Topic 4.3: Covalent structures Flashcards
Definition of lewis structure
A diagram of molecules in which the valence e- of the atom are represented by dots, and the sharing of e- to form a covalent bond is shown
How to write a lewis structure (6 steps)
a) Calculate the total number of valence e- in the molecule
b) Draw the skeletal structure of the molecule
c) Determine the central atoms (the lowest electronegative)
d) Draw single bonds to the central atom
e) Put all remaining valence e- on atoms as lone pairs
f) Turn lone pairs into double or triple bonds to give every atom an octet (except H)
g) Put the Lewis structure in a square bracket with the charge shown outside if it’s an ion
Definition of coordinate covalent bonds
Type of covalent bond in which both shared electrons in a molecule come from the same atom
Symbol of a coordinate covalent bond
An arrow on the head of the bond is used to indicate the origin of the electron pair
Definition of “octet rule”
Tendency of atoms to gain a valence shell with a total of 8 e-
Exceptions to octet rule
a) Be and B form stable molecules such as BeCl2 and BF3 (incomplete)
b) Elements in period 3 and below may expand their octet by using d-orbitals
Definition and properties of electron deficient molecules
Molecules with incomplete octets, which tend to accept an electron pair from a molecule with a lone pair. (NH4)
Valence Shell Electron Pair Repulsion (VSEPR) Theory
The total number of electron domains (# lone or bond pairs) determines the shape of a covalent molecule
a) Electron domains in the same valence shell carry the same charge, they repel each other and spread themselves as far as possible
Explanation of the order of electron repulsion among lone pair and bonding pair
LP – LP > LP – BP >
BP – BP
a) Orbitals that hold lone pairs are rounder and shorter and spread more easily
b) Orbitals that hold bonding pair are more elongated
Tendency of # lone pairs and bond angles
Since lone pairs cause more repulsion than bonding pairs, the angle is reduced as the # lone pairs increase
Difference between electron domain geometry and molecular geometry
a) The electron domain geometry is determined by the positions of all the electron domains,
b) Molecular geometry (arrangement of atoms in space) depends on the positions of the bonded atoms
Shape of molecules with 2 electron domains and 0 lone pairs
a) Molecular geometry
b) Angle
c) Example
a) Linear
b) 180°
c) CO2
Shape of molecules with 3 electron domains and 0 lone pairs
a) Molecular geometry
b) Angle
c) Example
a) Planar triangular
b) 120°
c) BF3
Shape of molecules with 3 electron domains and 1 lone pair
a) Molecular geometry
b) Angle
c) Example
a) Bent (V shaped)
b) < 120°
c) SO2
Shape of molecules with 4 electron domains and 0 lone pairs
a) Molecular geometry
b) Angle
c) Example
a) Tetrahedral
b) 109.5°
c) CH4
Shape of molecules with 4 electron domains and 1 lone pair
a) Molecular geometry
b) Angle
c) Example
a) Triangular pyramidal
b) < 109.5°
c) NH3
Shape of molecules with 4 electron domains and 2 lone pairs
a) Molecular geometry
b) Angle
c) Example
a) Bent (V shaped)
b) < < 109.5°
c) H2O
Shape of molecules with 5 electron domains and 0 lone pairs
a) Molecular geometry
b) Angle
c) Example
a) Triangular bipyrimidal
b) 90° / 120° / 180°
c) PF5
Shape of molecules with 5 electron domains and 1 lone pair
a) Molecular geometry
b) Angle
c) Example
a) See saw
b) 180° / 90° / < 120°
c) SF4
Shape of molecules with 5 electron domains and 2 lone pairs
a) Molecular geometry
b) Angle
c) Example
a) T-shaped
b) 90° / 180°
c) BrF3
Shape of molecules with 5 electron domains and 3 lone pairs
a) Molecular geometry
b) Angle
c) Example
a) Linear
b) 180°
c) I5 -
Shape of molecules with 6 electron domains and 0 lone pairs
a) Molecular geometry
b) Angle
c) Example
a) Octahedral
b) 90° | 180°
c) SF6
Shape of molecules with 6 electron domains and 1 lone pair
a) Molecular geometry
b) Angle
c) Example
a) Square pyramidal
b) 90° | 180°
c) BrF5
Shape of molecules with 6 electron domains and 2 lone pairs
a) Molecular geometry
b) Angle
c) Example
a) Square planar
b) 90° | 180°
c) XeF4
Conditions for a molecule to be dipole
a) It must have polar bonds
b) Arrangement of atoms is asymmetrical. Dipole bonds will oppose each other and cancel out.
Definition of resonance structures
Set of two or more LS that collectively describe the electronic bonding a single polyatomic species
Definition of resonance hybrid
Actual structure described by individual resonance structures
a) Bonding e-, insted of being confined to one location, are shared between more than one bonding positions (delocalized), providing more stability.
Types of covalent structures
a) Simple covalent
b) Giant covalent
Simple covalent structures
a) # atoms
b) Melting point
c) Conductivity
a) Few atoms held together by covalent bonds
b) Low boiling points due to weak intermolecular forces
c) Non conductive due to no free electron or an overall charge
Giant covalent structures
a) # atoms
b) Melting point
c) Conductivity
a) A lot of non-metals held by covalent bonds
b) Very high melting points due to strong covalent bonds
c) Variation in conductivity since some contain free electrons
Allotropes of carbon
a) Diamond
b) Graphite
c) C60 Fullerene
d) Graphene
Structure of Graphite
a) Covalently bonded in a trigonal planar arrangement to three other C atoms to form hexagonal layers
b) London forces between the layers, allowing the bonds to slide over each other easily
c) Covalent layer lattice
Structure of diamond
a) Covalently bonded to four others in a tetrahedral arrangement
b) Giant covalent structure
Structure of Fullerene C60
a) Each C atom is covalently bonded to 3 others in a sphere of 50 C atoms, consisting of 12 pentagons and 20 hexagons.
b) Molecular structure
Structure of Graphene
a) Each C atom is covalently bonded to 3 others forming hexagons in a trigonal planar arrangement.
b) It exists as a single layer
Electrical conductivity of Graphite
a) Good electrical conductor
b) Delocalized e- can move through the layers
Electrical conductivity of Diamond
a) No conductor of electricity
b) All electrons are bonded and so non-mobile
Electrical conductivity of Fullerene C60
Semiconductor
Electrical conductivity of Graphene
a) Very good electrical conductor
b) Delocalized e- can move through the layers
Thermal conductivity of Graphite
Not a good conductor of heat
Thermal conductivity of Diamond
Very efficient thermal conductor
Thermal conductivity of Fullerene C60
Very low thermal conductivity
Thermal conductivity of Graphene
Best thermal conductivity
Appearance of Graphite
Not lustrous, grey crystalline solid
Appearance of Diamond
Highly transparent
Appearance of Fullerene C60
Yellow crystalline solid
Appearance of Graphene
Transparent
Special properties of Graphite
a) Soft and slippery due to slippage of layers
b) High melting point
Special properties of Diamond
a) Hardest
b) High melting point
Special properties of Fullerene C60
a) Very light and strong
b) Low melting point
Special properties of Graphene
a) Thinnest and strongest material ever to exist
b) High melting point
Silicon giant covalent structure
Si forms 4 covalent bonds with other Si atoms to form a giant lattice structure based on a tetrahedral arrangement
SiO2 giant covalent structure
Si forms 4 covalent bonds with other O atoms to form a giant lattice structure based on a tetrahedral arrangement
a) Strong
b) Insoluble in H2O
c) High melting point
d) Nonconductor of electricity
