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Relation of Charge and Distance to Force:
Coulomb's Law (Scalar):
F = k(|q₁ · q₂| ÷ r²)
F: Electrostatic Force
Higher value → Stronger force.
k: Coulomb's Constant (8.99 × 10⁹ N·m²/C²)
q₁ and q₂: Magnitudes of the two charges
r: Distance between sources of the charges
You are not required to solve problems with Coulomb's Law in AP Chemistry.
This law shows how the strength of electrostatic forces are proportional to the difference and distance between the charges.
What to Understand:
Opposite charges (positive and negative) atrract.
Like charges repel.
Greater charge magnitudes result in stronger forces.
Smaller distances between charges result in stronger forces.
Coulomb's Law helps explain concepts like lattice energy (the energy required to break ionic lattices) and bond length.
Together with Fajan's rule, they can influence how soluble ionic compounds are.
Potential Energy and Length of Bonds
Atoms have positively-charged nuclei and negatively-charged electrons. During bonding, an atom's electrons experience an attractive force toward the nucleus of another atom, and vice versa. However, if atoms get too close, repulsive forces between their electron clouds and nuclei increase.
More attraction → Lower potential energy (system is more stable).
More repulsion → Higher potential energy (system is less stable).
Therefore, there is an optimal distance where the potential energy of a bond can be minimized by balancing the objectives of maximizing attraction while reducing repulsion.
What to Remember:
Smaller atoms form shorter bond lengths.
Higher bond order means shorter bond lengths and lower potential energy. (Triple < Double < Single)
Greater charge magnitudes lead to shorter ionic bonds.
Interpreting Bond Energy Graphs
Bond energy graphs show a comparison between the internuclear distance and the bond's potential energy.
If two atoms/ions are too close, they experience more repulsion.
If two atoms/ions are too far, they experience less attraction than ideal.
The minima of the curve represents the optimal bond length, as lower potential energy means more stability.
According to the y-axis, the lower the minima, the stronger the bond.
Comparison Between Bonds
H₂ (Grey Curve):
Shortest bond length, as hydrogen atoms have the smallest radius (only possessing 1s orbitals).
N₂ (Blue Curve):
Shorter bond length and lower minima than O₂ and F₂ due to triple bond, most stable bond.
O₂ (Red Curve):
Shorter bond length and lower minima than F₂ due to double bond.
F₂ (Lime Curve):
Longest bond length due to lowest bond order among period 2 diatomics.
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