Bonding Flashcards
Metals have
Low ionisation energy meaning positive ions are formed, delocalised electrons meaning they can move freely in the lattice
Metallic bonding
Electrostatic attraction of the positively charge metal ions to the delocalised electrons, non-directional bond
Non-directional charge
Attraction is equal in all directions
Properties of metallic bonding
High melting and boiling points, high thermal and electrical conductivity, malleable and ductile
Why does metallic bonding cause high melting and boiling points
strength of metallic bonds between cations and delocalised electrons determined by No. delocalised electrons and atomic radius (Increased number of valence electrons= increased melting and boiling points, smaller atomic radius= stronger metallic bond
Why does metallic bonding cause high thermal and electrical conductivity
When heated, delocalised electrons gain energy and move faster, colliding with other electrons and transferring energy, delocalised electrons can move freely in a solid state
Why does metallic bonding mean metals are malleable and ductile
when force is applied, delocalised electrons shield the cations from each other to they slide past one another without breaking the bonds meaning it changes shape
Covalent bonding
Results from the electrostatic attraction between a shared pair(s) of electrons and adjacent positive nuclei, directional bonds
Directional bonding
not equal attraction between atoms
Covalent molecular substances
Made up of molecules with the atoms within the molecules held together by the electrostatic attraction between the shared electrons and positive nuclei, strong covalent bond but weak intermolecular forces
Properties of covalent bonding
Low melting and boiling points, form soft solids, non-conductors and insulators
Why does covalent bonding cause low melting and boiling points
Weak intermolecular forces meaning less energy is required to separate molecules
Why does covalent bonding form soft solids
Weak intermolecular forces
Why does covalent bonding cause non-conductors and insulators
No free electrons meaning molecules are uncharged
Examples of covalent bonding
Cl2- share 2 electrons, H2O- two single covalent bonds
Octet rule exceptions
Beryllium is stable with 4 electrons, Boron is stable with 6 electrons, Phosphorous has 10 electrons in its valence shell- PCl5, Sulfur has 12 electrons in its valence shell- SF6
Covalent network substances
Each atom is covalently bonded to each other to form a 3 dimensional lattice, no weak intermolecular forces because each atom is covalently bonded
Properties of covalent network structures
High melting and boiling points, hard and brittle, no conductivity
Why do covalent network structures have a high melting and boiling point
Many strong covalent bonds which extend through the lattice which need a lot of energy to break apart
Why are covalent network structures hard and brittle
Hard due to the strong covalent bonds in the lattice, shatter if hammered with enough energy
Why do covalent network substances have no conductivity
No mobile ions or delocalised electrons
Covalent network substance exeption
Graphite has conductivity as it has one delocalised electron per carbon atom
Examples of covalent network substances
Diamond, silicon, carbide, quartz, granite
Ionic bonding
Result of the electrostatic attraction between positive ions (cations) and negative ions (anions), non-directional
What do metal atoms do to form cations
Lose elctrons
What do non-metal atoms do to form cations
Gain electrons
Stable arrangement in ionic bonding is achieved by
the transfer of electrons resulting in the formation of atoms held together in a 3 dimensional crystal lattice by electrostatic attraction
Properties of ionic bonding
High melting and boiling points, poor electrical conductors when solid, good electrical conductors when molten or in aqueous solution, brittle
Why do ionic bonds cause high melting and boiling points
strong electrostatic forces which extend through the lattice, through heating, ions vibrate faster and becomes less firmly held in place
Why does ionic bonding cause poor electrical conductors when solid
When solid, ions are held tightly together by electrostatic bonds meaning ions can’t move
Why does ionic bonding cause good electrical conductivity in a molten state or in an aqueous solution
Once the electrostatic forces are broken the ions are free to move meaning they carry charge, ionic compounds that are soluble in water means water molecules come between the ions breaking the lattice meaning the ions are mobile
Why does ionic bonding cause susbtances to be brittle
When force is applied, the ions shift positions so the same charge may come together, likes charges repel each other meaning the ions repel each other meaning the latice breaks
Examples of ionic bonding
NaCl, CaCl2
Allotropes
Elements that can exist with their atoms in several different structure arrangements which give them different physical forms, bonded in different ways
Examples of allotropes
Oxygen- O2 and O3, Carbon- diamond, graphite and amorphous carbon
Properties of diamond
Very hard, sublimes, non-conductor, brittle
Structure of diamond
Covalent network structure
Structure of graphite
Covalent layer lattice
Properties of graphite
soft, greasy, conductive, slippery
Structure of amorphous carbon
irregular structure, many varieties
Properties of amorphous carbon
Non-crystalline, cheap, conductive
Nanomaterial
Molecule smaller than 100 nanometers
Properties of nanomaterials
Increased strength, Increased chemical reactivity because of larger surface area, increased conductivity
Example of a nanomaterial
Carbon nanotubes
