3.1.3 Bonding (physical chemistry) Flashcards
define metallic bonding
The strong electrostatic attraction between positive metal ions (cations) and a sea of delocalised electrons
draw a diagram for the bonding in:
-sodium
-aluminium
-magnesium
explain why the bonding in magnesium is stronger than the bonding in sodium
-Mg2+ ions have a higher charge than Na+ ions
-Magnesium has twice as many delocalised electrons per metal ion than sodium
-Mg2+ ions have a smaller ionic radius so have a higher charge density
-this means the attraction between the metal cations and the delocalised e- is stronger
4 properties of metals
1.conductivity
-good electrical and thermal conductors as the delocalised e- help transfer energy as they flow through the metal
2.strength
-very strong as there is a strong electrostatic attraction between cations and delocalised e-
3. malleable and ductile
-layers of metal ions can slide past one another
4.melting and boiling points
-directly linked to strength of metallic bond
-stronger the bonds, higher the melting point and boiling point
definition of covalent bonding
shared pair of electrons between 2 atoms
5 macromolecular structures
1.Diamond
2.Graphite
3.Graphene
4. Silicon
5.Silicon dioxide
allotropes
different structural forms of the same element
describe structure and bonding of diamond
-Each carbon atom is bonded to 4 other carbons by strong covalent bonds
-To melt, many strong covalent bonds need to be broken
-This requires alot of energy to overcome
-Melting point is very high
-Diamond is very strong
-No delocalised e- so doesnt conduct electricity
describe structure and bonding of graphite
-Each carbon is covalently bonded to 3 other carbon atoms - strong covalent bonds
-1 delocalised e-/ carbon atom and so can conduct heat and electricity as it can move through the structure
-Graphite is arranged in layers held together by weak intermolecular forces
-These layers can slide over eachother making graphite a soft structure
describe structure and bonding of graphene
-a single layer of graphite
-conducts electricity as it has delocalised e-
-very strong and has strong covalent bonds between atoms
simple molecular structures forces and properties
-low melting points and boiling points
-only relatively weak IMF need to be broken
-no delocalised electrons or ions so not conductors of electricity
Ionic bonding definition
The strong electrostatic attraction between oppositely charged ions
-e- are always transferred from the metal to the non-metal
-the metal will always form a cation and the non metal an anion
physical properties of ionic compounds
-high melting points
-strong electrostatic attraction between ions requires lots of energy to overcome
conductivity of ionic compounds
-as a solid no delocalised e- or mobile ions so doesnt conduct electricity
-when molten or aqueous ions mobile and can easily carry charge so conducts electricity
formula of ammonium
[NH4]+
formula of hydroxide
[OH]-
formula of nitrate
[NO3]-
formula of nitrite
[NO2]-
formula of hydrogencarbonate
[HCO3]-
formula of chlorate (I)
[ClO]-
formula of chlorate (V)
[ClO3]-
formula of carbonate
[CO3]2-
formula of sulfate
[SO4]2-
formula of sulfite
[SO3]2-
formula of dichromate
[Cr2O7]2-
formula of phosphate
[PO4]3-
define molecule
a group of atoms which are covalently bonded to one another
define coordinate bond
a shared electron pair which have both come from the same atom
describe a coordinate bond
-coordinate bonds are always represented by an arrow
-the direction of the arrow points from the atom which donates the pair of electrons to the atom which accepts the pair
-once a coordinate bond has been formed it behaves in exactly the same way as a covalent bond i.e same length and strength
electron repulsion strength
lone to lone pair>lone to bond pair>bond to bond pair
-the strength of the repulsions determines the bond angles between the bond to bond pair
3 bonding pairs
triganol planor
120 degrees
2 bonding pairs
1 lone pair
V-shaped
117.5 degrees
4 bonding pairs
Tetrahedral
109.5 degrees
3 bonding pairs
1 lone pair
triganol pyramidal
107 degrees
replace e with lone pair
2 bonding pairs
2 lone pairs
V-shaped
104.5 degrees
replace e with lone pair
5 bonding pairs
triganol bipyramidal
90 degrees
120 degrees
4 bonding pairs
1 lone pair
see saw
87.5 degrees
117.5 degrees
replace e with lone pair
3 bonding pairs
2 lone pairs
triganol planor
120 degrees
6 bonding pairs
octahedral
90 degrees
5 bonding pairs
1 lone pair
square pyramidal
87.5 degrees
4 bonding pairs
2 lone pairs
square planor
90 degrees
define electronegativity
the power of an atom to attract the pair of electrons in a covalent bond towards itself
how is electronegativity measured
-the pauling scale
-each element assigned a number between 0 and 4
-the higher the number the more electronegative it is
factors which affect how electronegative an element is
1.nuclear charge, as nuclear charge increases so does electronegativity
2. shielding, as shielding increases electronegativity decreases
3. atomic radius, as atomic radius decreases, electronegativity increases
(as a rule of thumb, the closer an element to fluorine, the more electronegative it is)
explain electronegativity across period 2
-nuclear charge increases across the period
-atomic radius decreases across the period
-electronegativity increases across the period
-nuclear attraction for the shared pair of electrons increases
polarity between Cl-Cl
-molecules made of atoms with no difference in electronegativity have their electrons distributed evenly
polarity between H-Cl
-The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical
-this produces a polar covalent bond
-the atom which has a greater electron density has a delta- charge and the atom which has a lower electron density has a delta+ charge
-H is delta positive and Cl delta negative
is BF3 polar
-all bonds polar
BUT
-molecule is symmetrical
-dipoles cancel out
-molecule not polar
is NH3 polar
-all bonds are polar
-molecule is unsymmetrical
-dipoles do not cancel out
-molecule is polar
is CHCl3 polar
-all bonds are polar
-molecule is unsymmetrical
-dipoles do not cancel out
-molecule is polar
what makes a molecule symmetrical
-all outer atoms the same
-triganol planor
-linear
-tetrahedral
-triganol bipyramidal
-octahedral
-square planor
three types of intermolecular forces
1.hydrogen bonding
2.permanent dipole dipole
3.induced dipole dipole
when hydrogen bonding occurs
-intermolecular force
-lone pair on N,O or F
H-F bond
H-O bond
H-N bond
when permanent dipole dipole intermolecular forces occur
-present in polar molecules
when induced dipole dipole intermolecular forces occur
-non-polar molecules
-The more electrons, the stronger the force
-between all molecules and group 18 elements
how does hydrogen bonding arise
-there is a very large difference in electronegativity between H and N/O/F
-this creates a large dipole on the bond
-The delta positive hydrogen strongly attracts a lone pair of electrons on the delta negative N/O/F in another molecule
how to draw hydrogen bonding
- Draw all lone pairs and dipoles
- Dashed hydrogen bond between delta + H and delta - O from another molecule
- Bond angle must be 180 degrees
how does permanent dipole dipole forces arise
-differences in electronegativity leads to polar bonds
-if dipoles do not cancel out due to symmetry there is an overall dipole
-delta positive on one molecule is attracted to the delta negative on another molecule
how to draw permanent dipole dipole forces
how does induced dipole dipole forces arise
-random movement of electrons leads to an uneven electron distribution and a transient/instantaneous/temporary dipole
-this induces a dipole in a neighbouring atom or molecule
-delta positive and delta negative dipoles attract one another
how to draw induced dipole dipole forces
hydrogen bonding in ice
-ice is less dense than liquid water
-2 hydrogen bonds per molecule
-holds molecules together in an open lattice structure
-more space between molecules than liquid water
hydrogen bonding in proteins
-proteins are held in a complex 3D shape by Hydrogen bonds between different parts of the chain
-3D shape is important for a functional protein
hydrogen bonding in DNA
-2 strands of DNA held together in a double helix by hydrogen bonds
describe the general trend in melting and boiling points across period 3
Na to Mg to Al
-Na, Mg, Al are all giant metallic lattices
-Electrostatic attraction between positive ions and delocalised electrons
-Ionic charge increases from Na to Mg to Al (+1,+2,+3)
-Number of delocalised electrons increases
-Metallic bonding is stronger
-More energy is required to overcome this
describe the general trend in melting and boiling points across period 3
Silicon
-Silicon is a macromolecular structure with strong covalent bonds between atoms
-Lots of energy is required to overcome the many strong covalent bonds
describe the general trend in melting and boiling points across period 3
P to S to Cl to Ar
-P to S to Cl to Ar are all simple molecular structures
-They have weak induced dipole dipole forces between molecules
P4 contains 60 e-
S8 contains 128 e-
Cl2 contains 34 e-
Ar contains 18 e-
-The more electrons there are the stronger the induced dipole dipole forces between molecules and so the more energy required to overcome this
-The melting points are much lower than Na to Si as induced dipole dipole forces are much weaker than metallic and covalent bonds
-These require less energy to overcome
Explain why Silicon has the highest melting point but Aluminium has the highest boiling point
-When Al is molten there is still attractions between cations and delocalised electrons
-To vaporise it, strong metallic bonds still need to be broken
-Alot of energy is required to overcome this
-When Si is molten the covalent bonds are already broken
-Less energy is required to vaporise it
