n this activity you will use a PhET simulation to explore molecule polarity. Part I: What factors affect molecule polarity? 1. Explore the Molecule Polarity simulation for a few minutes with a partner. In each of the three tabs, try to find all of the controls and figure out how they work. Two Atoms tab 2. Describe all of the ways you can change the polarity of the two-atom molecule. 3. Explain how the representations below help you understand molecule polarity. Three Atoms tab 4. Describe any new ways you can change the polarity of the three-atom molecule. 5. Explain the relationship between the bond dipoles and the molecular dipole. Page 2 of 2 6. Can a non-polar molecule contain polar bonds? Explain your answer with an example. Real Molecules tab 7. Predict the polarity of 6 real molecules. First, draw the molecules and any bond dipoles. Then draw any molecular dipoles. Explain your reasoning before you check your predictions with the simulation.

Question 1: 2. Describe all of the ways you can change the polarity of the two-atom molecule.

In the "Two Atoms" tab of the PhET simulation, you can change the polarity of the two-atom molecule by:

• Adjusting the electronegativity of Atom A and Atom B: A larger difference in electronegativity between the two atoms leads to a more polar bond and thus a more polar molecule. If the electronegativities are equal, the molecule is nonpolar.
• Changing the bond length: While the primary determinant of polarity is electronegativity difference, the magnitude of the dipole moment (which quantifies polarity) is also dependent on the distance between the partial charges. Increasing the bond length for a given charge separation will increase the dipole moment.

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Question 2: 3. Explain how the representations below help you understand molecule polarity.

The representations in the PhET simulation help understand molecule polarity by visually depicting the distribution of electron density:

• Partial Charges ($\delta^+$ and $\delta^-$): These labels directly show which atom has a partial positive charge (less electron density) and which has a partial negative charge (more electron density) due to unequal sharing of electrons.
• Bond Dipole Arrow: This arrow points from the partially positive end to the partially negative end of the bond, indicating the direction and magnitude of electron shift within the bond. A longer arrow signifies a stronger bond dipole.
• Electrostatic Potential Map (or Charge Distribution): This color-coded surface (often red for negative, blue for positive) provides a visual representation of the overall electron distribution around the molecule. Areas with higher electron density appear red/negative, and areas with lower electron density appear blue/positive, clearly showing regions of charge separation.

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Question 3: 4. Describe any new ways you can change the polarity of the three-atom molecule.

In the "Three Atoms" tab, in addition to changing the electronegativity of the individual atoms (similar to the two-atom case), new ways to change the polarity of the molecule include:

• Adjusting the molecular geometry/bond angle: Even if individual bonds are polar, the overall molecular polarity depends on the spatial arrangement of these bonds. By changing the bond angle, you can alter whether the individual bond dipoles cancel each other out or sum up to create a net molecular dipole.
• Changing the electronegativity of the central atom relative to the terminal atoms: This allows for more complex scenarios where the central atom can be more or less electronegative than the surrounding atoms, influencing the direction and magnitude of multiple bond dipoles.

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Question 4: 5. Explain the relationship between the bond dipoles and the molecular dipole.

The relationship between bond dipoles and the molecular dipole is that the molecular dipole is the vector sum of all individual bond dipoles within the molecule.

• A bond dipole arises from the unequal sharing of electrons between two atoms in a covalent bond due to a difference in their electronegativities. It is a vector quantity, having both magnitude and direction.
• The molecular dipole represents the overall polarity of the entire molecule. If the individual bond dipoles are arranged symmetrically such that they cancel each other out (like equal and opposite vectors), the molecule will have no net molecular dipole and will be nonpolar. If the bond dipoles do not cancel due to asymmetry in the molecular geometry, they will sum to create a net molecular dipole, making the molecule polar.

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Question 5: 6. Can a non-polar molecule contain polar bonds? Explain your answer with an example.

Yes, a non-polar molecule can contain polar bonds. This occurs when the molecular geometry is symmetrical, causing the individual bond dipoles to cancel each other out, resulting in a net molecular dipole of zero.

Example: Carbon Dioxide ($\ce{CO2}$)

• Polar Bonds: Each $\ce{C=O}$ bond is polar because oxygen is significantly more electronegative than carbon, pulling electron density towards itself. This creates a bond dipole pointing from carbon towards each oxygen.
• Non-polar Molecule: Carbon dioxide has a linear molecular geometry. The two $\ce{C=O}$ bond dipoles are equal in magnitude and point in opposite directions ($\leftarrow \ce{O=C=O} \rightarrow$). Because they are equal and opposite, they perfectly cancel each other out, leading to a net molecular dipole of zero. Therefore, $\ce{CO2}$ is a non-polar molecule despite having polar bonds.

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Question 6: 7. Predict the polarity of 6 real molecules. First, draw the molecules and any bond dipoles. Then draw any molecular dipoles. Explain your reasoning before you check your predictions with the simulation.

Here are predictions for the polarity of six real molecules:

1. Water ($\ce{H2O}$)

• Drawing with Bond Dipoles: $$\ce{H} - \overset{\delta^-}{\ce{O}} - \ce{H}$$ (Arrows point from H to O for each bond)
• Molecular Geometry: Bent (V-shaped)
• Molecular Dipole: $$\ce{H} - \overset{\delta^-}{\ce{O}} - \ce{H}$$ (Net dipole points upwards, bisecting the H-O-H angle)
• Reasoning: Oxygen is significantly more electronegative than hydrogen, making each $\ce{O-H}$ bond polar. Due to the bent molecular geometry, the two bond dipoles do not cancel out; instead, they add