Answer :
To understand why acetone and sodium chloride have different properties despite having a similar mass, we need to take a closer look at their chemical nature and structure. Let's go through the points step-by-step to explain the differences.
1. Chemical Nature:
- Acetone ([tex]$C_3H_6O$[/tex]): Acetone is a covalent compound. Covalent compounds are formed by the sharing of electrons between non-metal atoms. In acetone, the carbon, hydrogen, and oxygen atoms share electrons to form stable molecules.
- Sodium Chloride ([tex]$NaCl$[/tex]): Sodium chloride is an ionic compound. Ionic compounds are formed by the transfer of electrons from one atom to another, resulting in a bond between positively charged cations (Na^+) and negatively charged anions (Cl^-).
2. Melting Point:
- Acetone: Acetone has a very low melting point of [tex]\(-94^{\circ} C\)[/tex]. This low melting point is due to the weak intermolecular forces (such as van der Waals forces) between the acetone molecules. Covalent compounds typically have low melting points because their molecules are held together by relatively weak forces.
- Sodium Chloride: Sodium chloride has a high melting point of [tex]\(801^{\circ} C\)[/tex]. The higher melting point is due to the strong electrostatic forces between the sodium and chloride ions in its crystal lattice. Ionic compounds generally have high melting points because the ionic bonds require significant energy to break.
3. State at Room Temperature:
- Acetone: Given its low melting point, acetone is a liquid at room temperature. This aligns with its covalent nature and the weak intermolecular forces that do not require much energy to overcome.
- Sodium Chloride: Sodium chloride is a solid at room temperature because of its high melting point and the strong ionic bonds maintaining its solid crystalline structure.
4. Electrical Conductivity:
- Acetone: Acetone has low electrical conductivity as it does not have free ions or electrons moving freely. Covalent compounds typically do not conduct electricity well in their pure form.
- Sodium Chloride: Sodium chloride has high electrical conductivity when molten or dissolved in water because the ions are free to move and carry electric charge. However, in solid form, the ions are fixed in the lattice and do not conduct electricity.
Given these points, we can now check the statements:
1. Acetone is a covalent compound, while sodium chloride is an ionic compound.
- This statement is true. (From our analysis of the chemical nature of the compounds)
2. Ionic compounds conduct electricity well in their solid (pure) form.
- This statement is false. (Ionic compounds conduct electricity well when dissolved in water or molten, but not in solid form)
3. Covalent compounds have weak attractions between molecules, resulting in low melting points.
- This statement is true. (We see this with acetone's low melting point and weak intermolecular forces)
Therefore, we conclude that the statements "Acetone is a covalent compound, while sodium chloride is an ionic compound" and "Covalent compounds have weak attractions between molecules, resulting in low melting points" are true. The statement "Ionic compounds conduct electricity well in their solid (pure) form" is false.
1. Chemical Nature:
- Acetone ([tex]$C_3H_6O$[/tex]): Acetone is a covalent compound. Covalent compounds are formed by the sharing of electrons between non-metal atoms. In acetone, the carbon, hydrogen, and oxygen atoms share electrons to form stable molecules.
- Sodium Chloride ([tex]$NaCl$[/tex]): Sodium chloride is an ionic compound. Ionic compounds are formed by the transfer of electrons from one atom to another, resulting in a bond between positively charged cations (Na^+) and negatively charged anions (Cl^-).
2. Melting Point:
- Acetone: Acetone has a very low melting point of [tex]\(-94^{\circ} C\)[/tex]. This low melting point is due to the weak intermolecular forces (such as van der Waals forces) between the acetone molecules. Covalent compounds typically have low melting points because their molecules are held together by relatively weak forces.
- Sodium Chloride: Sodium chloride has a high melting point of [tex]\(801^{\circ} C\)[/tex]. The higher melting point is due to the strong electrostatic forces between the sodium and chloride ions in its crystal lattice. Ionic compounds generally have high melting points because the ionic bonds require significant energy to break.
3. State at Room Temperature:
- Acetone: Given its low melting point, acetone is a liquid at room temperature. This aligns with its covalent nature and the weak intermolecular forces that do not require much energy to overcome.
- Sodium Chloride: Sodium chloride is a solid at room temperature because of its high melting point and the strong ionic bonds maintaining its solid crystalline structure.
4. Electrical Conductivity:
- Acetone: Acetone has low electrical conductivity as it does not have free ions or electrons moving freely. Covalent compounds typically do not conduct electricity well in their pure form.
- Sodium Chloride: Sodium chloride has high electrical conductivity when molten or dissolved in water because the ions are free to move and carry electric charge. However, in solid form, the ions are fixed in the lattice and do not conduct electricity.
Given these points, we can now check the statements:
1. Acetone is a covalent compound, while sodium chloride is an ionic compound.
- This statement is true. (From our analysis of the chemical nature of the compounds)
2. Ionic compounds conduct electricity well in their solid (pure) form.
- This statement is false. (Ionic compounds conduct electricity well when dissolved in water or molten, but not in solid form)
3. Covalent compounds have weak attractions between molecules, resulting in low melting points.
- This statement is true. (We see this with acetone's low melting point and weak intermolecular forces)
Therefore, we conclude that the statements "Acetone is a covalent compound, while sodium chloride is an ionic compound" and "Covalent compounds have weak attractions between molecules, resulting in low melting points" are true. The statement "Ionic compounds conduct electricity well in their solid (pure) form" is false.
