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Concept of Coulomb's Law
F = k(|q₁ · q₂| ÷ r²)
F: Electrostatic 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. However, this law shows how 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.
If F < 0, the force is attractive (lower potential energy).
If F > 0, the force is repulsive (higher potential energy).
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. 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
The minima of the curve represents the optimal bond length, as it possesses the lowest potential energy.
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 are the smallest.
Low minima for bond order due to strong 1s orbitals.
N₂ (Blue Curve):
Shorter bond length and lower minima than O₂ and F₂ due to triple 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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