- When comparing the bond formation in MgO versus H₂O, which analysis correctly describes the fundamental differences in their bonding mechanisms?
Answer: MgO forms through complete electron transfer creating discrete ions, while H₂O forms through electron sharing with partial charges
MgO forms ionic bonds through complete electron transfer from Mg (forming Mg²⁺) to O (forming O²⁻), creating discrete ions held by electrostatic attraction. H₂O forms polar covalent bonds where electrons are shared unequally, creating partial charges (δ+ on H, δ- on O). Option B incorrectly classifies MgO as covalent. Option C reverses the bond types. Option D incorrectly suggests H₂O involves electron transfer and ionic bonding.
- A compound exhibits the following properties: dissolves in water to conduct electricity, has a melting point of 850°C, and forms crystals that cleave along specific planes. When the electronegativity values are examined, the difference between the bonding atoms is 1.8. What conclusion about the bonding is most appropriate?
Answer: The compound is ionic because the physical properties are consistent with ionic bonding despite the moderate electronegativity difference
When you encounter questions about chemical bonding, remember that you must consider both electronegativity differences AND physical properties together to make accurate classifications.
The correct answer is D because the physical properties strongly indicate ionic bonding. The compound dissolves in water and conducts electricity (indicating ion formation), has a high melting point of 850°C (showing strong electrostatic forces), and forms crystals with specific cleavage planes (typical of ionic lattice structures). While the electronegativity difference of 1.8 is moderate, these compelling physical properties override the borderline electronegativity value.
Option A is incorrect because it relies too rigidly on the "2.0 rule" for electronegativity differences. This rule is a guideline, not an absolute threshold, and physical evidence should take precedence when properties clearly indicate ionic character.
Option B is wrong because metallic bonding doesn't explain the water solubility or the specific crystalline cleavage patterns described. Metals conduct electricity in solid state, but this compound only conducts when dissolved, indicating ion mobility rather than electron sea conduction.
Option C is incorrect because the physical properties aren't ambiguous—they clearly point toward ionic bonding. The evidence is sufficient for classification despite the moderate electronegativity difference.
Remember for the HESI: Physical properties are often more reliable indicators of bonding type than electronegativity differences alone. When you see high melting points, water solubility with electrical conductivity, and crystalline structure together, think ionic bonding regardless of borderline electronegativity values.
- A student is examining the structure of ammonium nitrate (NH₄NO₃). To completely describe the bonding in this compound, the student must identify:
Answer: covalent bonds within each polyatomic ion and an ionic bond between the two ions.
When analyzing the bonding in complex ionic compounds like ammonium nitrate (NH₄NO₃), you need to examine bonding at two levels: within the polyatomic ions themselves and between the ions.
Ammonium nitrate contains two polyatomic ions: ammonium (NH₄⁺) and nitrate (NO₃⁻). Within each polyatomic ion, the atoms are held together by covalent bonds. In NH₄⁺, nitrogen shares electrons with four hydrogen atoms through covalent bonding. Similarly, in NO₃⁻, nitrogen forms covalent bonds with three oxygen atoms. However, these two polyatomic ions are held together by ionic attraction - the positively charged ammonium ion is electrostatically attracted to the negatively charged nitrate ion.
Answer C correctly identifies both types of bonding: covalent bonds within each polyatomic ion and an ionic bond between the two ions.
Answer A is incomplete because it only identifies the ionic interaction between the polyatomic ions while ignoring the covalent bonding within each ion. Answer B makes the opposite error - it recognizes the internal covalent bonding but misses the ionic attraction that holds the compound together. Answer D incorrectly suggests metallic bonding, which occurs in metals where electrons move freely in an "electron sea," not in ionic compounds with discrete polyatomic ions.
For HESI chemistry questions involving polyatomic ions, always consider bonding at multiple levels. Remember that polyatomic ions are covalently bonded groups of atoms that carry a charge, and these charged groups interact ionically with other ions to form compounds.
- In a molecule of hydrogen fluoride (HF), a bond is formed between a hydrogen atom (electronegativity ≈ 2.2) and a fluorine atom (electronegativity ≈ 4.0). How is this bond best classified?
Answer: Polar covalent, because the significant electronegativity difference leads to unequal sharing.
When you encounter questions about chemical bonding, the key is determining bond type based on electronegativity differences between atoms. Electronegativity measures an atom's ability to attract electrons in a bond.
To classify the HF bond, calculate the electronegativity difference: 4.0−2.2=1.8. This significant difference means fluorine pulls the shared electrons much more strongly than hydrogen, creating an unequal sharing situation. The electrons spend more time near fluorine, making it partially negative (δ-) and hydrogen partially positive (δ+). This unequal electron sharing defines a polar covalent bond.
Choice A incorrectly classifies this as ionic. While the 1.8 difference exceeds the common 1.7 threshold, HF doesn't form discrete ions. Instead, the atoms share electrons unequally. The 1.7 rule is a guideline, not an absolute cutoff, and bonding exists on a spectrum.
Choice C misapplies the nonmetal rule. While both atoms are nonmetals, this doesn't automatically mean nonpolar covalent bonding. Nonpolar covalent bonds occur when electronegativity differences are very small (typically less than 0.4), resulting in equal electron sharing.
Choice D incorrectly suggests metallic bonding. Although hydrogen is in Group 1, it behaves as a nonmetal in most compounds. Metallic bonding involves delocalized electrons in a "sea" around metal atoms, which doesn't describe HF.
Remember: electronegativity differences determine bond polarity. Small differences (< 0.4) suggest nonpolar covalent, moderate differences (0.4-1.7) indicate polar covalent, and large differences (> 1.7) typically suggest ionic character. - Diamond is a form of pure carbon. It is a solid with an extremely high melting point, similar to ionic compounds, but it is a very poor electrical conductor. What type of bonding best explains these properties?
Answer: Covalent network bonding, where atoms are held in a large lattice by covalent bonds.
When you encounter questions about material properties like melting point, hardness, and electrical conductivity, you need to connect these macroscopic observations to the underlying bonding and structure at the atomic level.
Diamond's unique properties stem from its covalent network structure. In diamond, each carbon atom forms four strong covalent bonds with neighboring carbon atoms, creating a three-dimensional lattice that extends throughout the entire crystal. This extensive network of covalent bonds explains why diamond has such a high melting point—you'd need to break countless strong C-C bonds to melt it. The poor electrical conductivity occurs because all electrons are localized in covalent bonds between atoms, with no mobile electrons available to carry current.
Answer A incorrectly suggests ionic bonding. While ionic compounds do have high melting points, carbon atoms have similar electronegativity values, so they share electrons rather than transfer them. Answer C proposes metallic bonding, but carbon lacks the "sea" of delocalized electrons characteristic of metals—in fact, this absence explains diamond's poor conductivity. Answer D suggests separate C₄ molecules, which would actually result in a soft material with weak intermolecular forces, the opposite of diamond's properties.
The key strategy here is recognizing that extreme properties often indicate network structures. When you see descriptions of materials that are exceptionally hard, have very high melting points, but poor electrical conductivity, think covalent network solids like diamond, quartz, or silicon carbide.
- A molecule of carbon dioxide (CO₂) contains two polar covalent bonds between carbon and oxygen. However, the molecule itself is nonpolar. What is the best explanation for this observation?
Answer: The molecule is linear, and the two bond dipoles cancel each other out due to symmetry.
When you encounter questions about molecular polarity, remember that molecular shape determines whether individual polar bonds create an overall polar or nonpolar molecule. This concept is crucial for understanding how molecules interact with each other and with solvents.
Carbon dioxide's linear geometry is the key to understanding this apparent contradiction. The molecule has a straight-line arrangement: O=C=O, with a 180° bond angle. Each C=O bond is indeed polar because oxygen is more electronegative than carbon, creating partial charges (δ+ on carbon, δ- on oxygen). However, because the molecule is perfectly linear and symmetrical, these two bond dipoles point in exactly opposite directions and cancel each other out completely. The result is a molecule with polar bonds but no net dipole moment—making it nonpolar overall.
Choice B incorrectly suggests complete electron transfer, which would create ionic bonds, not the covalent bonds actually present in CO₂. Choice C contains a fundamental error: it claims the C=O bonds are nonpolar, contradicting the given information that these bonds are polar covalent. Choice D confuses cause and effect—while CO₂ does dissolve in nonpolar solvents (supporting its nonpolar nature), this solubility behavior doesn't explain why the bonds themselves are polar.
For HESI success, remember this pattern: molecular geometry beats individual bond polarity in determining overall molecular behavior. Always consider both bond polarity and molecular shape—symmetrical molecules often have surprising properties due to dipole cancellation.
- A compound has the chemical formula MgCl₂. What is the primary type of bond holding this compound together, and how are the electrons arranged?
Answer: Ionic; one magnesium atom transfers two electrons total, one to each of two chlorine atoms.
When you encounter chemical formulas like MgCl₂, you need to identify the bonding type by examining the elements involved. Metals bonding with nonmetals typically form ionic bonds, while nonmetals bonding together form covalent bonds.
Magnesium (Mg) is a metal in Group 2, meaning it has 2 valence electrons that it readily loses to achieve a stable electron configuration. Chlorine (Cl) is a nonmetal in Group 17, needing 1 electron to complete its outer shell. In MgCl₂, the magnesium atom loses both of its valence electrons, becoming Mg²⁺, while each chlorine atom gains 1 electron to become Cl⁻. This complete transfer of electrons creates the ionic bond that holds the compound together.
Option A incorrectly describes this as covalent bonding with electron sharing. Covalent bonds involve sharing electrons between atoms, not the complete transfer seen here. Option C is wrong because it describes an impossible scenario—two magnesium atoms cannot form a double bond with one chlorine atom, and this doesn't match the given formula. Option D makes a counting error, suggesting that one magnesium transfers two electrons to just one chlorine atom, but the formula MgCl₂ clearly shows two chlorine atoms are present.
Option B correctly identifies the ionic bonding and the proper electron distribution: one magnesium transfers its two electrons (one to each chlorine), satisfying both the 2:1 ratio in the formula and each atom's need for electron stability.
Remember: when you see metal + nonmetal combinations, think ionic bonding with electron transfer, not sharing.
- A compound has a melting point of 2800°C, conducts electricity when molten but not when solid, and is brittle. However, its constituent elements have an electronegativity difference of 1.5. What is the most likely explanation for these properties?
Answer: The compound is ionic despite the moderate electronegativity difference, with properties dominated by lattice energy
The properties (high melting point, conductivity when molten but not solid, brittleness) are characteristic of ionic compounds. While 1.5 electronegativity difference is moderate, these physical properties definitively indicate ionic bonding. The high melting point suggests high lattice energy from small, highly charged ions. Option A incorrectly identifies it as covalent network. Option C describes an unrealistic bonding scenario. Option D incorrectly suggests metallic bonding, which would show conductivity in solid state.
- Compare the bond formation in LiCl and HCl. Both involve chlorine, but they exhibit vastly different properties. Which explanation best accounts for this difference?
Answer: LiCl involves electron transfer creating Li⁺ and Cl⁻ ions, while HCl involves unequal electron sharing between H and Cl atoms
When you encounter questions about chemical bonding, focus on the fundamental difference between ionic and covalent bonds. The key is understanding how electronegativity differences between atoms determine the type of bond that forms.
LiCl and HCl demonstrate this principle perfectly. Lithium (Li) is a metal with very low electronegativity, while chlorine (Cl) is a highly electronegative nonmetal. This massive electronegativity difference means lithium completely transfers its valence electron to chlorine, creating Li⁺ and Cl⁻ ions held together by electrostatic attraction - an ionic bond. In contrast, HCl forms between two nonmetals with a moderate electronegativity difference. Here, hydrogen and chlorine share electrons unequally, creating a polar covalent bond where chlorine pulls the shared electrons closer, giving it a partial negative charge.
Choice A is incorrect because bond strength isn't determined by the number of bonds an atom can theoretically form, and hydrogen doesn't form multiple bonds with chlorine in HCl. Choice B incorrectly states that both involve electron sharing - LiCl involves complete electron transfer, not sharing. Choice D reverses the bond types entirely: LiCl is ionic (not covalent), and hydrogen doesn't easily lose electrons to form H⁺ ions in HCl.
The correct answer is C because it accurately describes the fundamental bonding difference: electron transfer in LiCl versus unequal electron sharing in HCl.
Remember this pattern: large electronegativity differences (typically metal + nonmetal) usually mean ionic bonding, while moderate differences (nonmetal + nonmetal) typically indicate polar covalent bonding.
- A student predicts that because both CaF₂ and SiF₄ contain fluorine, they should have similar properties. Which analysis best explains why this prediction is incorrect?
Answer: CaF₂ has ionic bonding creating a crystalline lattice, while SiF₄ has covalent bonding creating discrete molecules
CaF₂ is ionic (Ca²⁺ and F⁻ ions) forming an extended crystalline lattice with high melting point and conductivity when molten. SiF₄ is covalent with discrete tetrahedral molecules, existing as a gas at room temperature. The bonding type, not just the presence of fluorine, determines properties. Option B incorrectly suggests similar bonding types. Option C focuses on atomic size rather than fundamental bonding differences. Option D incorrectly describes the bonding in both compounds.
- Methane (CH₄) is a covalently bonded molecule, while sodium chloride (NaCl) is an ionically bonded compound. Which statement is a likely prediction of their relative physical properties?
Answer: At room temperature, methane is a gas, while sodium chloride is a crystalline solid.
When you encounter questions comparing covalent and ionic compounds, focus on how their different bonding types lead to distinct physical properties. Covalent bonds involve shared electrons between atoms, while ionic bonds form through electron transfer, creating charged ions that arrange in crystal lattices.
The correct answer is C because the bonding differences directly affect physical states at room temperature. Methane's covalent bonds create discrete molecules held together by weak intermolecular forces (van der Waals forces). These weak attractions are easily overcome by thermal energy at room temperature, making methane a gas. Sodium chloride's ionic bonds create strong electrostatic attractions between Na⁺ and Cl⁻ ions in a rigid crystal structure, requiring much more energy to disrupt, so it remains a solid.
Option A is incorrect because ionic compounds like NaCl have much higher melting points than small covalent molecules like methane. NaCl melts at 801°C, while methane melts at -182°C.
Option B is wrong because methane doesn't dissociate into ions in water, so it cannot conduct electricity. NaCl dissolves into Na⁺ and Cl⁻ ions, making its solution an excellent electrical conductor.
Option D reverses the solubility relationship. Methane is nonpolar and follows "like dissolves like," making it poorly soluble in polar water. NaCl is highly polar and dissolves readily in water through ion-dipole interactions.
Remember this pattern: ionic compounds typically have high melting points, conduct electricity when dissolved, and dissolve in polar solvents, while small covalent molecules often exist as gases, don't conduct electricity, and prefer nonpolar solvents.
- What is the defining characteristic of a polar covalent bond?
Answer: The unequal sharing of valence electrons between two atoms, creating partial positive and negative charges.
When you encounter questions about chemical bonding, focus on how electrons behave between atoms—whether they're transferred, shared equally, or shared unequally.
A polar covalent bond forms when two atoms with different electronegativities share electrons unequally. The more electronegative atom pulls the shared electrons closer, creating a partial negative charge (δ-) on itself and leaving a partial positive charge (δ+) on the other atom. This unequal distribution creates polarity within the molecule, like in water (H₂O) where oxygen is more electronegative than hydrogen.
Looking at the incorrect options: Choice A describes ionic bonding, where electrons are completely transferred from metals to nonmetals, creating full charges rather than partial ones. Choice B represents nonpolar covalent bonding, which occurs between atoms of identical electronegativity (like in H₂ or Cl₂) where electrons are shared equally. Choice D describes metallic bonding, where metal atoms form a "sea" of delocalized electrons that can move freely throughout the structure.
The key distinction is that polar covalent bonds involve sharing (not transferring) electrons, but the sharing is unequal due to electronegativity differences, creating partial charges rather than complete charges or equal distribution.
Study tip: Remember the bonding spectrum from most to least polar: ionic (complete transfer) → polar covalent (unequal sharing) → nonpolar covalent (equal sharing). When you see "partial charges" or "unequal sharing," think polar covalent. This concept frequently appears on chemistry sections of health science entrance exams.
- Based on their positions in the periodic table, which of the following pairs of elements is most likely to form a covalent bond?
Answer: Sulfur (S) and Oxygen (O)
When you encounter questions about chemical bonding, focus on the types of elements involved and their positions in the periodic table. Covalent bonds form when atoms share electrons, which typically occurs between nonmetals with similar electronegativity values. Ionic bonds form when electrons transfer from metals to nonmetals due to large electronegativity differences.
Looking at the periodic table positions, sulfur and oxygen (choice C) are both nonmetals located on the right side of the periodic table. They have similar electronegativity values, making electron sharing favorable. When these elements bond, they form covalent compounds like sulfur dioxide (SO₂) or sulfur trioxide (SO₃).
Choice A (potassium and fluorine) represents a classic ionic bond. Potassium is an alkali metal that readily loses electrons, while fluorine is a highly electronegative halogen that gains electrons easily. Choice B (calcium and oxygen) also forms an ionic bond, as calcium is an alkaline earth metal and oxygen is a nonmetal with high electronegativity. Choice D (lithium and chlorine) creates another ionic bond between a metal and a halogen.
The key pattern to remember: metal + nonmetal = ionic bond, while nonmetal + nonmetal = covalent bond. On the HESI, these bonding questions often test whether you can quickly identify element types from their periodic table positions. Practice recognizing metals (left side), nonmetals (right side), and metalloids (diagonal line) to quickly determine the most likely bonding type.
- A nurse is preparing a solution of sodium bicarbonate (NaHCO₃). What types of chemical bonds are present within one formula unit of this compound?
Answer: Both ionic and covalent bonds
When analyzing chemical compounds, you need to identify the types of atoms present and how they're connected. Sodium bicarbonate (NaHCO₃) contains both a metal (sodium) and nonmetals (hydrogen, carbon, oxygen), which creates multiple bonding types within the same compound.
The correct answer is C because NaHCO₃ contains both ionic and covalent bonds. The sodium ion (Na⁺) forms ionic bonds with the bicarbonate ion (HCO₃⁻) through electrostatic attraction between oppositely charged ions. Meanwhile, within the bicarbonate ion itself, covalent bonds hold the hydrogen, carbon, and oxygen atoms together through shared electron pairs.
Option A is incorrect because while ionic bonds do exist between Na⁺ and HCO₃⁻, the atoms within the bicarbonate ion are covalently bonded, so it's not "only" ionic bonds. Option B is wrong for the opposite reason - while covalent bonds exist within the bicarbonate ion, the overall compound is held together by ionic attraction between Na⁺ and HCO₃⁻. Option D is incorrect because metallic bonds occur between metal atoms in pure metals or alloys, not in ionic compounds like sodium bicarbonate.
Study tip for the HESI: When you see a compound with both metals and nonmetals, immediately consider that it likely contains both ionic bonds (between the metal and nonmetal groups) and covalent bonds (within polyatomic ions). Memorize common polyatomic ions like bicarbonate, sulfate, and phosphate - they frequently appear in questions testing your understanding of mixed bonding types.
- The bonds between oxygen and hydrogen within a single water molecule (H₂O) are essential for its life-sustaining properties. These bonds are best described as:
Answer: Polar covalent, because oxygen is significantly more electronegative than hydrogen.
When you encounter questions about chemical bonding, focus on electronegativity differences and how electrons are shared or transferred between atoms.
In a water molecule, oxygen and hydrogen atoms share electrons, but not equally. Oxygen has a much higher electronegativity (3.44) than hydrogen (2.20), meaning oxygen pulls the shared electrons closer to itself. This unequal sharing creates a polar covalent bond, where the oxygen end becomes slightly negative (δ-) and the hydrogen ends become slightly positive (δ+). This polarity gives water its unique properties like surface tension and its ability to dissolve many substances.
Let's examine why the other options are incorrect:
A) is wrong because hydrogen doesn't transfer its electron completely to oxygen. In ionic bonding, electrons are fully transferred from one atom to another, typically between metals and nonmetals with very large electronegativity differences.
C) misses the crucial point about electronegativity differences. While both atoms are nonmetals that share electrons, the significant difference in their electronegativity values makes the bond polar, not nonpolar.
D) confuses intramolecular and intermolecular forces. Hydrogen bonds occur between different water molecules, not within a single H₂O molecule. They're the attractions between the positive hydrogen of one molecule and the negative oxygen of another.
For HESI success, remember that bond type depends on electronegativity differences: large differences create ionic bonds, moderate differences create polar covalent bonds, and small differences create nonpolar covalent bonds. Always distinguish between bonds within molecules versus forces between molecules.
- The electronegativity values for four elements are: A = 3.5, B = 3.0, C = 1.5, and D = 0.8. Which pair of elements would form a bond with the most ionic character?
Answer: A and D
When you encounter questions about ionic character in chemical bonds, remember that ionic character depends on the electronegativity difference between two elements. The greater the difference in electronegativity values, the more ionic (and less covalent) the bond becomes.
To find the most ionic bond, calculate the electronegativity difference for each pair. For option A (A and B): 3.5−3.0=0.5. For option B (B and C): 3.0−1.5=1.5. For option C (A and D): 3.5−0.8=2.7. For option D (C and D): 1.5−0.8=0.7. The largest difference is 2.7, making A and D the correct answer.
Option A is incorrect because the small electronegativity difference (0.5) indicates these elements would form a largely covalent bond with minimal ionic character. Option B shows a moderate difference (1.5) that would create a polar covalent bond with some ionic character, but not the most. Option D has an even smaller difference (0.7) than option B, resulting in a bond that's primarily covalent.
Remember this rule: electronegativity differences less than 0.5 are nonpolar covalent, 0.5-1.7 are polar covalent, and greater than 1.7 are considered ionic. On the HESI, always calculate the actual differences rather than estimating visually—this prevents careless errors and ensures you identify the largest gap between electronegativity values. - When predicting the bonding in a series of fluorides (LiF, BeF₂, BF₃, CF₄), which trend correctly describes the progression from ionic to covalent character?
Answer: Ionic character decreases from LiF to CF₄ as the central atom's electronegativity increases and electron sharing becomes favorable
Moving from Li to C, electronegativity increases (Li: 1.0, Be: 1.6, B: 2.0, C: 2.5), making electron sharing more favorable and reducing ionic character. LiF is ionic, BeF₂ has some covalent character, BF₃ is predominantly covalent, and CF₄ is purely covalent. Option A ignores the effect of the central atom's electronegativity. Option B incorrectly describes the relationship between charge density and bonding character. Option D reverses the trend direction.
- Which statement accurately contrasts the role of valence electrons in the formation of an ionic bond versus a covalent bond?
Answer: In ionic bonds, valence electrons are transferred from one atom to another; in covalent bonds, they are shared between atoms.
Understanding chemical bonding requires recognizing how atoms achieve stable electron configurations through their valence electrons—the electrons in their outermost shell that participate in bonding.
In ionic bonding, atoms with very different electronegativities interact. A metal atom (like sodium) will completely transfer one or more valence electrons to a nonmetal atom (like chlorine). This creates charged ions: the metal becomes a positively charged cation, and the nonmetal becomes a negatively charged anion. These oppositely charged ions then attract each other electrostatically.
In covalent bonding, atoms with similar electronegativities share valence electrons. Rather than one atom taking electrons from another, both atoms contribute electrons to form shared electron pairs that occupy the space between the nuclei, holding the atoms together.
Option B correctly describes this fundamental difference—ionic bonds involve electron transfer while covalent bonds involve electron sharing.
Option A incorrectly describes metallic bonding (the "sea" of electrons) rather than ionic bonding. Option C reverses the concepts, suggesting ionic bonds involve sharing when they actually involve complete transfer. Option D is incomplete because it only describes half of each bonding type and doesn't capture the essential difference between transfer and sharing.
Remember this key distinction: if you see "transfer" or "give/take" electrons, think ionic bonding between metals and nonmetals. If you see "share" electrons, think covalent bonding between nonmetals. This fundamental difference drives most of the contrasting properties between ionic and covalent compounds.
- When analyzing the bonding in BeF₂ and MgF₂, both formed from Group 2 metals and fluorine, which statement best explains the difference in their bonding characteristics?
Answer: BeF₂ has significant covalent character due to beryllium's small size and high charge density, while MgF₂ is predominantly ionic
BeF₂ has significant covalent character because Be²⁺ is very small with high charge density, causing polarization of the F⁻ electron cloud (Fajans' rules). MgF₂ is predominantly ionic because Mg²⁺ is larger with lower charge density. Option A ignores the effect of cation size on bonding character. Option C incorrectly suggests larger cations promote covalent bonding. Option D incorrectly assumes formula type determines bonding character regardless of atomic properties.
- A scientist analyzes an unknown liquid. Which of the following findings would most strongly suggest that the substance is a molecular compound held together by covalent bonds?
Answer: The substance has a low boiling point and is a poor electrical conductor.
When you encounter questions about molecular compounds and covalent bonding, focus on the characteristic properties that distinguish them from ionic compounds. Molecular compounds consist of atoms sharing electrons through covalent bonds, creating discrete molecules rather than extended lattice structures.
Choice C correctly identifies key properties of molecular compounds. Low boiling points occur because molecular compounds are held together by relatively weak intermolecular forces (like van der Waals forces) rather than strong ionic or metallic bonds. Poor electrical conductivity is another hallmark—since electrons are localized in covalent bonds between specific atoms, they aren't free to move and carry electric current.
Choice A is incorrect because a boiling point of 1250°C indicates extremely strong bonding, typical of ionic compounds or network covalent solids like diamond, not simple molecular compounds. Choice B describes ionic compounds, which form rigid crystal lattices due to the electrostatic attraction between oppositely charged ions—this creates hard, brittle structures that shatter along cleavage planes. Choice D describes ionic compounds that dissociate in water, releasing mobile ions that conduct electricity well.
The contrast is clear: molecular compounds have weak intermolecular attractions leading to low melting/boiling points and poor conductivity, while ionic compounds have strong electrostatic forces creating high melting points, crystal structures, and good conductivity when dissolved.
Remember this pattern: when identifying molecular compounds, look for properties reflecting weak intermolecular forces—low boiling/melting points, poor electrical conductivity, and often solubility in nonpolar solvents rather than water.
