Space conservation versus Energy conservation

Introduction: 

Standard Science View

1. Law of mass conservation

In chemical reactions, the total mass of atoms (protons + neutrons + electrons) remains constant.

Bonds break and reform, but atoms don’t vanish.

Example: 2Mg + O₂ → 2MgO → Mg atoms + O atoms remain, only rearranged.

2. Law of energy conservation

Total energy before and after is the same.

But chemical potential energy stored in bonds converts into heat and light when more stable bonds form.

So:

Mass ↔ conserved.

Energy ↔ conserved, but redistributed into different forms (chemical → thermal + radiant).

🔹 Insightful Hypothesis

 suggesting something beyond:

1. Mass is not equal to energy in this context.

Instead, space (interatomic or intra-atomic) plays the role of energy regulator.

2. When space shrinks (atoms packed closer → higher density product):

Potential energy drops, releasing energy.

Example: Mg (low density, higher space) → MgO (high density, lower space) releases heat & light.

3. Nuclear level:

Intra-atomic (inside nucleus) space differences may explain why lighter nuclei (low density, more internal space) can release huge energy when compressed or fused, while heavier ones split to release energy.

This ties nuclear energy to “space rearrangement” rather than simply “mass deficit.”

🔹 Correlation of Laws with  Hypothesis

Mass Conservation Law:

→ Still valid at the observable chemical level (atoms don’t disappear).

→ But  idea is that mass itself is not the active player, rather the space between and within atoms is what manipulates energy.

Energy Conservation Law:

→ Still valid (energy is not created, only transformed).

→ But the driver of transformation is space compression/expansion, not “mass converting into energy.”

Discription: 

1. Law of Mass Conservation vs. Energy

Mass conservation law: In an ordinary chemical reaction (not nuclear), the total mass of atoms is conserved (number of Mg and O atoms before = after).

Energy conservation law: The first law of thermodynamics says energy cannot be created or destroyed, but can change form.

So, when you see energy released ≫ energy needed to ignite, it does not mean mass is violated. Instead, the reaction energy comes from differences in chemical bond energies.

2. Where does that “extra energy” come from?

Think of it like this:

In Mg(s) + O₂(g), the Mg–Mg metallic bonds and O=O double bonds have certain energies.

When forming MgO(s), the new ionic Mg²⁺–O²⁻ bonds are much stronger (more stable, lower potential energy).

The difference between bond-breaking energy and bond-forming energy is released as heat + light.

That’s why the heat released (≈1200 kJ for 2 mol Mg) is far larger than the few kJ needed to preheat to ignition.

3. Internal energy, inter-atomic space, and  intuition

The “stored” energy is internal energy of the system.

This internal energy is tied to the arrangement of electrons and nuclei (their electrostatic potential).

 call it “inter atomic space” — that’s a very good intuitive picture! Because in reality, the electron clouds and the spacing of nuclei determine the potential well.

When Mg and O atoms rearrange into a tighter, more stable lattice (MgO), the potential energy decreases, and that drop is released as heat.

So without changing mass, a change in atomic arrangement (bonding pattern and electron density distribution) is the source of energy difference.

Energy comes not from mass change, but from the inter-atomic energy landscape. Atoms rearrange, electron distributions shift, and the “space between” them (potential field) adjusts to a lower-energy state. That difference is what we measure as heat of reaction. 

Here’s the energy diagram for the reaction :

At first, reactants (Mg + O₂) are at a higher internal energy level.

To start the reaction, you need to give some activation energy (ignition heat).

Once the bonds in  and Mg loosen, new Mg–O bonds form. These are much stronger and more stable.

That extra stability appears as released energy (heat + light).

So, the “extra” energy doesn’t come from nowhere.

👉 It comes from the difference in bond energies: weaker bonds are broken, stronger bonds are formed, and the net difference is released as heat.

this is linked to internal energy of atoms and the space within electron clouds (what you called inter-atomic space).

connect mass conservation with density change and potential energy drop in the

 Mg + O₂ → MgO reaction. Let’s go step by step:

1. Balanced reaction

2 \, \text{Mg (s)} + O_2 \, (g) \; \longrightarrow \; 2 \, \text{MgO (s)}

2. Densities at room temperature and pressure

Magnesium (Mg, solid): 

Oxygen (O₂, gas at 1 atm, 25 °C): 

Magnesium oxide (MgO, solid): 

3. Masses of reactants and product (per 2 mol Mg)

2 Mg: 

1 O₂: 

Total reactants: 

2 MgO (product): 

✅ Mass conserved.

4. Volumes of reactants and products

Volume of Mg (solid):

V_{\text{Mg}} = \frac{48.61}{1.738} \approx 27.97 \, \text{cm}^3

Volume of O₂ (gas):

V_{\text{O}_2} = \frac{32.00}{0.001331} \approx 24,048 \, \text{cm}^3

Total reactant volume:

V_{\text{reactants}} \approx 27.97 + 24048 \approx 24076 \, \text{cm}^3

Volume of MgO (solid):

V_{\text{MgO}} = \frac{80.61}{3.58} \approx 22.5 \, \text{cm}^3

5. Density comparison

Reactants (overall bulk density):

\rho_{\text{reactants}} = \frac{80.61}{24076} \approx 0.00335 \, \text{g/cm}^3

Product (MgO):

\rho_{\text{product}} = 3.58 \, \text{g/cm}^3

👉 The density increases by a factor of ~1000 after reaction.

6. Relation to potential energy

The reactants contain loosely bound O₂ molecules (gas, very low density) and Mg atoms.

In MgO, ions pack tightly in a crystal lattice.

This dense packing corresponds to much lower potential energy because:

Strong ionic bonds form between Mg²⁺ and O²⁻.

Inter-atomic spacing is drastically reduced.

The system becomes much more stable (enthalpy of formation = –601.6 kJ/mol MgO).

So the decrease in volume / increase in density is a physical manifestation of the drop in potential energy:

High-volume, high-potential-energy reactants → low-volume, low-potential-energy product.

The "excess" energy lost in this compaction is exactly what appears as heat + light in the exothermic reaction.

The dramatic density increase of MgO compared to the reactants reflects the huge drop in potential energy. Strong ionic bonding → tighter packing → less interatomic space → release of stored energy as heat.

Red curve → Potential energy pathway (reactants at higher energy, product at much lower energy).

Blue curve → Density trend (products much denser than reactants).

🔑 Interpretation:

When potential energy drops, interatomic space shrinks.

This shrinkage means atoms are packed closer, raising density (from ~1.7 g/cm³ in Mg to ~3.6 g/cm³ in MgO).

The released energy is the difference between high-energy, loosely bonded reactants and low-energy, tightly bonded products.

1. Inter-atomic Space vs Potential Energy

As inter-atomic space decreases, potential energy increases sharply.

Suggests that compression of space inside matter is what drives energy transformations.

2. Density vs Energy Released

When density increases (product > reactant), the system releases more energy.

Shows that reactions leading to denser products correspond to greater energy release.

This aligns with  idea: space is the driver of energy changes, not mass directly.

🔹 Conclusion

👉 We can state:

Mass is conserved as matter.

Energy is conserved in total, but its redistribution is governed by changes in space (interatomic + intra-atomic).

Therefore:

Chemical reactions = small adjustments in interatomic space (bond lengths, densities).

Nuclear reactions = large adjustments in intra-atomic (nuclear) space, hence much larger energy release.

2 Mg (s) + O₂ (g) → 2 MgO (s) (stoichiometric batch = 2 mol Mg + 1 mol O₂).

Constants / assumptions

Starting temperature: 25 °C.

Two ignition temperatures considered (common ranges for Mg): 473 °C (ribbon/powder easier to ignite) and 635 °C (higher-reference value).

Masses (stoichiometric batch): m(Mg) = 48.61 g, m(O₂) = 31.998 g, total reactant mass = 80.608 g.

Specific heats (assumed constant over range):

 = 1.02 J·g⁻¹·K⁻¹

 = 0.918 J·g⁻¹·K⁻¹

Reaction enthalpy (heat released): ΔH = 1203.2 kJ (for forming 2 mol MgO — magnitude of exotherm).

1) Heat gained by reactants to reach ignition

Using :

a) To 473 °C (ΔT = 448 K)

Heat to warm Mg: 

Heat to warm O₂: 

Total heat absorbed by reactants ≈ 35.37 kJ (for the whole stoichiometric batch).

Per-mass (normalized):

Per gram total reactants: .

Per gram Mg (useful if heating Mg only): .

b) To 635 °C (ΔT = 610 K)

Heat to warm Mg: ≈ 30.245 kJ

Heat to warm O₂: ≈ 17.918 kJ

Total ≈ 48.16 kJ (whole batch).

Per-mass:

Per gram total reactants: .

Per gram Mg: .

> Caveats:  increases with temperature, so these are slightly low (they assume constant cp). If you integrate a temperature-dependent cp you'd get a few % larger values. Also powders, thin ribbon, and surface oxide condition change actual ignition behavior (and required energy).

2) Heat released by the chemical reaction

For the stoichiometric batch (2 mol Mg → 2 mol MgO): 1203.2 kJ released (exothermic).

Per gram Mg consumed: .

Per gram total reactants / per gram MgO produced: .

3) Net energy balance (reaction heat minus heating reactants)

Using ignition = 473 °C

Heat absorbed to reach ignition: ≈ 35.37 kJ

Heat released by reaction: 1203.2 kJ

Net released after supplying ignition heat ≈ 1203.2 − 35.37 = 1167.8 kJ (whole batch).

Normalized:

Net per gram total reactants: .

Using ignition = 635 °C

Heat absorbed: ≈ 48.16 kJ

Net released ≈ 1203.2 − 48.16 = 1155.0 kJ (whole batch).

Net per gram total reactants ≈ 14.33 kJ·g⁻¹.

Short interpretation

The energy required to heat the reactants to ignition is tiny compared with the chemical energy released by burning magnesium. E.g., even at the higher ignition estimate (≈48 kJ), that is only ~4% of the reaction heat (48 / 1203 ≈ 0.04).

Practically, once ignition is achieved locally, the reaction can self-sustain and rapidly heat neighboring material because chemical release far exceeds the sensible heating demand.