Covalent vs Ionic Bonding and Energy Curves — Transcript Notes
Electronegativity and Covalent Bonding
Definition: electronegativity is a relative measure of an atom's ability to attract electrons. It describes how strongly an atom holds onto electrons or is willing to attract more electrons.
Covalent bonding basics:
Occurs typically between atoms with similar electronegativities, because they are willing to share electrons.
Example discussed: carbon (in CH"); carbon shares electrons with hydrogen atoms to achieve a more stable electron configuration.
Carbon and bonding basics relevant to biology:
Biological molecules are predominantly carbon-based and rely on covalent bonds for structure and function.
Covalent bonds stabilize molecules by sharing electrons rather than transferring them.
Covalent Bonding in Detail
Carbon valence and octet context:
Carbon has an s orbital and p orbitals that can hold a total of eight electrons in its outer shell when fully shared.
In CH
t (methane, CH extsubscript{4}), carbon shares electrons with four hydrogen atoms, allowing carbon to fulfill the octet rule per valence considerations.
General intro chem reference: covalent bonding is illustrated here with methane (CH extsubscript{4}) as an introductory example. If you need deeper review, resources can be provided.
Electronegativity: Examples and Implications
Fluorine as the most electronegative element: it highly desires one more electron to complete its outer shell.
Atoms on the far end of the spectrum that readily give up electrons tend to form ionic bonds when paired with highly electronegative partners.
Ionic Bonding
Definition: ionic bonding is a bond between a positive ion and a negative ion created by an actual transfer of electrons.
Key requirement: a large difference in electronegativity is typically necessary to drive electron transfer.
Classic example: sodium chloride (NaCl).
Sodium tends to lose an electron to become Na extsuperscript{+} (a stable cation).
Chlorine gains an electron to become Cl extsuperscript{−} (a stable anion).
The electrostatic attraction between Na extsuperscript{+} and Cl extsuperscript{−} forms the ionic bond.
Note: ionic bonding tends to occur between atoms that are sufficiently far apart on the periodic table in terms of electronegativity, leading to electron transfer rather than sharing.
Interatomic Potential Energy and Bond Lengths
Energy vs distance plot (conceptual):
Y-axis: potential energy (in kJ/mol).
X-axis: internuclear distance (in picometers, pm).
At large separation: atoms interact negligibly, E ≈ 0.
As distance decreases: attractive forces dominate at first, lowering the energy.
At an optimal distance: a balance of attractive and repulsive forces yields a minimum in energy.
If atoms are pushed even closer: repulsive forces dominate and energy rises again.
Bond length and minimum energy:
The interatomic spacing at the energy minimum is the bond length (r extsubscript{bond}).
Example: H extsubscript{2} bond length is ~
r{ ext{H}2} = 74 ext{ pm}.The minimum energy corresponds to the most energetically favorable bonding arrangement.
Energetic stability concept:
Energetically favorable states correspond to the lowest possible energy on the curve.
To break a bond, energy must be supplied; to form a bond, energy is released.
Energy of Bonds: Breaking vs. Forming
Key rule to avoid common confusion:
Breaking a bond requires energy input (endothermic): the dissociation energy is positive.
Forming a bond releases energy (exothermic): energy is released as the bond forms.
Quantitative example: H extsubscript{2} bond energy (bond dissociation energy), typically reported as
Approximately
ext{BDE}_{ ext{H–H}}
ightarrow ext{around } 432 ext{ to } 436 ext{ kJ mol}^{-1}.This is the energy required to break one mole of H–H bonds into 2 separate H atoms.
Important nuance: real systems are not isolated single bonds in vacuum.
In ATP hydrolysis context, breaking a phosphoanhydride bond alone requires energy input, but net energy release in cells arises from subsequent steps (e.g., hydration and other reactions) modulated by solvent effects and hydrogen-bond rearrangements.
The takeaway: energy to break a bond is not the same as the energy released during the full process in solution; the environment matters significantly.
Conceptual takeaway:
When a bond forms, energy is released; when a bond breaks, energy is absorbed.
The magnitude of the energy change depends on the specific bond type and the surrounding chemical environment.
Force and Energy Derivatives (Conceptual)
Relationship between energy and force:
The interatomic force is the negative derivative of the potential energy with respect to distance:
F(r) = -\frac{dE}{dr}.
At the energy minimum:
The slope of the energy curve is zero: \frac{dE}{dr} = 0.
Consequently, the force is zero: F(r) = 0.
Force curve intuition:
To the left of the minimum, the slope is negative, corresponding to a positive force that pulls the atoms together.
To the right of the minimum, the slope is positive, corresponding to a repulsive or pulling force that pushes the atoms apart.
Sign convention for energy:
Potential energies are often plotted as negative values for bonded states; more negative means a more stable, lower-energy state.
When comparing energies, “less negative” corresponds to a higher energy and thus a less stable configuration.
H–H Bond vs. Other Covalent Bonds (Context for Practice Question)
The class exercise asks you to compare which bond is more energetically favorable: C–C vs C–H.
General guidance (based on typical bond energies):
C–H bond energy is typically around E_{ ext{C–H}} \approx 413 \ ext{kJ mol}^{-1}.
C–C single bond energy is typically around E_{ ext{C–C}} \approx 348 \ ext{kJ mol}^{-1}.
Therefore, C–H bonds are generally stronger (harder to break) and more energetically favorable to form than C–C single bonds, all else being equal.
Implication for energy curves:
The energy minimum for a C–H bond would be deeper (more negative) and at a shorter bond length compared to a C–C bond, reflecting greater stability and larger dissociation energy.
Practice Question and Suggested Sketch
Question: Which bond is more energetically favorable: C–C or C–H? Which one is harder to break?
Answer (based on typical bond energies): C–H bonds are stronger; C–H is harder to break than C–C, given the larger dissociation energy.
Suggested exercise: sketch an energy vs distance curve for both bonds on the same axes.
Axis scales (as suggested):
Energy: from 0 to negative 500 kJ/mol
Distance: from 0 to 200 pm
Both curves should show a minimum at their respective bond lengths, with the C–H curve reaching a more negative energy minimum and typically shorter equilibrium distance than the C–C curve.
Indicate the dissociation energy as the vertical distance from the minimum to zero energy, and show how the slope changes as you move away from the minimum (derivative crossing zero at the minimum).
Quick Connections and Takeaways
Covalent bonds dominate in carbon-based biology due to favorable electronegativity balance and the desire to complete octets via sharing.
Ionic bonds rely on electron transfer and strong electrostatic attractions between ions with large electronegativity differences; example NaCl illustrates the concept.
Understanding energy curves helps explain bond lengths, bond strengths, and the nature of chemical reactivity.
Real biochemical processes (e.g., ATP hydrolysis) are context-dependent; environment and subsequent reactions significantly influence net energy changes, not just the intrinsic bond dissociation energy.
Key Definitions and Formulas (Recap)
Electronegativity: a relative measure of an atom's ability to attract electrons.
Covalent bond: sharing of electrons between atoms of similar electronegativity.
Ionic bond: bond formed by transfer of electrons between ions of opposite charge.
Bond energy (dissociation energy): energy required to break a bond in a mole of bonds, units \text{kJ mol}^{-1}.
Bond length: the internuclear distance at which the potential energy is minimized, i.e., the bond minimum.
Potential energy curve: E(r) as a function of interatomic distance r, with a minimum at the bond length.
Force from energy: F(r) = -\frac{dE}{dr}.
Sign convention: more negative E indicates a more stable bond; energy to break is positive; energy released on bond formation is negative in the sense of energy change of the system.
Note on Units and Notation
Distances: picometers (pm).
Energies: kilojoules per mole (kJ mol$^{-1}$).
Example values used:
H–H bond length: r{ ext{H}2} = 74\,\text{pm}.
H–H bond dissociation energy: \Delta E_{ ext{diss}} \approx 432 \text{ to } 436\ \text{kJ mol}^{-1}.
C–H bond energy: approximately E_{ ext{C–H}} \approx 413\ \text{kJ mol}^{-1}.
C–C single bond energy: approximately E_{ ext{C–C}} \approx 348\ \text{kJ mol}^{-1}.