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Flashcards in Bonding and Structure Deck (17):
1

How are ions formed?

By electron loss or gain.

2

What are the charges of:

Metals in group 1, 2 and 3: +, 2+ and 3+
Non-metals in group 5, 6 and 7: 3-, 2-, -
Silver: Ag+
Copper: Cu2+
Iron (II): Fe2+
Iron (III): Fe3+
Lead: Pb2+
Zinc: Zn2+
Hydrogen: H+
Hydroxide: OH-
Ammonium: NH4+
Carbonate: CO32-
Nitrate: NO3-
Sulphate: SO42-

3

What is the ‘glue’ in ionic bonding?

The electrostatic forces of attraction between the positively charged cations and the negatively charged anions.

4

Why do compounds with giant ionic lattices have high melting and boiling points?

The electrostatic forces of attraction between the positive actions and negative anions are very strong, with each ion being attracted to each ion of the other charge, pulling them together in a tight, solid lattice. To break down these bonds, a lot of heat energy is needed.

5

Why don’t ionic compounds conduct electricity when in solid state, but do when molten or in aqueous solution?

In solid ionic structures, the positively charged and negatively charged ions are held firmly in place in the lattice, so they cannot move and carry the current. If we melt them or dissolve them, the lattice breaks down and the ions can move and carry a current. Here, it is not electrons but ions carrying the charge.

6

What is a covalent bond formed by?

The sharing of a pair of electrons.

7

What is the ‘glue’ in covalent bonds?

The electrostatic forces of attraction between the shared electrons and the positively charged nuclei of the atoms.

8

Why are substances with a simple molecular structure gases or liquids, or solids with low melting and boiling points?

Although the intramolecular forces (forces within the molecules) are very strong, the London forces (intermolecular forces) are extremely weak, and only require small amounts of heat energy to be broken down.

9

Why are substances with giant covalent structures solids with high melting and boiling points?

To melt giant covalent structures, a huge number of strong covalent bonds have to be broken. This requires a lot of energy. The very strong covalent bonds are also responsible for giant covalent structures’ strength.

10

Why doesn’t a diamond conduct electricity?

Each atom is held together with strong covalent bonds. These atoms can therefore not move (they are fixed) as well as being neutral. They can therefore not conduct.

11

Why does graphite conduct electricity?

The delocalised electrons between the layers of graphite can move and carry current.

12

Why are Buckminster-Fullerenes poor conductors of electricity?

Each carbon in a C60 fullerene has one delocalised electron, however, these electrons can’t move between the molecules.

13

Why are Buckminster-Fullerenes soft?

C60 molecules are only held together by intermolecular forces, so they can slide over each other.

14

Why do covalent compounds not usually conduct electricity?

In covalent compounds, there tend to be no charged particles to carry the current, and the atoms tend to be in a fixed position (can’t move).

15

What is the ‘glue’ in metallic bonding?

The electrostatic forces of attraction between the negatively-charged delocalised electrons and the positive metal cations.

16

What is the structure of a metal?

A giant lattice of positive ions in a ‘sea’ of delocalised electrons.

17

List the typical physical properties of metals and why they have these properties.

- Good conductors of electricity (the delocalised electrons can move and carry current as they are charged particles).
- Malleable and ductile (they consist of layers of atoms which can easily slide over each other).
- High melting and boiling points, tough (the electrostatic forces of attraction between the metal actions and the negatively charged electrons are very strong, and therefore difficult to break, requiring a lot of heat energy).