Covalent bonding Flashcards
Covalent bond definition
The electrostatic force of attraction between the 2 nuclei of atoms and the shared pair of electrons in the bond
How is a covalent bond formed?
When an atomic orbital with a single electron in it from one electron overlaps with an orbital of another atom that contains an electron
What atomic orbitals overlap in covalent bonding?
Any 2
4 ways an orbital can overlap in covalent bonding
2 ends of s orbitals overlap
2 ends of p orbitals overlap
Sideways overlap of p orbitals
End of s and p orbitals overlap
What is a sigma bond?
When the ENDS of 2 atomic orbitals overlap eg 2s orbital ends or 2p orbitals end
What is a pi bond?
When a sideways overlap is formed eg 2p orbitals bond sideways
When can a pi bond be formed?
Can only be formed if sigma bond had already been formed = only in double or triple bonded atoms
S orbital and p orbital overlap
Overlaps end on
Only those of DIFFERENT elements eg not Cl-Cl
Forms a polar covalent bond due to invariably of the elements
Examples of formation of sigma bonds which result in covalent bonding
H2 because both electrons exist in 1s orbital
Ends of these overlap to form a new molecular orbital, and these electrons occupy this in the overlap
Therefore the highest electron density is between the 2 nuclei of the hydrogen
Examples of formation of pi bonds which result in covalent bonding
In ethene where in C=C, one bond between carbon formed is a sigma bond whereas the other is pi, this is because of the way the p subshell is structured that a head on overlap between 2 p orbitals in different atoms (sigma bond) can result in up to 2 sideways overlaps (pi bonds)
Electron density of a pi bond
highest above and below the molecule where the 2 orbitals overlap
Bond length
Distance between between nuclei of 2 atoms that are covalently bonded together
This distance is the point where the repulsion between nuclei and attraction nuclei has to electrons is balanced
Bond strength
Amount of energy required to break every bond in one mole of that bonded molecule when it’s in a gaseous state
General relationship between bond strength and bond length between bonds of a similar nature
Higher the electron density (eg number of electrons in a bond) the stronger force of attraction the nucleus has to shared electrons
So higher bond enthalpy (more energy stored) and shorter bond chain because the higher forces pull nuclei closer
Why is it important to only compare similar molecule with each other in terms of bond comparing strength and length?
Because there are extraneous reasons as to why a bond strength is weaker, rather than just a larger bond length ie a lone, non-bonded pair of electrons present repel each other = weaker bond
Electronegativity
An atom’s electronegativity is its ability of an atom to attract a bonding pair of electrons
Trend in electronegativity in periodic table
DECREASES going DOWN the group
INCREASES from LEFT to RIGHT
Most electronegative element
Fluorine
Electron density
Shows regions around nuclei in bonded atoms where the electrons are distributed (Eg in the overlap of orbitals)
Electron density in identically bonded atoms
The distribution around the nuclei will be roughly symmetrical for each atom
Because it’s the same element = same electronegativity = so same ability to attract an element
What is formed when 2 atoms of different electro-negatives bond?
A polar covalent bond
Contour lines in a polar covalent molecule
the lines are closer packed to atom of higher electronegativity showing it has a higher density than the other
Charges in a polar molecule
The atom with the largest electronegativity thus higher electron density becomes slightly negatively charged so represented as d-
Whereas atom with lower electron density near it becomes d+
Ways of representing a polar bond in a stick diagram of a covalent bond
Have an arrow point in the direction of ‘electron drift’ so which has a higher electron density (so higher electronegativity)
Eg: H–>–Cl
