Mock Revision Flashcards
Giant ionic lattice
• An ionic bond is an electrostatic force of attraction between a positively charged metal (cation) ion and a negatively charged non-metal (anion) ion
• The metal becomes positively charged as it transfers electrons to the nonmetal which then becomes negatively charged
• When an ionic compound is formed, the attraction between the ions happens in all directions
• The type of lattice formed depends on the sizes of the positive and negative ions which are arranged in an alternating fashion
Covalent lattice
Covalent bonds are bonds between nonmetals in which electrons are shared between the atoms
• Covalent compounds can be arranged in simple molecular or giant molecular lattices
• Simple molecular lattices: iodine, buckminsterfullerene (Co) and ice
• Giant molecular: silicon(IV) oxide, graphite and diamond
Metal attic structure
Metals form giant metallic lattices in which the metal ions are surrounded by a
‘sea’ of delocalised electrons
• The metalions are often packed in hexagonal layers or in a cubic arrangement
• This layered structure with the delocalised electrons gives a metalits key properties
Alloy + strength of metallic attraction
• If other atoms are added to the metal structure, such as carbon atoms, this creates an alloy
• Alloys are much stronger than pure metals, because the other atoms stop the layers of metalions sliding over each other easily
The strength of the metallic attraction can be increased by:
• Increasing the number of delocalised electrons per metal atom
• Increasing the positive charges on the metal centres in the lattice
• Decreasing the size of the metalions
Malleability
Metallic compounds are malleable
• When a force is applied, the metal layers can slide
• The attractive forces between the metalions and electrons act in all directions
• So when the layers slide, the metallic bonds are re-formed
• The lattice is not broken and has changed shape
Strength + electrica conductivity
• Metallic compounds are strong and hard
• Due to the strong attractive forces between the metalions and delocalised electrons
Metals can conduct electricity when in the solid or liquid state
• In the solid and liquid states, there are mobile electrons which can freely move around and conduct electricity
• When a potential difference is applied to a metallic lattice, the delocalised electrons repel away from the negative terminal and move towards the positive terminall
• As the number of outer electrons increases across a period, the number of delocalised charges also increases:
Thermal conductive
Metals are good thermal conductors due to the behaviour of their cations and their delocalised electrons
• When metals are heated, the cations in the metal lattice vibrate more vigorously as their thermal energy increases
• These vibrating cations transfer their kinetic energy as they collide with neighbouring cations, effectively conducting heat
The delocalised electrons are not bound to any specific atom within the metal lattice and are free to move throughout the material
• When the cations vibrate, they transfer kinetic energy to the electrons
• The delocalised electrons then carry this increased kinetic energy and transfer it rapidly throughout the metal, contributing to its high thermal conductivity.
Meeting and boiling point
Metals have high melting and boiling points
• This is due to the strong electrostatic forces of attraction between the cations and delocalised electrons in the metallic lattice
• These require large amounts of energy to overcome
• As the number of mobile charges increases across a period, the melting and boiling points increase due to stronger electrostatic forces
Effects or banding on physical properties ionic
lonic compounds are strong
• The strong electrostatic forces in ionic compounds keep the ions strongly together
• They are brittle (meaning ionic crystals split apart easily)
• lonic compounds have high melting and boiling points
• The strong electrostatic forces between the ions in the lattice act in all directions and keep them strongly together
• Melting and boiling points increase with charge density of the ions due to the greater electrostatic attraction of charges
• lonic compounds are soluble in water as they can form ion - dipole bonds
• lonic compounds only conduct electricity when molten or in solution
• When molten or in solution, the ions can freely move around and conduct electricity
• In the solid state they’re in a fixed position and unable to move around
Effects os bonding on physical properties metal
Metallic structures are malleable
• When a force is applied, the metal layers can slide
• The attractive forces between the metalions and electrons act in all directions
• So when the layers slide, the metallic bonds are re-formed
• The lattice is not broken and has changed shape
• Metallic lattices are strong and hard
• Due to the strong attractive forces between the metalions and delocalised electrons
• Metals have high melting and boiling points
• Pure metals are insoluble in water
• Metals can conduct electricity when in the solid or liquid state
• As both in the solid and liquid state there are mobile electrons which can freely move around and conduct electricity
Covalent lattice structure
• Simple covalent lattices have low melting and boiling points
• These compounds have weak intermolecular forces between the molecules
• Only little energy is required to break the lattice
• Most compounds are insoluble with water
• Unless they are polar and can form hydrogen bonds (such as sucrose)
• They do not conduct electricity in the solid or liquid state as there are no charged particles
• Some simple covalent compounds do conduct electricity in solution, but this is a reaction with the water than produces ions such as HCl which forms H* and Clions
Giant covenant structure
• Giant covalent lattices have very high melting and boiling points
• These compounds have a large number of covalent bonds linking the whole structure
• Alot of energy is required to break the lattice
• The compounds can be hard or soft
• Graphite is soft as the forces between the carbon layers are weak
• Diamond and silicon(IV) oxide are hard as it is difficult to break their 3D network of strong covalent bonds
• Most compounds are insoluble with water
• Most compounds do not conduct electricity however some do
• Graphite has delocalised electrons between the carbon layers which can move along the layers when a voltage is applied
• Diamond and silicon(IV) oxide do not conduct electricity as all four outer electrons on every carbon atom are involved in a covalent bond so there are no freely moving electrons available
The valence still electron repulsion theory
• When determining the shape and bond angles of a molecule, the following VSEPR rules should be considered:
• Valence shell electrons are those electrons that are found in the outer shell
• Electron pairs repel each other as they have the same charge
• Lone pair electrons repel each other more than bonded pairs
• Repulsion between multiple and single bonds is treated the same as for repulsion between single bonds
• Repulsion between pairs of double bonds are greater
• The most stable shape is adopted to minimize the repulsion forces
• The order of repulsion is therefore: lone pair - lone pair > lone pair - bond pair > bond pair - bond pair
Shapes of molecules
2 bp no ep = linear / 180•
3 bp no ep = trigonal planar / 120•
2 bp 2 ep = non linear /104.5
3 bp 1 ep = pyramidal /107
4 bp no ep = tetrahedral /109.5
5 bp no ep = trigonal bipyramid / 120 between 2 close together bonds and 90
6 bp and no ep = octahedral / 90
4 bp 2 ep = square planar /90 but ep goes at opposite poles
Electronegativates trend
Down a group
There is a decrease in electronegativity going down the group
• The nuclear charge increases as more protons are being added to the nucleus
• However, each element has an extra filled electron shell, which increases shielding
• The addition of the extra shells increases the distance between the nucleus and the outer electrons resulting in larger atomic radil
• Overall, there is decrease in attraction between the nucleus and outer bonding electron
Across a period
• Electronegativity increases across a period
• The nuclear charge increases with the addition of protons to the nucleus
• Shielding remains relatively constant across the period as no new shells are being added to the atoms
• The nucleus has an increasingly strong attraction for the bonding pair of electrons of atoms across the period of the periodic table
• This results in smaller atomic radii
Permanent dipole
The molecule will always have a negatively and positively charged end
Forces between two molecules that have permanent dipoles are called permanent dipole - dipole forces
• The + end of the dipole in one molecule and the o- end of the dipole in a neighbouring molecule are attracted towards each other
Relative strength
For small molecules with the same number of electrons, permanent dipoles are stronger than induced dipoles
Hydrogen bonding
Hydrogen bonding is the strongest form of intermolecular bonding
• Intermolecular bonds are bonds between molecules
• Hydrogen bonding is a type of permanent dipole - permanent dipole bonding
• For hydrogen bonding to take place the following is needed:
• A species which has an O, Nor F (very electronegative) atom bonded to a hydrogen
• When hydrogen is covalently bonded to an O, N or F, the bond becomes highly polarised
• The H becomes so delta positvly* charged that it can form a bond with the lone pair of an O, N or F atom in another molecule
Enthalpy change of vapourisection op hydrides
The high enthalpy change of evaporation of water suggests that instantaneous dipole-induced dipole forces are not the only forces present in the molecule – there are also strong hydrogen bonds, which cause the high boiling point
High surface tension
Water has a high surface tension
• Surface tension is the ability of a liquid surface to resist any external forces (i.e. to stay unaffected by forces acting on the surface)
• The water molecules at the surface of liquid are bonded to other water molecules through hydrogen bonds
• These molecules pull downwards the surface molecules causing the surface of them to become compressed and more tightly together at the surface
• This increases water’s surface tension
Density
• Solids are denser than their liquids as the particles in solids are more closely packed together than in their liquid state
• The water molecules are packed into an open lattice
• This way of packing the molecules and the relatively long bond lengths of the hydrogen bonds means that the water molecules are slightly further apart than in the liquid form
• Therefore, ice has a lower density than liquid water by about 9%
Welling and boiling points + properties of water
Hydrogen bonding in water, causes it to have anomalous properties such as high melting and boiling points, high surface tension and a higher density in the liquid than the solid
Water has high melting and boiling points due to the the strong intermolecular forces of hydrogen bonding between the molecules in bothice (solid H2O) and water (liquid H2O)
• A lot of energy is therefore required to separate the water molecules and melt or boil them
Structural and stereoisonerism
Structural isomers are compounds that have the same molecular formula but different structural formulae
Stereoisomers are compounds that have the same atoms connected to each other, however the atoms are differently arranged in space
E/2 isomerism
Atom with the highest atomic number takes priority each side
Carbon monoxide+ nitrogen oxide equations
Carbon monoxide
2CO +O2—>2CO2
2CO +2NO—> 2CO2 + N2
Nitrogen oxide
2CO+ 2NO —>2CO2 +N2
Percentage uncertainty
Uncertainty
——————— x 100
Value taken
Remember uncertainty is range eg +_ 5 would have an uncertainty of 10.
