Molecule Polarity In 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 a two-atom molecule, polarity is determined by the difference in electronegativity between the two atoms. You can change the polarity by:

• Changing the electronegativity of atom A.
• Changing the electronegativity of atom B.
• The greater the difference in electronegativity, the more polar the bond (and thus the molecule).

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

The representations help understand molecule polarity by visually depicting the unequal sharing of electrons and the resulting charge distribution:

• Partial charges ($\delta^+$ and $\delta^-$): These indicate which atom has a slight positive charge (less electronegative) and which has a slight negative charge (more electronegative) due to electron displacement.
• Bond dipole arrow: This arrow points from the less electronegative (partially positive) atom to the more electronegative (partially negative) atom, showing the direction of electron density shift within the bond.
• Electrostatic potential map/surface: This uses a color gradient (e.g., red for electron-rich/negative, blue for electron-poor/positive) to show the overall distribution of electron density across the molecule, highlighting regions of partial positive and negative charge.

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

In a three-atom molecule, polarity depends on both the polarity of individual bonds and the molecule's geometry. New ways to change polarity compared to a two-atom molecule include:

• Changing the bond angles/molecular geometry: Even if individual bonds are polar, a symmetrical arrangement can cause bond dipoles to cancel out. Changing the angle (e.g., from linear to bent) can prevent cancellation and make the molecule polar.
• Changing the electronegativity of the central atom or the outer atoms: Similar to the two-atom case, altering the electronegativity of any of the three atoms will affect the polarity of the individual bonds, which in turn influences the overall molecular polarity.

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

A bond dipole represents the polarity of an individual bond due to the electronegativity difference between the two atoms forming that bond; it is a vector pointing from the less electronegative to the more electronegative atom. The molecular dipole is the vector sum of all individual bond dipoles in the molecule. If the bond dipoles are arranged symmetrically such that they cancel each other out, the molecule will have a net molecular dipole of zero and be nonpolar. If the bond dipoles do not cancel due to an asymmetrical arrangement, the molecule will have a net molecular dipole and be considered 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 molecule has a symmetrical geometry that causes the individual bond dipoles to cancel each other out, resulting in a net molecular dipole of zero.

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

• Each $\ce{C=O}$ bond is polar because oxygen is significantly more electronegative than carbon, creating a bond dipole pointing from carbon towards each oxygen.
• However, $\ce{CO2}$ has a linear molecular geometry. The two $\ce{C=O}$ bond dipoles are equal in magnitude and point in opposite directions, effectively canceling each other out.
• Therefore, despite having two polar bonds, the overall $\ce{CO2}$ molecule is nonpolar.

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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.

As an AI, I cannot physically draw molecules or interact with a simulation. However, I can describe the process for predicting polarity for a few common real molecules, including how one would draw the dipoles and explain the reasoning.

Molecule 1: Water ($\ce{H2O}$)

• Drawing/Geometry: Water has a bent molecular geometry (approximately 104.5° bond angle) due to two lone pairs on the central oxygen atom.
• Bond Dipoles: The $\ce{O-H}$ bonds are polar because oxygen is more electronegative than hydrogen. Bond dipoles point from each hydrogen towards the oxygen.
• Molecular Dipole: Because of the bent geometry, the two $\ce{O-H}$ bond dipoles do not cancel. They add up to create a net molecular dipole moment pointing from the region between the hydrogens towards the oxygen atom.
• Reasoning: Water has polar bonds and an asymmetrical (bent) molecular geometry, leading to a net molecular dipole.
• Prediction: Polar.

Molecule 2: Carbon Tetrachloride ($\ce{CCl4}$)

• Drawing/Geometry: Carbon tetrachloride has a tetrahedral molecular geometry, with the carbon atom at the center and four chlorine atoms at the vertices.
• Bond Dipoles: The $\ce{C-Cl}$ bonds are polar because chlorine is more electronegative than carbon. Bond dipoles point from the central carbon towards each chlorine atom.
• Molecular Dipole: The four $\ce{C-Cl}$ bond dipoles are equal in magnitude and are arranged symmetrically in a tetrahedral shape. Due to this symmetry, the vector sum of these bond dipoles is zero.
• Reasoning: Carbon tetrachloride has polar bonds, but its highly symmetrical tetrahedral geometry causes all bond dipoles to cancel out.
• Prediction: Nonpolar.

Molecule 3: Ammonia ($\ce{NH3}$)

• Drawing/Geometry: Ammonia has a trigonal pyramidal molecular geometry, with the nitrogen atom at the apex and three hydrogen atoms forming the base, plus one lone pair on the nitrogen.
• Bond Dipoles: The $\ce{N-H}$ bonds are polar because nitrogen is more electronegative than hydrogen. Bond dipoles point from each hydrogen towards the nitrogen.
• Molecular Dipole: Due to the trigonal pyramidal geometry and the presence of the lone pair on nitrogen (which contributes to electron density on that side), the three $\ce{N-H}$ bond dipoles do not cancel. They add up to create a net molecular dipole moment pointing from the base of hydrogens towards the nitrogen atom (and the lone pair).
• Reasoning: Ammonia has polar bonds and an asymmetrical (trigonal pyramidal) molecular geometry, leading to a net molecular dipole.
• Prediction: Polar.