Structure and Bonding (Physical) (complete) Flashcards
Describe the structure and bonding of a metal element.
- (m.p and b.p + how affected)
- solubility in water
- electrical conductivity
- strength
- giant lattice structure of (named) metal ions.
- electrostatic attractions between positive metal ions and delocalised negative electrons.
- contains metallic bonding, so high melting and boiling points, increased by larger + charge therefore more outer e-.
- insoluble in water.
- conducts electrically due to delocalised electrons.
- strong however malleable and ductile.
Describe the structure and bonding of an ionic compound.
- (m.p and b.p + how affected)
- solubility
- electrical conductivity
- strength
- lattice structure of positively and negatively charged ions.
- ionic bonding - electrostatic forces of attraction between positive and negative charges.
- high melting and boiling points, electrostatic attractions strong, increased by larger charge size and smaller ion size.
- dissolves in water as attracts water due to both -ve and +ve charges.
- conducts electrically when dissolved/melted in water.
- brittle, can shatter easily
Describe the structure and bonding of a giant covalent structure.
- (m.p and b.p + how affected)
- solubility
- electrical conductivity
- strength
- giant lattice structure in which all atoms are joined to others by covalent bonds, continuous pattern with no inter-molecular forces.
- very high melting and boiling points as need to break strong covalent bonds, increased by stronger covalent bonds.
- insoluble in water.
- (diamond/Si/SiO2) don’t; (graphite/graphene) do.
- strong.
Describe the structure and bonding of a simple molecular structure.
- (m.p and b.p + how affected)
- solubility
- electrical conductivity
- strength
- individual molecules with weak forces between them, (atoms within molecules joined by covalent bonds)
- low melting and boiling points due or weak forces between molecules, increased by stronger inter-molecular forces.
- insoluble in water.
- don’t conduct electrically.
- brittle.
Describe the structure and bonding of a monatomic element.
- (m.p and b.p + how affected)
- solubility
- electrical conductivity
- individual atoms with very weak forces between them, contain no bonding.
- very low melting and boiling points due to weak forces between atoms.
- insoluble in water.
- don’t conduct electrically.
How do you strengthen covalent bonds?
- the shorter the bond, the stronger the bond (usually).
- triple bonds stronger than double bonds stronger than single bonds.
How to identify a co-ordinate bond?
- where both the electrons being shared in a covalent bond come from the same species.
- drawn with arrows rather than sticks to show the bond, coming from the species with both electrons.
- identical to covalent bonds once formed.
Name the 12 potential shapes (taking into account lone pairs) that molecules can be.
- give the:
> total number of electron pairs
> the number of bonding pairs of electrons
> the number of lone pairs of electrons.
- give the bond angles present.
- linear - 2 total e- pairs ; 2 bonding pairs ; 0 lone pairs ; 180 degrees
- trigonal planar - 3 total 3- pairs ; 3 bonding pairs ; 0 lone pairs ; 120 degrees
- bent (v-shape) - 3 total e- pairs ; 2 bonding pairs ; 1 lone pair ; 118 degrees
- tetrahedral - 4 total e- pairs ; 4 bonding pairs ; 0 lone pairs ; 109.5 degrees
- trigonal pyramidal - 4 total e- pairs ; 3 bonding pairs ; 1 lone pair ; 107 degrees
- bent (v-shape) - 4 total e- pairs ; 2 bonding pairs ; 2 lone pairs ; 104.5 degrees
- trigonal bipyramidal - 5 total e- pairs ; 5 bonding pairs ; 0 lone pairs ; 90 degrees and 120 degrees
- n/a name - 5 total e- pairs ; 4 bonding pairs ; 1 lone pair ; 119 degrees and 89 degrees
- t-shape - 5 total e- pairs ; 3 bonding pairs ; 2 lone pairs ; 89 degrees
- octahedral - 6 total e- pairs ; 6 bonding pairs ; 0 lone pairs ; 90 degrees
- square pyramid - 6 total e- pairs ; 5 bonding pairs ; 1 lone pair ; 89 degrees
- square planar - 6 total e- pairs ; 3 bonding pairs ; 3 lone pairs ; 90 degrees
Explain the reasoning behind the certain shape of a molecule if:
- the shape has no lone pairs.
- the shape has lone pairs.
- no lone pairs - the bonding pairs repel equally, so are placed as far apart from each other as possible to minimise repulsion.
- lone pairs - lone pairs repel more than bonding pairs, so the bonding pairs are forced closer together to minimise repulsion.
Define electronegativity.
Name and explain the three factors that affect electronegativity.
- The power of an atom to attract the two shared pair of electrons in a covalent bond.
- Number of protons - more protons can more easily attract electrons (opposite charges)
- Nuclear shielding - less nuclear shielding means less between attracting nucleus and electrons.
- Atomic size - closer to the nucleus, more attraction between nucleus and electrons.
Define electronegativity.
Name and explain the three factors that affect electronegativity.
- The power of an atom to attract the two shared pair of electrons in a covalent bond.
- Number of protons - more protons can more easily attract electrons (opposite charges)
- Nuclear shielding - less nuclear shielding means less between attracting nucleus and electrons.
- Atomic size - closer to the nucleus, more attraction between nucleus and electrons.
Describe the trend in electronegativity:
- down a group of
- across a period
- down a group - general decrease; larger atomic radius, more nuclear shielding, so less attraction between nucleus and bonding pair of electrons.
- across a period - general increase; atomic radius decrease but number of protons increase and same nucleus shielding, so more attraction between nucleus and bonding pair of electrons.
What does it mean when a covalent bond is:
- Polar (when does it happen?)
- Non-polar (when does it happen?)
Define a bond-dipole moment.
- Polar - when the two atoms in a covalent bond have different electronegativity, so the two electrons are not shared equally.
- Non-polar - when the two atoms in a covalent bond have identical electronegativity, so the two electrons are shared equally.
- occurs in polar bonds, it is a measure in the strength and direction of the polarity in the bond. the bigger the difference in electronegativity, the greater the bond-dipole moment.
Name the weakest type of intermolecular forces (usually) and how and where they occur.
- how can they be strengthened?
- Van der waals forces -
> present in all molecular substances between the atoms.
> occur because the e- are constantly moving around randomly, therefore an uneven electron distribution at any one time, causing a temporary dipole within a molecule.
> this temporary dipole induces a temporary dipole in a neighbouring molecule, then there is an attraction between the two. this is a temporary induced dipole-dipole attraction (van der waals force).
> can be strengthened by having a bigger molecule (so more electrons) to create more dipole-dipole moments.
Name the second weakest (usually) intermolecular force, and how and where they occur.
- Permanent dipole-dipole attractions -
> present between polar molecules, (not necessarily polar bonds)
> formed due to permanent dipole-dipole moments being created between atoms, as the dipoles are always int eh same place, therefore are ‘permanent’.
