Topic 4: Chemical Bonding and Structure Flashcards
Define intermolecular and intramolecular bonds.
Intermolecular bonds are between molecules. These are weak. Eg. London dispersion forces, dipole-dipole and hydrogen bonds.
Intramolecular bonds are within atoms or ions. These are strong. Eg. ionic, covalent and metallic.
Define electronegativity.
The ability of an atom to attract the bonding electron pair in a covalent bond to itself.
Describe how electronegativity affects the type of bonding that elements undergo.
Elements with low electronegativity form ionic bonds with atoms of high electronegativity.
Elements with moderate to high electronegativity form covalent bonds with each other.
Describe the 1.7 rule.
Atoms that have electronegativity differences greater than 1.7 form ionic bonds, atoms with electronegativity differences less than 1.7 form polar covalent bonds. The smaller the electronegativity difference the less polar the bond. If the difference is zero the bond is totally covalent.
Define an ionic bond.
The electrostatic attraction between oppositely charged ions. It is formed by the transfer of an electron. It is an example of a strong intramolecular bond.
Describe the physical properties of ionic compounds.
- High melting points/boiling points.
- Only conducts electricity when molten or in solution.
- Strong electrostatic forces.
- Brittle.
- Soluble in water.
Describe an ionic lattice.
Ions in a lattice are arranged in a regular repeating pattern so that positive charges cancel out negative charges. Therefore the final lattice is overall electrically neutral.
Define metallic bonding.
The attraction between the orderly arrangement of positive metal ions and the delocalised valence electrons. It is an example of intramolecular bonding.
Describe the physical properties of metallic compounds.
- Good electrical conductors.
- Can conduct heat.
- Ductile.
- Malleable.
Define covalent bonding.
A shared pair of electrons with one electron being supplied by each atom either side. It is an example of intramolecular bonding. NON-METAL ATOMS. There are two types:
1) Discrete covalent molecules
2) Giant covalent networks (lattice structure in 3D)
Define dative covalent bonding.
Both bonding electrons are supplied by one atom.
Describe the physical properties of discrete covalent molecules.
- Low melting point/boiling point.
- Poor conductor of electricity as no mobile electrons.
- Insoluble.
When can lewis diagrams be drawn? And give the rules for drawing.
Lewis diagrams are drawn to represent discrete covalent molecules.
Rules:
- Count the total number of electrons.
- Identify the central atom from symmetry or it being the least electronegative atom.
- Draw bonding electrons and non bonding electrons until all are used up.
*Square brackets with charge for polyatomic ions.
Describe the differences between bond lengths and strengths.
The greater the forces of attraction between electrons and nuclei, the more the atoms are pulled closer to each other. This decreases the bond length of a molecule and increases the strength of the covalent bond.
Triple bonds are the shortest and strongest covalent bonds due to the large electron density between the nuclei of the two atoms.
Give a summary of the shapes of molecules with 2 regions of electron density around the central atom.
For 2 bonding pairs, 0 non bonding pairs the shape is LINEAR. The bond angle is 180°. Eg. CO2.
