Thermochemistry Properties Flashcards
(46 cards)
Bent (4 ED)
Tetrahedral arrangement:
- 4 ED
- 2 B
- 2 NB
- 109.5°
Trigonal Pyramid
Tetrahedral arrangement:
- 4 ED
- 3 B
- 1 NB
- 109.5°
Tetrahedral
Tetrahedral arrangement:
- 4 ED
- 4 B
- 0 NB
- 109.5°
Bent (3 ED)
Trigonal Planar arrangement:
- 3 ED
- 2 B
- 1 NB
- 120°
Linear
Linear arrangement:
- 2 ED
- 2 B
- 0 NB
- 180°
Trigonal Planar
Trigonal Planar arrangement:
- 3 ED
- 3 B
- 0 NB
- 120°
Molecular structure
Type of solid: Molecular
Particle: Molecules
Force: Weak intermolecular forces
Strength: Weak
Molecular substances are made up of discrete molecules held together with weak intermolecular forces. These molecules are not charged.
Metal structure
Type of solid: Metal
Particle: Atoms
Force: Metallic bonds
Strength: Strong
Metals are made up of a 3D lattice with strong non-directional electrostatic attractions between the positively charged metal nuclei and the sea of delocalised valence electrons called metallic bonding.
Ionic structure
Type of solid: Ionic
Particle: Ions - cations and anions
Force: Ionic bonds
Strength: Strong
Ionic substances are made up of a 3D lattice with alternating positive and negative ions held together with strong directional ionic bonds. The ions are fixed in place.
Diamond structure
Type of solid: Covalent network
Particle: Atoms
Force: Covalent bonds
Strength: Strong
Diamond is a large 3-D network of atoms bonded together by strong covalent bonds. Each carbon atom is covalently bonded to 4 other carbon atoms forming a strong 3-D covalent network lattice with no free electrons.
Silica structure
Type of solid: Covalent network
Particle: Atoms
Force: Covalent bonds
Strength: Strong
Silica is a large 3-D network of atoms bonded together by strong covalent bonds. Each silicon atom is covalently bonded to 4 oxygen atoms and each oxygen atom is covalently bonded to two silicon atoms in a 3-D covalent network lattice with no free electrons.
Graphite structure
Type of solid: Covalent network
Particle: Atoms
Force: Covalent bonds
Strength: Strong
Graphite is a large 3-D network of atoms bonded together by strong covalent bonds. Each C atom in graphite is covalently bonded to 3 other carbon atoms resulting in hexagonal rings which are arranged in layers. Each C atom, therefore, has a delocalised valence electron. There are weak inter-layer bonds formed between the delocalised electrons of two consecutive layers.
Molecular solids properties
> Hardness:
Molecular solids are softer than ionic, metallic, and covalent network solids due to weak intermolecular forces.
> Forces:
London dispersion forces (weakest), dipole-dipole interactions, and hydrogen bonds hold molecules together.
> Brittleness:
- Molecules held by weak forces (e.g., van der Waals, dipole-dipole) make molecular solids brittle.
- Lack of delocalized electrons prevents easy rearrangement, so molecular solids break under stress.
> Density:
- Molecular solids have lower densities due to weaker intermolecular forces that prevent tight packing of molecules.
> Melting & Boiling Points:
- Molecular solids have low melting and boiling points because of weak intermolecular forces.
- Larger molecules with stronger forces (e.g., hydrogen bonding) have higher melting and boiling points.
> Conductivity:
- Molecular solids are poor conductors. - Electrons are localized, preventing electrical and thermal conductivity in molecular solids.
> Solubility:
- Polar molecular solids dissolve in water due to dipole-dipole interactions; non-polar solids generally do not.
- Substances with hydrogen bonds (e.g., alcohols) dissolve well in water.
Ionic solids properties
> Hardness:
- Ionic solids are generally harder than molecular solids due to strong electrostatic attraction between oppositely charged ions.
- Ionic bonds require significant energy to break, contributing to their hardness.
- Softer than covalent network solids (e.g., diamond).
> Brittleness:
- Rigid lattice structure of alternating positive and negative ions.
- Stress shifts ions, aligning like charges, leading to repulsion and cracks (brittle, not malleable).
> Density:
- High density due to close packing in the lattice and the mass of the ions.
- Denser than molecular solids, but comparable to metallic and covalent network solids.
> Melting/Boiling Points:
- High due to strong ionic bonds, requiring significant energy to break.
- Higher than molecular solids, comparable to metals and covalent network solids.
> Electrical Conductivity:
- Non-conductive in solid form (ions fixed in place).
- Conductive when molten or dissolved (ions free to move).
> Solubility:
Soluble in water! water molecules surround and separate ions (hydration).
Metallic solids properties
> Hardness:
- Metals are generally hard due to strong metallic bonds, but hardness varies depending on the specific metal and its atomic arrangement.
- Unlike brittle ionic and covalent solids, metallic bonds allow ions to move without breaking.
> Malleability and Ductility:
Metals are malleable (can be hammered) and ductile (can be drawn into wires) because delocalised electrons enable metal atoms to slide past each other.
> Density:
Metallic solids typically have high densities due to the closely packed lattice structure of metal atoms.
> Melting and Boiling Points:
Metals have high melting and boiling points due to the strong attraction between metal ions and delocalised electrons, requiring large amounts of energy to break.
> Conductivity:
Metals are excellent conductors of electricity and heat because of the free movement of delocalised electrons.
> Solubility:
Most metals are not soluble in water because the strong metallic bonds cannot be overcome by water’s polar molecules.
Covalent network: diamond properties
> Hardness:
The rigid, continuous network of covalent bonds throughout the structure makes diamond extremely hard, as breaking these bonds requires a lot of energy.
> Melting Point:
The strong covalent bonds between carbon atoms throughout the lattice result in an extremely high melting point, as significant energy is needed to break all the bonds.
> Conductivity:
Diamond does not conduct electricity because all electrons are involved in covalent bonding, leaving no free electrons to carry charge.
> Brittleness:
Despite its hardness, diamond is brittle; applying force breaks the covalent bonds rather than allowing the structure to deform.
> Insolubility:
Diamond is insoluble in water because the strong covalent bonds cannot be broken by water molecules, preventing it from dissolving.
Covalent network: silica properties
> Hardness:
Silicon is hard because of its strong covalent bonds, although not as hard as diamond due to differences in atomic structure and bond strength.
> Melting Point:
Silicon has a very high melting point because breaking the covalent bonds throughout the lattice requires a large amount of energy.
> Conductivity:
- Silicon is a poor conductor of electricity at room temperature because, like diamond, it lacks free electrons.
- However, when heated, some electrons gain enough energy to move, making it a semiconductor.
> Brittleness:
Silicon is brittle as its rigid structure causes the covalent bonds to break under stress, rather than allowing deformation.
> Insolubility:
Silicon is insoluble in water, as the covalent bonds in the lattice structure cannot be broken by interactions with water molecules.
Covalent network: graphite properties
> Hardness:
Graphite is soft because the weak Van der Waals forces between layers allow them to slide over each other easily, making it slippery and soft compared to diamond.
> Melting Point:
Strong covalent bonds within the layers require a high amount of energy to break, resulting in a high melting point.
> Conductivity:
Graphite conducts electricity because each carbon atom has one free electron that can move between layers, allowing for the flow of charge.
> Brittleness:
Graphite is brittle when force is applied perpendicular to the layers, as this disrupts the weak Van der Waals forces between the layers.
> Insolubility:
Graphite is insoluble in water because water molecules cannot break the strong covalent bonds within the layers.
M
= molar mass (m/n)
n
= number of moles (m/M)
m
=mass (nxM)
One mole
=6.022 × 10²³
Comparing Polarities
There are _____ ______ _____ bonds because the _____ atom is more electronegative than the ______ atom. Due to the _______ shape, the ________ bonds are ________ arranged around the central _____ atom. Therefore the bond dipoles ______ cancel each other out and thus _____ is a __________ molecule.
Comparing Molecules
There are ____ areas of electron density around the central ___ atom. The electrons repel each other with maximum separation and minimal repulsion, forming a _______ arrangement. There are ______ areas of bonded electrons and ____ lone pairs making a ______ shape. This means a bond angle of ________°.
