice versa.

Chain reactions can be endothermic, but such a reaction would need to be in contact with a heat source in order to carry out its chain process.

The fastest reactions in chemistry, by far, are acid-base reactions in water, as the molecules don’t have to touch each other. The 2nd fastest reactions are generally the electron-transfer reactions. However, depending on the context, photochemical reactions are faster, such as photoisomerization.

Combination reactions favor the higher oxidation state, replacement reactions favor the lower oxidation state.

For an explosion to occur, the decomposition or combustion of the explosive must be exothermic (so that a detonation wave spreads), and the reaction products must be gaseous. The most dangerous endothermic reactions are when gases are produced (which would be from a hot gas to a cold gas). And the most dangerous reactions where no gases are produced, are exothermic reactions going from a cold solid to a hot liquid (or something almost a gas, due to the pressure change).

For all cases, 1 should never mix strong oxidants with oxidizable material whenever there is the possibility of a reaction to produce gas.

The act of oxygen combining with iron to form rust is hugely exothermic, about -826 kJ/mol, which means that about 100 grams of iron, in the course of rusting, would release enough energy to raise 1 liter of water's temperature about 200 degrees C. However, rust happens so slowly (over the course of years to decades) that that energy comes out in such a small trickle, we don't notice it.

Gibbs free energy.

It is impossible to predict S and H a priori without any computer quantum calculations. This lack of predictions, will explain whether a combination reaction will be exothermic or endothermic, and whether something that dissolves in water will be exothermic or endothermic.

If solubility decreases with higher temperature, then change in S must be negative. Then the salt will dissolve if change in H is negative enough to give an overall negative result.

These all dissolve better in cold water (exothermic):

Ammonium bromoplatinate, sulfate octahydrate, hydroxystannate, sodium selenite, sodium dihydrogen pyrophosphate hexahydrate, sulfate octahydrate, ytterbium sulfate, and virtually any substance that's a gas at ordinary temperatures, including nitrogen, oxygen, hydrogen, helium, carbon dioxide, and ammonia.

Gases.

All gas reactions are either exothermic and spontaneous, and not endothermic and spontaneous, or, if you flip the reaction around, are endothermic and non-spontaneous, and not exothermic and non-spontaneous.

Gas reactions are further either combination or decomposition reactions, all that are exothermic in their spontaneous-side. (So if the gas reactions are spontaneous, then they are supposed to be exothermic.). You could, theoretically, have a endothermic spontaneous reaction, in their decomposition form, which would require the gases to absorb heat from their surroundings. But these would not be dependent on the identity of the gas, only the controlled environments.

Advanced - finding the temperature of a reaction, at equilibrium.

Reactions tend to happen after an activation energy, Ea. The Ea is the same whether at equilibrium. The equilibrium constant, K, symbolizes equilibrium when it is 1. Make sure ΔG and R are in the same units, so if using kJ for ΔG, move the decimal 3 places to the right.

So take 2Na + Cl2 -> 2NaCl. In real life, we use a flame on the sodium metal when it's in a tube of chlorine gas. What is the minimum temperature needed for the reaction to react?

The answer is any temperature, for some reaction to occur, just that even at subzero temperatures, the speed is microscopic. But for the reaction at equilibrium (K = 1), T comes to be above their melting and boiling point. (There is a formula, for 1 = e(ΔG/RT), but, the values for ΔG and R are defined for 298 K, so you can't use that formula).

When K = 1, ΔG0 = 0, and T = ΔH/ΔS. (Because ΔG0 = -RT ln(K), and ln(1) = 0.).

Do not use the product of formation for ΔH and ΔS, use the products - reactants.

So for Na + 2Cl2 -> 2NaCl, T = 4547 K.

ΔH = 2* -410 - (2*0 +0) = -820.
ΔS = 2*.072 - (2*.051 + .2223) = -.1803.

T = -820/-.1803 = 4547 K.

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There are some reactions that don't need a flame to activate it. Such as.

-Forming of rust. 4Fe + 3O2 -> 2Fe2O3.
-Formation of quicklime. CaO + H2O -> Ca(OH)2.
-Decomposition of hydrogen peroxide. 2H2O2 -> 2H2O + O2.

In organic chemistry.

All Lewis acids are electrophiles, all Lewis bases are nucleophiles.

The 6 main types:

SN2: losing a halogen and doing a replacement. C temporarily has 5 bonds, no carbocation or carbanion. The nucleophile bonds to the electrophilic carbon as the substrate leaves. SN2 reactions require nucleophiles to be strong.

SN1: rearrangement (losing halogens and forming carbocations). Something falls off 1st, then something else bonds on. A substitution / replacement, involving weak nucleophiles that are not strong enough to do SN2.

E1: carbocation formed. A base abstracts a proton from the C adjacent to the carbocation. E1 reactions almost always occur together with SN1.

E2: requires a strong base present (like SN2, so no carbocations formed) but SN2 mechanism is blocked due to being hindered. So a strong base abstracts a proton on a C adjacent to the C with the leaving group, forming a double bond on the adjacent carbon. This makes E2 a concerted reaction as bonds break and form at the same time. No rearrangements.

SN2 and SN1 both stand for nucleophilic substitution, and we do have addition reactions.

Nucleophilic addition reaction: the nucleophilic carbon of the reagent attacks the electrophilic carbon of the carbonyl group.

Electrophilic addition reaction: involve the addition of electrophiles to a double or triple bond. Such as the addition of hydrogen halides to alkenes. Or adding water (hydration) to alkenes.

Electrophilic substitution reactions involve the substitution of a hydrogen atom in an organic compound with an electrophile. Common examples include electrophilic aromatic substitution reactions.

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Examples in nature for the above reactions:

DNA replication (SN2). During DNA replication, DNA polymerases catalyze the incorporation of nucleotides into a growing DNA strand. This process involves SN2 reactions where a nucleophile (incoming nucleotide) attacks the electrophilic phosphate group on the DNA backbone.

Protein folding uses both SN2 and SN1, while photosynthesis and fermentation uses both E1 and E2.

Glutathione (a tripeptide containing a thiol group), serves as a mild reducing agent to detoxify peroxides and maintain the cysteine residues of hemoglobin in the reduced state. It can also detoxify alkylating agents, where the thiol of glutathione reacts with methyl iodide by an SN2 reaction, making the methyl iodide harmless and preventing its reaction with other molecules in the body.

Some notes about aromatic rings.

Common aromatic systems have 2, 6, or 10 pi electrons, whereas anti-aromatic systems have 4, 8, or 12 pi electrons. Benzene (and pyridine and pyrrole) have 6.

Deactivating groups (electron-withdrawing) are for nucleophilic substitution and activating groups (electron-donating) are for electrophilic aromatic substitution.

Electron-withdrawing substituents (such as nitro groups and carboxyl groups) activate the ring toward nucleophilic aromatic substitution. These deactivating groups (such as carbonyls (aldehydes, ketones) and halogens) withdraws electron density by both the inductive effect (through the sigma bonds) and the resonance effect which involves pi systems.

Likewise, things that increase electron density of the ring (like hydroxyl groups, amine groups, and methyl groups), boosts electrophilic substitution.

An electron-withdrawing substituent deactivates primarily the ortho and para positions (such as nitro groups, for nucleophilic substitution), and an electron-donating substituent activates primarily the ortho and para positions (such as phenols, for electrophilic aromatic substitution). However, by deactivating the ortho and para position, we mean it primarily activates the meta position.

So an application of this, is when methylbenzene (toluene) reacts with strong acids (activating), they react a lot quicker than benzene does, and they react on the ortho and para positions. But when nitrobenzene reacts with strong acids (deactivating), it is much less reactive than benzene is, and the products that do form, form primarily in the meta position. (With strong bases like NaOH or Grignards, they are not strong enough to pull hydrogens from benzene or methylbenzene, but can from hydroxybenzene (phenol)).

Do note that in nitro groups, no matter how we position the electrons in a Lewis dot diagram, the nitrogen always has a formal positive charge, so it inductively withdaws electron density from the aromatic ring.

Electrophilic aromatic substitutions are far more important than nucleophilic aromatic substitutions.

In polymer chemistry.

2 types of polymer reactions: step-growth (which includes condensation reactions), and chain-growth (which includes free-radical and ionic reactions). Step-growth has a slow rate of increasing the molecular weight, whereas chain-growth has a fast rate of increasing the molecular weight. In step-growth reactions, water or carbon dioxide are typically a byproduct, and not in chain-growth. In chain growth, free-radical reactions are generally faster than ionic reactions, due to less stabilization.

Step-growth reactions tend to be endothermic, whereas chain-growth reactions tend to be exothermic after the initiator (the initiator is endothermic).

In toxicology.

Chemicals are dangerous depending on the type of reactions and route of exposures. Examples of types of reactions as listed:

“Damaging” reactions.

Includes acids and bases (dangerous to touch). Strong bases dissolve tissues. Sulfuric acid is both a strong acid, a strong oxidizer, and a strong dehydrating agent (absorbs water), whereas HNO3 is both a strong acid and a strong oxidizer, and HCl is just a strong acid and KMnO4 is just a strong oxidizer. The most dangerous commercial chemical is HF, a weak acid, but not because of the acid, but because of the fluorine.

“Replacement” reactions.

Compounds that are safe to touch, but dangerous to breathe or digest. Examples are carbon monoxide and cyanide gas (CO and HCN), as well as solid KCN and LiCN.

For example, in cyanide gases, cyanide bonds to Fe(III) in ferricytochrome oxidase enzyme, preventing reduction to Fe(II) in the oxidative phosphorylation process by which the body utilizes O2 (which prevents utilization of O2 in cells, so metabolic processes cease). And CO binds to hemoglobin to convert oxyhemoglobin to carboxyhemoglobin, at 210 times more than oxygen (so the pain is a lack of oxygen, and a lack of oxygen ceases production of ATP).

Other examples of chemicals besides cyanide, are H2S, azides, formates, NO-radicals, and PH3.
Other examples of chemicals besides carbon monoxide, are chemicals that form methemoglobin, such as nitroglycerin, sulfonamides, chlorobenzene, and arsine.

“Prevention” reactions.

Examples are organophosphates, where they inhibit (or prevent) acetylcholinesterase (an enzyme essential for nerve function), which accumulates excess acetylcholine at cholinergic synapses, with overstimulation of muscarinic and nicotinic cholinergic receptors. Excess acetylcholine causes paralysis of muscles needed for breathing and heartbeat. These are through absorption of the skin.

Another example is atropine, an alkaloid which can counter the effects of pesticides and nerve gas by blocking the receptors they over-activate.

Case study: the Haber process.

Haber-Bosch uses 8.6 * 1018 J of world's energy in 2021 (about 80% efficient). About .18 gigatonness (180 million tonnes) of ammonia produced, or about 3.5 * 1018 J. The flip side is about .42 gigatons of carbon dioxide are produced. Ammonia is about the same price as a $2 gallon of gasoline. Nearly 50% of the nitrogen found in human tissus originated from the Haber-Bosch process, which could explain why this is the donator of the population explosion, enabling the world population to increase from 1.6 billion in 1900 to 7.7 billion by Nov. 2018.

2 ways to get energy out of ammonia gas:

Combustion: ammonia can be combusted in the presence of oxygen to release heat energy.

4NH3(g) + 3O2(g) -> 2N2(g) + 6H2O(g) ΔH = -1269.6 kJ/mol.

This reaction can be utilized in combustion engines or power plants to generate heat, which can then be converted into electricity or used for other purposes.

Another method is to electrolyze ammonia. Ammonia fuel cells: ammonia is oxidized at the anode and reduced at the cathode, with the overall reaction producing electricity and nitrogen gas as byproducts. The specific reactions depend on the type of fuel cell used. In a direct ammonia fuel cell (DAFC), the following reactions occur:

Anode (oxidation): 6NH3 + 6H2O -> 6NH4+ + 6e-.
Cathode (reduction): 6NH4+ + 6e- + 6O2 -> 6 NO3- + 12H2O.
Overall reaction: 6NH3 + 6O2 -> 6NO3- + 12H2O + 6e-.

The electrons generated from the oxidation reaction flow through an external circuit, producing electricity. Meanwhile, the oxidized ammonia (NO3-) combines with water to form nitrate ions.