Bonding and Structure Flashcards
covalent bonding definition
electrostatic attraction between positive nuclei and the shared pair of negative electrons.
list of metallic elements
Li, Be, Na, Mg, Al, K, Ca
list of covalent molecular (diatomic) elements
H2, N2, O2, F2, Cl2
list of discrete covalent molecular elements
P4, S8, C60 (fullerene)
list of covalent network elements
B, C, Si
list of monatomic elements
He, Ne, Ar… (noble gases)
metallic bonding definition
electrostatic attraction between positive metal ions and a sea of delocalised negative outer electrons
London dispersion forces definition
the weak forces between all atoms and molecules including monatomic elements. (weakest form of intermolecular bonding)
strength of LDF
the strength of LDF is related to the number of electrons within an atom or molecule. The more electrons the stronger the LDF will be.
difference between intermolecular and intramolecular forces
intramolecular- within molecules
intermolecular- between molecules
metallic bonding properties
- conduct electricity when solid or liquid due to delocalised electrons.
- high density due to closely packed lattice structure
- usually high bp (lots of bonds to break)
mp and bp of metallic elements trend
- increase across period because more electrons in outer shell, the stronger the metallic bond.
- decrease down a group because the additional filled electron shells weaken electrostatic forces between nuclei and delocalised electrons.
monatomic elements properties
- exist as single atoms
- form no bonds due to full outer shell (no intramolecular bonding)
- do not conduct electricity because there are no free electrons.
- contain LDF
covalent diatomic properties
-low mp/bp because it is the weak intermolecular forces which break between the molecules.
covalent diatomic mp/bp trend
down the group, mp/bp increase because the strength of the LDF is increasing due to larger atoms with more electrons.
Fullerene
- one of the three forms of carbon.
- smallest is C60, known as buckminsterfullerene
- does not conduct electricity as no free electrons.
- NOT A COVALENT NETWORK
- large molecule=strong LDF
covalent network properties
- every atom linked to another by strong covalent bonds
- high mp/bp because lots of energy needed to break strong covalent bonds
covalent network forms of carbon
DIAMOND
-tetrahedral structure
-does not conduct electricity (no free electrons)
-hardest natural substance (used for drills/cutting tools)
GRAPHITE
-layered structure with LDF between layers
-very soft (layers break away due to weak LDF)
-conducts electricity (delocalised electrons)
allotrope
different form of same element
discrete covalent molecular mp/bp
low because it is the weak intermolecular forces which break between molecules not covalent bonds within molecule
noble gases mp trend
increases down the group as larger atoms with more electrons mean stronger LDF.
3 types of Intramolecular bonding
- covalent molecular (polar or non-polar)
- Metallic
- Ionic
covalent bonding definition
- shared pair of electrons
- both nuclei try to pull electrons towards themselves
- creates a strong bond between the two atoms.
pure/non-polar covalent bonding definition
when both atoms have an equal ‘pull’ on the shared electrons because they have the same electronegativity.
the electrons sit in the middle of the two nuclei.
Eneg difference 0.0-0.3
polar covalent bond definition
- a bond formed when the shared pair of electrons in a covalent bond are not shared equally.
- this is due to different elements having different electronegativities.
- Eneg difference 0.4-2.5ish
dipole moment
shows direction electrons are being pulled in
the bonding continuum
a spectrum which starts at pure/non-polar covalent and ends at ionic
how to work out if a molecule is polar or non-polar
- is there a polar bond in the molecule (E difference>0.3) –> if not, non-polar
- is the molecule symmetrical? yes=non-polar no=polar
testing polarity
non-polar: charged rod does no attract or repel a stream of water.
polar: charged rod attracts/repels stream of water.
Ionic bond
- usually between metal and non-metal with a large E difference
- forms a lattice of + and - ions
- electrostatic force of attraction between positively and negatively charged ions
- non metal (high eneg) gains electron to form negative ion
- element with low eneg loses electron to form positive ion.
- formula tells you ratio of ions
increasing ionic character/polarity
increasing electronegativity difference
3 types of intermolecular/ Van der Waals’ forces
- London Dispersion Forces
- Permanent dipole-permanent dipole interactions
- Hydrogen bonds
what do intermolecular bonds dictate
melting/boiling points of molecules
are intra or inter stronger?
intramolecular bonds are stronger (what holds atoms together)
permanent dipole-permanent dipole
- occurs between polar molecules
- the permanent dipole in one molecule is attracted to the permanent dipole of another
- higher mp/bp than just LDF because pdp-pdp are stronger and take more energy to break.
what is needed for a fair comparison between the mp/bp of a polar and non-polar molecule
similar formula mass
polar or non-polar if only has LDF
non-polar
polar or non-polar if molecule has LDF and pdp-pdp
polar
hydrogen bonding
- strongest intermolecular bonding
- occurs between molecules where there is an atom of H joined to an atom of N,O,F.
- higher mp/bp as they are held most tightly together.
- extreme case of pdp-pdp due to very large difference in electronegativity
properties related to intermolecular forces: melting and boiling points
stronger bonds (pdp-pdp/hydrogen) require more energy to break as molecules are held more tightly together.
viscosity
how thick/gloopy a liquid is
increasing viscosity trend
more OH groups=more H bonding=molecules are held more tightly=more viscous
miscibility
how well liquids mix together
miscible liquids mix thoroughly without any visible boundary
why are some liquids immiscible
if one is polar, other is non-polar; unable to bond to one another
what molecules will mix well with water
polar molecules
why does ‘like dissolve like’
due to attraction between the charges in each compound
polar substance+polar solvent–>soluble
density
how tightly packed the particles are
density of ice
ice is less dense than liquid water because hydrogen bonding between H20 molecules in ice gives a very open and expanded structure.
what decides bonding type
difference in Eneg can decide bonding type
ultimately properties of a substance which dictate bonding type
ionic bonding definition
electrostatic forces of attraction between positive and negative ions, usually between metal and non-metal.
delocalised definition
delocalised electrons, in metallic bonding, are free from attachment to any one metal ion and are shared amongst the entire structure.
dipole definition
an atom or molecule in which a concentration of positive charges is separated from a concentration of negative charges.
isoelectronic
means having the same arrangement of electrons
Eg: Neon, Na+ and Mg2+
lattice
a lattice is a regular 3D arrangement of particles in space. The term is applied to metal ions in a solid, and to positive and negative ions in an ionic solid.
lone pairs
pairs of electrons in the outer shell of an atom which take no part in bonding.
