Metallic Bonding and Properties of Metals Quiz
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5 questions, one answer each
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1. Which of the following best describes metallic bonding?
- A. The electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons
- B. The covalent bonds formed between adjacent metal atoms in the lattice
- C. The ionic bonds between metal atoms and non-metal atoms
- D. The van der Waals forces between individual metal atoms
2. Why are metals good conductors of electricity?
- A. Because metal atoms are tightly packed together in a regular lattice
- B. Because the delocalised electrons are free to move through the structure
- C. Because metal ions are positively charged and repel each other
- D. Because metals contain covalent bonds that allow charge transfer
3. Metals have high melting and boiling points. What is the reason for this?
- A. Metal atoms are very large compared to non-metal atoms
- B. The metallic bonds between the ion lattice and electron sea are very strong
- C. Metal lattices contain a high number of atoms in a small volume
- D. Metallic bonding only occurs in solid metals and breaks down when heated
4. Explain why metals are malleable and ductile, whereas ionic compounds are brittle.
- A. Metals have more electrons than ionic compounds
- B. In metals, layers of ions can slide over each other whilst maintaining metallic bonding; in ionic compounds, shifting layers brings like charges together and causes repulsion
- C. Metals are less dense than ionic compounds, making them easier to deform
- D. The delocalised electrons prevent the formation of strong bonds in metals
5. Why are alloys typically harder than their constituent pure metals?
- A. Alloys contain more electrons in their delocalised electron sea
- B. The different sized atoms in an alloy distort the regular lattice, making it harder for layers of ions to slide over each other
- C. Alloys have stronger metallic bonds than pure metals
- D. Alloys have lower melting points, which makes them harder at room temperature