Chemical bonding Flashcards
- Metallic
BONDs
electrostatic forces of attraction between metal cations and the sea of delocalised electrons. Bonds are non directional
STRUCTURE
metal lattice comprising of layers of metal cations immersed in sea of delocalised electrons. (SODE) Each metal atom contributes its loosely bound electrons to form SODE
Physical properties of metals
- High MP/BP
- Good electrical conductors
- Thermal conductivity
- Malleable
- Soluble in other metals
Factors affecting strength of metallic bonds
- amount of electrons contributed by each metal atom to SODE (greater, stronger)Io
- Ionic radius (smaller, stronger)
Ionic
BONDs
efa between oppositely charged anions and cations in an ionic compound. ionic bonds are non directional. Each ion is attracted equally to several oppositely charged surrounding ions
STRUCTURE
involves complete transfer of electrons from an atom (forms cation) to another atom (forms anion)
Giant ionic crystal 3D lattice structure
oppositely charged ions held in fixed positions in an orderly manner in a regular structure. (FOR)
Factors affecting strength of ionic bonds
In other words, factors affecting magnitude of L.E.
|L.E| ∝ |q+ x q- / r+ + r- |
1. Ionic radius (Interionic distance of compound) shorter, greater
2. Charge of ions (Product of charges of anion and cation) greater, stronger
Physical properties of ionic solids
- High MP/BP
- Hard and brittle
- Different electrical conductivity in different physical states
- in aq
- as a solid
- molten - Soluble in water and other polar solvents (usually)
Covalent bonds
BONDs
efa of positively charged nuclei of each bonding atom for the shared pair of electrons/the efa between the electron density and the bonded nuclei. Sharing a pair of electrons between two atoms of similar electronegativity
Is formed as a result of maximal overlap of valence atomic orbitals cont. one electron each. Maximal overlap ensures that the electron density is concentrated between the nuclei and can hold the atoms against the mutual repulsion of the nuclei. Bonds are localised and directional as electron pair is confined between nuclei of two bonding atoms
Structure
Simple covalent structure
Giant covalent compounds/structures
Dative covalent bond
When both electrons come from only one of the atoms
When a filled valence orbital of an atom overlaps with a vacant valence orbital of another atom. An atom donates a lone pair of electrons to another atom which has an empty low lying orbital to accommodate the electrons.
denoted by “ —–>”
examples AlCl3 , NH4 ion
Types of covalent bond
- Sigma σ
the collinear overlap of two atomic orbitals
there can only be one sigma bond between two atoms - pi π
the collateral overlap of two atomic p orbitals
π vs σ strength
Sigma is stronger as there is greater degree of overlap of orbitals than in pi bonds
a pi bond is formed after a sigma bond is formed. so pi bonds are only present in multiple bonds (double bonds/ triple bonds)
Bond length
the distance between the nuclei of the two bonding atoms in a covalent bond
a balance between
1. the maximum attraction between the nuclei for the shared electron density
2. the minimum repulsion between two positive nuclei and between electron clouds
Stronger the covalent bond shorter the bond length
Bond energy
can be used to infer degree of orbital overlap between two bonding pair of atoms and hence the strength of covalent bond.
The energy required to break one mole of covalent bonds between atoms in a gaseous molecule
Factors affecting strength of covalent bonds
BEBT
- Bond order
- the number of covalent bonds formed between two atoms (double bonds-2/ triple bonds-3)
- higher the bond order greater the no.of orbitals overlapped. increased electron density bet.bonding atoms greater att. bet. bonding nuclei and shared electrons. stronger covalent bond - Effectiveness of orbital overlap
- more effective the orbital overlap, stronger the bond
- large orbitals are more diffuse. hence when large orbitals overlap with each other there is less effective orbital overlap and lower percentage of electron density between the two nuclei. hence bond weaker - Bond polarities
- Presence of partial charges increase the attraction between bonding atoms on top of existing cov. bond. The more electronegative the atom, the greater the partial charges that arise - *type of hybridisation
Physical properties of covalent cmpds
Simple molecular (discrete molecules)
- Low MP/BP
- Soluble in non polar organic solvents usually
- non electrical conductors
Giant covalent
- Very high MP/BP
- Non electrical conductors (except graphite)
- Hard
- Insoluble in all solvents
- Slippery/lubricating property (Graphite)
- Electrical conductor (Graphite)
Giant covalent
Diamond/ Silicon dioxide
- Rigid 3D tetrahedral structure. Each C atom bonded to 4 other C/ Each Si bonded to 4 O
by strong covalent bonds throughout lattice
Graphite
-Network of planar hexagonal ring layers
- C bonded to 3 other C by covalent bonds
- Between each layers weak intermolecular forces
Intermediate bond types
- Pure ionic bond
- ions exist as discrete point charges with no electron density bet.them
- does not occur in reality in ionic cmpds - Polarised ionic bond
- cations attracts and distorts electron cloud of anion
- partial sharing of electrons as electron density drawn into region bet.2 nuclei
-occurs in ionic cmpds
3.Polar covalent bond
-Electron density not symmetrically distributed in a bond bet. diff atoms
- more electronegative atoms has a greater share of the shared electron density
- there is permanent separation of partial charges and compds exist as polar molecules - Pure covalent bond
- Electron density symmetrically distributed in a bond bet. identical atoms
like Br2
Factors affecting polarisation
- Charge to size ratio of the cation/ Charge density
- Size of the anion
greater the polarisation greater the covalent character in the cmpd
Molecular geometry and polarity
refer to chart
VSEPR theory
the geometry of a covalent compd affects its polarity and its physical properties
Deviation of bond angle predicted by VSEPR theory
due to
- Different electronegativity of central atoms given same terminal atoms
as electronegativity of the central atom increases electron density in the bond pairs is drawn closer to the central atom. Increased electron density around the central atom results in increasing bond pair-bond pair repulsion. Hence bond angle increases
-Different electronegativity of terminal atoms given same central atoms
electron density of bond pairs drawn to the more electronegative terminal atoms away from central. Lesser electron density around central atom hence lesser bond pair bond pair repulsion hence bond angle decreases
Polar and Non polar molecules
greater the difference in electronegativity the greater the ionic character in covalent bond
polar : there is a difference in polarity and partial charges arise
individual bond dipoles do not cancel out each other giving rise to net dipole moment
non polar : no partial charges arise due to negligible difference in electronegativity
individual bond dipoles cancel out each other
How to determine the polarity of a molecule
- Polarity of bonds in the molecule
2. Shape/Geometry of molecule
instantaneous dipole- induced dipole
interaction between non polar molecules and noble gases
As electrons are in constant motion at some instant, there is a temporary shift of electrons to one side of the atom and results in an instantaneous dipole
the instantaneous dipole induces a similar dipole on an adjacent atom an induced dipole is formed
this repeats
dipoles are temporary but net attraction they produce is permanent
Factors affecting strength of id-id
- Size of electron cloud
2. Surface area in contact between molecules
Permanent dipole-permanent dipole
interactions between polar molecules
stronger than id-id
Hydrogen bonding
efa between protonic H atom in H-F / H-O/ H-N bond and a lone pair on an electronegative atom F,O,N in a neighbouring molecule
When H is bonded to a highly electronegative atom, the highly electronegative atom attracts bonding electrons towards itself leaving H with a very small share of the electron pair. H acquires a large partial positive charge and electronegative atom a large partial negative charge. H behaves like a bare proton as a result
Conditions required for H-bonds
- Protonic H in H-F/H-O/H-N
2. Lone pair(s) of electrons on an electronegative F,O,N atom in neighbouring molecule
Factors affecting strength of hydrogen bonds
- Extensiveness of hydrogen bonding (greater, stronger)
determined by average no.of H-bonds formed per molecule
which can be determined by
-counting total no.of protonic H per molecule
-count total no.of lone pairs on F,O,N
and taking the lower value of the two - Polarity of H-Y bond (more electronegative Y, stronger)
Boiling melting point of cmpds
SSCLC
can also be determined by the electrons cloud size of species when considering strength of id-id vs pd-pd vs H bond
Solubility of compounds in solvents
3 types of interactions involved in determining solubility of a solute in a solvent
- solute-solute interaction
- solvent-solvent interaction
- solute-solvent interaction
PSSSI
Solubility in ionic cmpds
ion dipole interactions formed in polar solvents
for eg in water, ion-dipole between oppositely charged ions and water molecules
ion dipole overcomes ionic bonds ionic lattice breaks down
acid disassociates in water. disassociated ions form ion dipole w water
sufficient to overcome H-bond
between water and covalent bond bet. acid molecules
Structure and properties of ice
Ice is less dense than water which is anomalous
why?
Dimerisation
Mr of molecule doubles
of ethanoic acid
-dimer formed due to H-bonds between two molecules
dimer not formed in (aq) state why?
AlCl3—-> Al2Cl6
-via dative bonding (describe the process)
NO2—> N2O4
-via covalent bond ( describe)
Intramolecular H-bonds
2-nitrophenol vs 4-nitrophenol
are isomers
2-nitrophenol lower BP than 4-nitrophenol
due to NO2 and OH groups, 2 nitrophenol forms intramolecular H bonds whereas 4 nitrophenol forms only intermolecular hydrogen bonding
less sites available for intermolecular H bonding in 2 nitrophenol
more energy required to overcome intermolecular H bonding in 4 nitrophenol hence high BP
