Questions on Metallic Bond
20 Multiple-Choice Questions: Metallic Bond
Questions
1. What type of particles are involved in a metallic bond?
A) Protons and neutrons
B) Positive ions and delocalized electrons
C) Neutrons and electrons
D) Atoms and protons
E) Covalent molecules
2. Which of the following best describes a metallic bond?
A) Transfer of electrons from one atom to another
B) Sharing of electrons between two atoms
C) Electrostatic attraction between metal cations and free electrons
D) Attraction between opposite ends of polar molecules
E) Hydrogen bonding between atoms
3. What gives metals their ability to conduct electricity?
A) Static lattice structure
B) Free movement of neutrons
C) Presence of delocalized electrons
D) Transfer of protons
E) Rigid covalent structure
4. Which property of metals is most directly related to metallic bonding?
A) Solubility in water
B) High ionization energy
C) Malleability
D) Lack of luster
E) Poor thermal conductivity
5. What term describes the mobile electrons in a metallic bond?
A) Stationary electrons
B) Valence protons
C) Delocalized electrons
D) Bonding ions
E) Nuclear electrons
6. Metals are good conductors of heat because:
A) They have low melting points
B) Their atoms are highly electronegative
C) Delocalized electrons can transfer energy quickly
D) They dissolve in acids
E) Their nuclei are mobile
7. Why are metals malleable and ductile?
A) Ionic bonds resist compression
B) Covalent bonds form layers
C) Metal atoms can slide past each other without breaking bonds
D) Electrons are tightly held
E) Metallic bonds are directional
8. What holds metal atoms together in a metallic bond?
A) Electron-electron repulsion
B) Nucleus-nucleus attraction
C) Attraction between cations and sea of electrons
D) Hydrogen bonding
E) Ionic attraction between atoms
9. Which of the following is a key characteristic of metallic substances?
A) Low density
B) Soft texture
C) High melting point
D) Poor electrical conductivity
E) Transparent appearance
10. In metallic bonding, the electrons are:
A) Shared between two specific atoms
B) Confined to one atom
C) Shared among all atoms
D) Found only in outermost shells
E) Found in molecules only
11. Which of the following best explains the luster of metals?
A) Metallic bonds absorb visible light
B) Delocalized electrons reflect light
C) Strong ionic bonds
D) Presence of covalent compounds
E) High specific heat
12. What happens to the strength of metallic bonds as the number of delocalized electrons increases?
A) It decreases
B) It remains the same
C) It increases
D) It becomes zero
E) It causes bond repulsion
13. Which of the following is NOT a typical property of metals explained by metallic bonding?
A) Electrical conductivity
B) Thermal conductivity
C) Brittleness
D) Malleability
E) Luster
14. What type of elements form metallic bonds?
A) Nonmetals only
B) Metalloids
C) Noble gases
D) Metals only
E) Transition elements only
15. Which metal has the strongest metallic bond?
A) Sodium
B) Potassium
C) Aluminum
D) Cesium
E) Lithium
16. What structure do metals typically form in solid state?
A) Amorphous
B) Tetrahedral molecules
C) Crystal lattice
D) Covalent chain
E) Non-crystalline solid
17. What explains the high melting points of most metals?
A) Weak intermolecular forces
B) Strong attraction between ions and delocalized electrons
C) Presence of hydrogen bonding
D) Electrostatic repulsion
E) Low electron density
18. Which physical state are metals in under standard conditions (except mercury)?
A) Gas
B) Plasma
C) Solid
D) Liquid
E) Solution
19. What is the “sea of electrons” in metallic bonding?
A) A term for electron orbitals
B) Free electrons that surround and move between metal ions
C) Water molecules around metals
D) Ionized protons
E) Electrons bound to individual atoms
20. Which of the following metals is liquid at room temperature and still exhibits metallic bonding?
A) Mercury
B) Gold
C) Silver
D) Lead
E) Copper
Answers with Explanations
1. B) Positive ions and delocalized electrons
→ Metallic bonds involve a lattice of positive metal ions and a "sea" of mobile electrons.
2. C) Electrostatic attraction between metal cations and free electrons
→ This attraction defines metallic bonding.
3. C) Presence of delocalized electrons
→ These electrons carry charge through the metal.
4. C) Malleability
→ Because atoms can slide over each other in a metallic bond without breaking.
5. C) Delocalized electrons
→ Not tied to one atom, they roam throughout the lattice.
6. C) Delocalized electrons can transfer energy quickly
→ This enables efficient heat conduction.
7. C) Metal atoms can slide past each other without breaking bonds
→ The non-directional nature of metallic bonds allows flexibility.
8. C) Attraction between cations and sea of electrons
→ The delocalized electrons bind the structure.
9. C) High melting point
→ Most metals melt at high temperatures due to strong bonding.
10. C) Shared among all atoms
→ Electrons move freely throughout the structure.
11. B) Delocalized electrons reflect light
→ The surface electrons reflect photons, giving luster.
12. C) It increases
→ More delocalized electrons strengthen the bonding.
13. C) Brittleness
→ Metals are usually not brittle—this is a property of ionic or covalent solids.
14. D) Metals only
→ Only metals exhibit metallic bonding.
15. C) Aluminum
→ Aluminum has three valence electrons contributing to bonding, increasing strength.
16. C) Crystal lattice
→ Metals form organized arrays of atoms.
17. B) Strong attraction between ions and delocalized electrons
→ These bonds require high energy to break.
18. C) Solid
→ All metals except mercury are solid at room temperature.
19. B) Free electrons that surround and move between metal ions
→ This is what keeps the metallic lattice together.
20. A) Mercury
→ The only metal that is liquid at room temperature.
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