Covalent bonds

covalent bonds

• electrostatic attraction between a shared pair of electrons and the positively charged nuclei

• electrons are shared rather than transferred

• 2 atomic orbitals overlap and molecular orbital is formed

• electrons are more stable when attracted to 2 nuclei rather than one

• electron pairs should be regarded as charge clouds, as they are in a state of constant motion

octet rule

• tendency of atoms to gain a valence shell with a total of 8 electrons

• being able to accommodate more than 8 electrons in the outer shell - expanding the octet

• accommodating less than 8 electrons - electron deficient

steps for drawing lewis structures

1. count total number of valence

2. draw skeletal structure

3. use dots/crosses to show an electron pair

4. add more electron pairs to complete the octets around the atoms

5. if not enough electrons to complete octet, add double/triple bonds

incomplete octet examples

BeCl2, BF3

• Be and B form stable compounds with 4 and 6 valence electrons respectively

bond energy

• energy required to break one mole of a particular covalent bond in the gaseous states, the larger the bond energy, the stronger the covalent bomd

bond length

• internuclear distance of 2 covalently bonded atoms

• the greater the forces of attraction are, the more the 2 atoms are pulled to eachother

• this decreases the bond length of a molecule and increases the strength of a covalent bond

• triple bonds are the shortest and strongest due to large electron density between the nuclei

coordinate bonds

• some molecules have a lone pair of electrons that can be donated to form a bond with an electron deficient atom

• both electrons share the same atom

• this is called coordinate bonding, use an arrow to indicate it

• e.g ammonium ion donating to H+ ion

VSEPR theory

• when an atom forms a covalent bond with another atom, the electrons in the different bonds and the non-bonding electrons in the outer shell all behave as negatively charged clouds, repelling eachother

• regions of negative cloud charge are known as domains,

• in order to minimise this repulsion, all outer shell electrons are spread out as far as possible in the space

• molecular shapes and their angles can be predicted through this theory

3 rules of VSEPR theory

1. all electron pair and all lone pairs arrange themselves as far apart in the space as possible

2. lone pairs repel more strongly than bonding pairs (because they’re pulled more closely to the central atom)

3. multiple bonds behave like single bonds

2 electron domains

• shape: linear

• bond angle: 180

3 electron domains

• shape: trigonal planar

• bond angle: 120

• if one of these domains is a lone pair, bond angle is slightly less than 120 (118) due to the stronger repulsion from lone pairs, forcing the bonding pairs closer together

• this shape is now called: bent linear

4 electron domains

• shape: tetrahedral

• bond angle: 109.5

• if one of the electron domains is a lone pair, the bond angle is slightly less than 109.5 (107)

• this shape is now called: trigonal pyramidal

• if two of these electron domains is a lone pair, the bond angle is slightly less than 109 (104.5)

• this shape is now called: bent linear

electronegativity

• the ability of an atom to draw electrons towards itself in a covalent bond

• in diatomic molecules, electron density is shared between the 2 atoms

• both atoms have the electronegativity value, and have an equal attraction for the bonding pair of electrons, this forms a covalent bond

polar bond

• when 2 atoms in a covalent bond have 2 different electronegativities, the covalent bond is polar and the electron is drawn to the more electronegative side

• electron distribution is asymmetric

• extent of polarity in a covalent bond varies, depending on how big a difference in electronegativity values between the 2 atoms

  • the higher the difference, the higher the polarity

dipole moment

• measure of how polar a bond is

• direction of the dipole movement is shown by an arrow, pointing towards the partially negative charged end of the pole

molecular polarity

• to determine whether a molecule with more than 2 atoms is polar consider:

  • polarity of each bond

  • geometry of the molecule

• some molecules have polar bonds, but are overall non-polar because the molecule is arranged in such a way that the dipole moments cancel eachother out

covalent lattices

• in some cases it is impossible to satisfy the bonding capacity of a substance in the form of a molecule, the bonds between atoms continue indefinitely and a large lattice is formed

• these are called giant covalent structures, most important examples are C and SiO2

allotrope

• different molecular arrangements of the same element in the same physical state

diamond

• each carbon to 4 others tetrahedral arrangement with a bond angle of 109.5

• results in giant lattice with strong bonds in all directions

graphite

• each carbon atom bonded to 3 others in a layered structure

• layers made of hexagons with bond angle of 120

• spare electron is delocalised and occupies space between layers

• all atoms in the same layer held together by strong covalent bonds, different layers held together by weak IM forces

buckminsterfullerene

• contains 60 carbon atoms, each is bonded to 3 others by covalent bonds

• 4th electron is delocalised so electrons can migrate throughout the structure, making it a semi conductor

graphene

• made of single layer of carbons bonded together in a repeating pattern of hexagons

silicon

• each silicon atom bonded to 4 others, tetrahedral arrangement

silicon (IV) oxide

• each silicon bonded to 4 oxygen atoms, each oxygen is shared by 2 silicon atoms tetrahedral arrangement

• tetrahedral units are repeated

charateristics of giant covalent structures

intermolecular forces

• forces of attraction between molecules in molecular covalent compounds

• three main types:

  • london forces

  • dipole dipole attraction

  • hydrogen bonding

london dispersion forces

• electrons are in a state of constant motion, therefore it is likely they are not always distributed symmetrically

• this is known as a temporary dipole

  • constantly appearing and disappearing as electrons keep moving

• an adjacent atom will be repelled and attracted by the dipole and will move accordingly

  • this is a temporary induced dipole

• london forces are present between all molecules but are generally very weak (weakest out of IM bonds)

• happens between non-polar species

• strength of the forces depends on:

  • number of electrons in molecule

  • surface area of molecule

number of electrons (london forces)

• the more electrons in a molecule, the greater possibility of a distortion, thus the greater frequency and magnitude of temporary dipoles

  • london forces are stronger and mpt/bpt is higher

surface area (london forces)

• the larger the SA, the larger contact with adjacent molecules

• the greater the ability to induce a dipole in an adjacent molecule, the greater the london forces and the higher the mpt/bpt

dipole-dipole attractions

• temporary dipoles exist in all molecules, in some molecules there is also permanent dipole

• in addition to london forces caused by temporary dipoles, molecules with permanent dipoles are also attracted to eachother

• this type of bonding slightly increases bpt than would be expected with just london forces, as well as strength of IM bonds

• second weakest out of IM bonds

dipole induced dipole attraction

• some mixtures contain both polar and non polar molecules

• the permanent dipole of the polar molecule can cause a temporary separation of charge in a non-polar molecule

• this force acts in addition to london forces between non polar molecules, and dipole dipole attraction in polar molecules

hydrogen bonding

• strongest type of IM force

• for hydrogen bonding to take place:

  • a species which has an O, N , or F (very electronegative) atom, with an available lone pair

  • a hydrogen attached to the O, N or F

• when this happens, the bond becomes very polarised

• the H becomes so delta positive that it can form a bond with a lone pair of O, N, F in another molecule

• hydrogen bonds represented by dashes

van der Waals forces

• used to refer to

  • london dispersion forces

  • dipole dipole attraction

  • dipole induced dipole attraction

  • Hydrogen bonding

melting and boiling point explained in covalent

• when covalent molecular substances change state, you are overcoming IM forces

• the stronger the forces, the more energy needed to break

• IM forces are much weaker than covalent bonds, this is why many covalent substances are liquid/gas at rtp

• volatility is high when mpt/bpt is low

solubility explained in covalent

• non-polar dissolves in non-polar solvents, london forces form between solvent and solute

• polar dissolves in polar solvents as a result of dipole-dipole attraction, or hydrogen bonding between solvent and solute

• as covalent molecules become larger, solubility can decrease as polar part of molecule is a smaller part of overall structure

• this is why giant covalent structures are generally insoluble

conductivity explained in covalent

• as covalent substances contain no mobile charged particles, they’re unable to conduct in solid or liquid state

• some giant covalent structures can conduct due to delocalised electrons, but are exceptions

Rf values in chromatography

• the extent of separation of the component molecules in the investigated sample depends on their solubility in the mobile phase and the extent of adhesion to the stationary phase

• Rf values are used to quantify the distance travelled relative to solvent front

resonance structures

• delocalisation of electrons explains why structures of some species don’t fit with lewis formula

  • delocalised electrons are electrons in a molecule that aren’t associated with a single covalent bond or atom

• the lewis structure for nitrate ion has a double bond and 2 single bonds

  • there are 3 ways to structure this molecule

  • these structures are resonance structures

• although, bonds are all equal in length and electron density is spread evenly between the 3 oxygen atoms

• the bond length is somewhere between a single and a double bond, and the structure is somewhere in between the 3, this is called a resonance hybrid

• dotted lines are used to show the delocalised electrons

criteria for resonance structures

• molecules must have a double bond that is capable of migrating from one part of the molecule to another

• lone pairs of electrons that can re-arrange themselves and allow the double bonds to be in different positions

benzene as a resonance structure

• each carbon atom in the ring forms 3 σ bonds using the sp² orbitals

• the remaining p orbitals overlap laterally w/ p orbitals of neighbouring carbons forming a π system

• extensive side ways overlap allows delocalisation of electrons, being able to spread freely around the ring.

• delocalisation of electrons renders the carbon-carbon bonds to have both double and single bond character (resonance).

• evidence includes:

  • bond length being between the single and double value

  • undergoes substitution reactions instead of addition reactions

expansion of the octet

• elements in period 3 and above can have more than 8 electrons in valence shell, because of the d-subshell being able to accommodate more electron pairs

5 electron domains

• 5 BP 0 LP: trigonal bimpyramidal,

• 4 BP 1 LP: see saw

• 3 BP 2 LP: t- shape

• 2 BP 3 LP: linear

6 electron domains

• 6 BP 0 LP: octahedral

• 5 BP 1 LP: square based pyramid

• 4 BP 2 LP: square planar

formal charge

• sometimes it’s difficult to determine which lewis structure is most appropriate for a molecule

• formal charge is the charge assigned to an atom in a molecule, assuming that all the electrons in the bonds are shared equally between atoms, regardless of differences in electronegativity

• FC= (number of valence electrons) - ½(number of bonding electrons) - (number of non-bonding electrons)

• the lewis formula that is preferred is:

  • the difference in FC of the atoms is closest to zero

  • negative charges located on most electronegative atoms

bond overlap

• each atom that combines has an atomic orbital with an unpaired electron

• when 2 atomic orbitals overlap, they form a molecular orbital containing 2 electrons

• the greater the overlap, the stronger the bond

sigma bonds

• formed from head on overlap of atomic orbitals

• electron density is concentrated along the bond axis (imaginary line between the 2 nuclei)

• single covalent bonds are always sigma bonds

• can be s+s, s+p or p+p

pi bonds

• sideways overlap of adjacent p orbitals

• the 2 lobes that makeup the pi bond are above and below the plane of the sigma bond

• the electron density is concentrated on opposite sides of the bond axis

• only found within double and triple bonds

hybridisation

• the electronic structure of the ground state of carbon would imply that it uses the unpaired 2p electrons to form covalent bonds

• this is not true, as carbon usually forms 4 covalent bonds

• a half full p-subshell is slightly lower in energy than a partially filled one

• the difference in energy between the 2s and the 2p subshell is small, so an electron can be promoted fairly easily from the 2s to the 2p

• the 2s and the 2p subshells blend together to form 4 new hybrid orbitals (because 1 s and 3 p)

• this would give 4 unpaired electrons capable of forming 4 covalent bonds

sp3 hybridisation

• 4 hybrid orbitals are produced when the 2s and and three 2p are blended together

• these hybrids are 25% s character and 75% p character

• the 4 sp3 orbitals space themselves out at 109.5 degrees in a tetrahedral shape

• carbon atom forms single bonds

  • sp3 orbitals merge with s orbitals in hydrogen forming 4 equal sigma bonds

• lone pairs can also be present in hybridisation e.g in ammonia

sp2 hybridisation

• 3 hybrid orbitals are produced when the 2s and two 2p are blended together

• these hybrids are 1/3 s character and 2/3 p character

• space themselves out at 120 degrees, in a trigonal planar shape

• carbon atom forms double bond

  • sp2 orbital merges with s orbitals in hydrogen, and the sp2 of an adjacent carbon molecule, forming 3 equal sigma bonds

  • double bond is created by sideways overlap of the unhybridised p orbitals, causing formation of 1 pi bond

sp hybridisation

• 2 hybrid orbitals are produced when the 2s and one 2p are blended together

• these hybrids are ½ s character and ½ p character

• space themselves out at 180 degrees, in a linear shape

• the carbon atom forms a triple bond

  • sp orbital merges with s orbital in hydrogen, and the sp of an adjacent carbon to form 2 equal sigma bonds

  • triple bond is created by head-on overlap of 2 pairs of unhybridised p orbitals (from the 2 carbons)