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’.
Name the strongest type of intermolecular force, and how and where they occur.
- Hydrogen bonding -
> special case of permanent dipole-dipole attractions, occur between a hydrogen atom and a very electronegative atom. (e.g: fluorine, nitrogen, oxygen)
> this is because the polar bond between the H and N/O/F leaves the H nucleus exposed, as it only has one electron., so there is a STRONG attraction between the lone pair on the N/O/F to the exposed nucleus of the H atom.
Describe the solubility of Ionic compounds:
- Water (polar)
- Alkanes (non-polar)
- Water - many Ionic compounds dissolve in water. Attractions form between polar water molecule and ions. (delta -ve oxygen attracted to positive ion) ( delta +ve hydrogen attracted to negative ion)
- Alkanes - Insoluble
Describe the solubility of Compounds with hydrogen bonding:
- Water (polar)
- Alkanes (non-polar)
- Water - usually dissolve in water. water has hydrogen bonds and substances with hydrogen bonds can form attractions through hydrogen bonds to the water molecules.
- Usually insoluble (sometimes slightly soluble. (hydrogen bonds not attracted to non-polar molecules).
Describe the solubility of Non-polar substances in:
- Water (polar)
- Alkanes (non-polar)
- Water - Usually insoluble or only slightly soluble (attractions between atoms stronger than to water)
- Alkanes - Usually dissolve well - intermolecular forces form between the solvent and solute molecules) (non-polar dissolves non-polar).
