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Questions and Answers
In covalent molecular bonding, what type of forces are present between molecules?
In covalent molecular bonding, what type of forces are present between molecules?
- Weak intramolecular forces
- Only strong covalent bonds
- Strong intermolecular forces
- Strong intramolecular covalent bonds, but weak intermolecular forces (correct)
Ionic bonding occurs when two nonmetals share electrons to achieve a full outer shell.
Ionic bonding occurs when two nonmetals share electrons to achieve a full outer shell.
False (B)
What distinguishes covalent network bonding from covalent molecular bonding in terms of structure?
What distinguishes covalent network bonding from covalent molecular bonding in terms of structure?
Covalent network bonding forms large, continuous 3D structures, while covalent molecular bonding forms small, distinct molecules.
In metallic bonding, metal atoms lose valence electrons, creating a 'sea of ________ electrons' that move freely among positive metal cations.
In metallic bonding, metal atoms lose valence electrons, creating a 'sea of ________ electrons' that move freely among positive metal cations.
Match the following bond types with the primary force that holds them together:
Match the following bond types with the primary force that holds them together:
Which type of intermolecular force is the strongest?
Which type of intermolecular force is the strongest?
Metals tend to have high electronegativity, which is why they readily lose electrons in metallic bonding.
Metals tend to have high electronegativity, which is why they readily lose electrons in metallic bonding.
Explain why water (H₂O) exhibits hydrogen bonding, while methane (CH₄) only exhibits dispersion forces.
Explain why water (H₂O) exhibits hydrogen bonding, while methane (CH₄) only exhibits dispersion forces.
Dispersion forces occur due to the formation of temporary ________ as a result of random electron movement.
Dispersion forces occur due to the formation of temporary ________ as a result of random electron movement.
Which of the following compounds is most likely to form a covalent network structure?
Which of the following compounds is most likely to form a covalent network structure?
In dipole-dipole forces, molecules align so that the partially positive end of one molecule is near the partially positive end of another molecule.
In dipole-dipole forces, molecules align so that the partially positive end of one molecule is near the partially positive end of another molecule.
Describe the role of electronegativity in ionic bond formation.
Describe the role of electronegativity in ionic bond formation.
In sodium chloride (NaCl), sodium (Na) loses one electron to form a ________, and chlorine (Cl) gains one electron to form an ________.
In sodium chloride (NaCl), sodium (Na) loses one electron to form a ________, and chlorine (Cl) gains one electron to form an ________.
Which type of bonding involves a 'sea of delocalized electrons'?
Which type of bonding involves a 'sea of delocalized electrons'?
Covalent bonds are generally weaker than intermolecular forces.
Covalent bonds are generally weaker than intermolecular forces.
Explain why diamond is so much harder than graphite, even though both are made of carbon atoms.
Explain why diamond is so much harder than graphite, even though both are made of carbon atoms.
Hydrogen bonding occurs when hydrogen is covalently bonded to _, _, or _.
Hydrogen bonding occurs when hydrogen is covalently bonded to _, _, or _.
Which of these molecules would exhibit dipole-dipole forces?
Which of these molecules would exhibit dipole-dipole forces?
Dispersion forces are stronger in small molecules with fewer electrons.
Dispersion forces are stronger in small molecules with fewer electrons.
Describe the difference between intramolecular and intermolecular forces. Give an example of each.
Describe the difference between intramolecular and intermolecular forces. Give an example of each.
Flashcards
Covalent Bonding
Covalent Bonding
Sharing of electrons between nonmetals to achieve a full outer shell.
Covalent Molecular Bonding
Covalent Molecular Bonding
Small, distinct molecules held together by covalent bonds with weak intermolecular forces.
Covalent Network Bonding
Covalent Network Bonding
Large, continuous 3D structure held by strong covalent bonds throughout.
Ionic Bonding
Ionic Bonding
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Metallic Bonding
Metallic Bonding
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Intermolecular Forces
Intermolecular Forces
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Dispersion Forces
Dispersion Forces
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Dipole-Dipole Forces
Dipole-Dipole Forces
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Hydrogen Bonding
Hydrogen Bonding
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Cation
Cation
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Anion
Anion
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Electronegativity
Electronegativity
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Delocalized Electrons
Delocalized Electrons
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Electrostatic Force
Electrostatic Force
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Octet Rule
Octet Rule
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Study Notes
- Bonding arises from interactions between atoms, driven by electrons, electrostatic forces, and the pursuit of energy stability.
Covalent Bonding
- Covalent bonding involves the sharing of electrons between two or more nonmetal atoms to achieve a full outer shell (octet rule).
- Nonmetals exhibit high electronegativity and do not readily lose electrons, making sharing favorable.
- Example: In H₂O, each hydrogen atom shares one electron with the oxygen atom, resulting in oxygen having eight valence electrons and each hydrogen having two.
- Covalent bonding results in either covalent molecular or covalent network structures.
Covalent Molecular Bonding
- In covalent molecular bonding, small, distinct molecules form, held together by covalent bonds.
- Example: In CO₂, carbon shares two pairs of electrons (double bond) with each oxygen atom.
- Strong intramolecular covalent bonds are present, but intermolecular forces are weak.
Covalent Network Bonding
- Covalent network bonding forms large, continuous 3D structures with strong covalent bonds throughout.
- Example: Diamond (C), where each carbon bonds with four others.
- Only strong covalent bonds are present, without weak intermolecular forces.
Ionic Bonding
- Ionic bonding occurs when a metal transfers electrons to a nonmetal.
- This transfer forms positive (cation) and negative (anion) ions.
- Metals have low electronegativity, readily losing electrons, while nonmetals have high electronegativity, easily gaining electrons.
- Example: In NaCl, sodium (Na) loses one electron to become Na⁺, and chlorine (Cl) gains one electron to become Cl⁻.
- The oppositely charged ions attract each other, establishing a strong electrostatic force (ionic bond).
- Strong electrostatic forces exist between the ions in a lattice structure.
Metallic Bonding
- Metallic bonding happens when metal atoms lose valence electrons.
- This loss creates a "sea of delocalized electrons" moving freely among positive metal cations.
- Metals readily lose electrons, dictated by their low electronegativity.
- Example: Copper (Cu), where copper atoms release valence electrons to form Cu²⁺ cations.
- Free-moving electrons create a strong attraction between the metal cations.
- Strong electrostatic attraction can be observed between cations and delocalized electrons.
Intermolecular Forces (Between Covalent Molecules)
- Intermolecular forces are weaker than bonds but influence melting/boiling points and solubility.
Dispersion Forces
- Dispersion forces are the weakest intermolecular force.
- Temporary dipoles form due to random electron movement.
- Electrons are constantly moving, leading to temporary charge imbalances.
- Example: In O₂, a temporary negative region in one molecule attracts the positive region of another.
Dipole-Dipole Forces
- Dipole-dipole forces arise when polar molecules align due to permanent partial charges.
- Unequal electron sharing in polar covalent bonds creates δ⁺ and δ⁻ ends.
- Example: In HCl, the δ⁺ hydrogen of one molecule attracts the δ⁻ chlorine of another.
Hydrogen Bonding
- Hydrogen bonding is the strongest intermolecular force.
- It occurs when hydrogen is covalently bonded to N, O, or F creating a strong dipole-dipole attraction.
- Nitrogen, oxygen, and fluorine are highly electronegative, drawing electrons away from hydrogen, making it highly δ⁺.
- Example: In H₂O, the δ⁺ hydrogen in one molecule attracts the δ⁻ oxygen in another.
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