2.2.2: Bonding & Structure Flashcards
Ionic bond
Electrostatic attraction between two oppositely charged ions.
Groups whose elements form ionic compounds
1
2
6
7
Properties of giant ionic structures
- High melting point
- Electrical insulators (except when molten or in solution)
- Brittle
- Soluble in polar molecules
Explain why ionic compounds have high melting points
- Must overcome the (almost infinite) electrostatic attractions between ions
- Huge numbers of ions ∴ large amounts of energy needed
Explain why ionic compounds are hard and brittle
- Ions locked in fixed position, will not move
- Electrostatic attractions are strong
- Will break if enough energy supplied –> structure becomes misaligned
- Formation of breakage at a discontinuity
What determines the way in which the ions are arranged in a lattice?
- Ratio of ions
- Size of ions
- Charge of ions
Effect of charge density on energy required to melt ionic compound
Charge density = charge/area
Ions with increased charge density attract more strongly∴more energy required to break electrostatic attraction and melt
Explain why ionic compounds dissolve in polar compounds e.g. water
- Solubility depends on overcoming electrostatic energy and replacing it with similarly strong interaction with water
- Need to match energy held in electrostatic attraction with energy given out making new bonds
- If energy out greater than energy required to break ionic bonds then substance will dissolve
Covalent bond
A shared pair of electrons. Occurs (generally) between non-metals.
Types of covalent structures:
- Giant covalent
* Simple molecular (finite)
Examples of giant covalent structures
Silicon dioxide, graphite, diamond
Examples of simple molecular structures
Carbon dioxide, water
Dative covalent bond
One of the atoms forming the bond donates both the bonding electrons; coordinate bond.
All single bonds are _____ bonds
sigma bonds
The sharing of electrons between atoms is achieved by:
the overlap of atomic orbitals containing 1 electron each.
Properties of simple molecular structures
- Low MBP
- Do not conduct electricity
- Soluble in organic, non polar solvents
Which is stronger, π bonds orσ bonds?
• σ bonds are stronger
Double bonds are made up of:
1 σ and 1 π bond
Explain why simple molecular structures have low MBP
• Weak London forces require little energy to break
Explain why simple molecular structures do not conduct electricity
• No free ions/electrons to carry charge
Explain why simple molecular structures are soluble in organic solvents
- Not polar so not soluble in polar solvents
* Soluble instead in organic solvents which are non-polar (“like dissolves like”)
Properties of Diamond
- Very high MBP
- Very hard
- High density
- Poor thermal/electrical conductivity because electrons held rigidly in very strong covalent bonds
Different arrangements of the same element e.g. graphite and diamond
Allotropes
Properties of Graphite
- High MBP because strong covalent bonds between carbon atoms require much energy to break
- Conducts electricity because delocalised electron between layers can carry charge across molecule
Shapes of covalent molecules determined by:
Valence Shell Electron Pair Repulsion theory
Order of repulsion
- Lone pair to lone pair = Most repulsion
- Lone pair to bonding pair
- Bonding pair to bonding pair = least repulsion
Intermolecular forces
Short range forces that exist between (simple covalent) molecules. They are overcome during physical changes (boiling/melting)
Intermolecular forces arise from what?
Differential movement of electrons in a molecule relative to nuclei which gives rise to temporary dipoles.
Types of intermolecular forces
- Permanent dipole-dipole attractions
- London forces: temporary instantaneous mutually induced dipole-dipole attraction
- Hydrogen bonding
All types of intermolecular forces are comparatively _______ than a typical covalent bond.
weaker
Bond angle in tetrahedral
109.5°
Bond angle in pyramidal
107.5°
Bond angle in linear
180°
Bond angle in angular
104.5°
Bond angle in octahedral
90°
Bond angle in trigonal planar
120°
Bond angle in square planar
180° / 90° depending on which bonds
Bond angle in trigonal bypyramidal
120° / 90° depending on which bonds
Shapes of molecules options
- Linear
- Angular
- Pyramidal
- Tetrahedral
- Octahedral
- Trigonal planar
- Square planar
- Trigonal bypyramidal
Sigma bonds can arise from:
• Overlap of atomic orbitals containing 1 electron each
• End of overlap of p orbitals
Pi bonds arise from
• Sideways overlap of p orbitals above and below the plane of the sigma bond
Factors affecting the strength of intermolecular forces (London forces)
- Number of e- in a molecule
* SA of the molecules (larger SA = more IMF)
