TOPIC 3: CHEMICAL BONDING Flashcards
[Definition] Covalent Bond
Covalent bond is the electrostatic forces of attraction between the shared electrons and the positively charged nuclei
What particles form covalent bond?
Non-metals,
BUT can be between metal & non-metal also
e.g. AlCl₃
Which part of a molecule has the highest electron density?
The covalent bond (region between the two nuclei)
Sigma (σ) bond VS Pi (π) bond
Head-on overlap of an orbital of each atom VS
sideways overlap of a p orbital of each atom
1 region of electron density VS
2 regions of electron density
More effective overlap (stronger) VS
less effective overlap (weaker)
π bond can only form after σ bond is formed
Corresponding multiple bonding to σ & π bond
single bond: 1 σ bond
double bond: 1 σ bond + 1 π bond
triple bond: 1 σ bond + 2 π bonds
Strength of dative bond
Same strength as a normal covalent bond
3 main types of dative / co-ordinate bonding
- just donate electrons ONLY
* e.g. NH₃BF₃ - donate electrons, then break original bond
* e.g.NH₃ + HCl –> NH₄Cl - share & donate electrons
* e.g. CO
[Definition] Electronegativity
Electronegativity refers to the ability of an atom to attract the shared pair(s) of electrons in a covalent bond.
Trend of electronegativity
Across period: Increases
Down group: Decreases
6 Terms involving bond polarity
- Difference in electronegativity
- Permanent separation of partial charges,
- resulting in formation of dipoles (atoms with partial charge)
- Partial positive / negative charge
- Dipole moment (strength of dipole)
- Polar covalent bond formed
C and H generally considered to be non-polar
[Definition] Bond Energy
Bond energy is the amount of energy required to break one mole of covalent bonds in the gaseous state.
[Factor] Factors affecting covalent bond strength
In order of importance,
1. Bond order
* Bond order increase,
* number of bonding electrons within the inter-nuclei region increases,
* electrostatic forces of attraction for these electrons increases
2. Effectiveness of orbital overlap
* large orbitals are more diffuse,
* less effective orbital overlap,
* weaker bond strength
3. Bond polarity
* more polar the bond, the stronger the bond
Why is the bond energy of F-F lower than Cl-Cl when it should be higher?
F atom very small, so bond length very short.
Non-bonding electrons on the F atom being in very close proximity with those of the other F atom
Repulsion occurs, which weakens the covalent bond
[SOP] Drawing dot-and-cross diagrams for covalent bonds
- Least electronegative atom as central atom (except H)
- Form bonds to achieve noble gas electronic configuartion
- Pair up remaining valence electrons to form lone pairs
*For anion, add electron(s) to the more electronegative atom
*For cation, remove electron(s) from the less electronegative atom
Octet rule still applicable?
Yes, but only for period 2 elements
Elements from period 3 onwards can expand octet due to their vacant and energetically accessible d orbitals
What is covalency?
The number of bonds that a molecule can form
E.g. covalency of oxygen is 2, covalency of carbon is 4
But elements in groups 15-17 in the third and subsequent rows can display more than one covalency
E.g. covalency of phosphorus is 3 or 5, chlorine is 1, 3 or 5
Types of ‘special’ molecules
- Electron deficient molecules (usually grp 2 and 13)
* E.g. BeCl₂, BH₃, AlCl₃
* But can still achieve noble gas ec through forming dative bonds / forming dimer - Radicals (substances with unpaired valence electrons)
* E.g. NO, NO₂
* Central atom may have <8 electrons
* These substances are usually very reactive
* Can still achieve noble gas ec through forming dimer - Expansion of octet structures
* E.g. PCl₅, SO₄²⁻
* Elements from period 3 onwards have vacant and energetically accessible d orbitals to expand beyond octet structure - Dative Covalent Bonds
* NO₂, H₃O⁺
Principle 1 of VSEPR Theory
Valence Shell Electron Pair Repulsion Theory
Electron pairs (both bond pairs and lone pairs) around a central atom are arranged as far apart as possible so as to minimise repulsion.
*For the purpose of VSEPR theory, single/double/triple bonds are assumed to occupy 1 region of electron density ONLY even though they may have multiple σ & π bonds
Principle 2 of VSEPR Theory
Valence Shell Electron Pair Repulsion Theory
lp-lp repulsion > lp-bp repulsion > bp-bp repulsion
Shape & bond angle of 2 electron pairs
2 electron pairs = 2 electron densities in case of VSEPR
2 BP, 0 LP: Linear, 180°
Shape & bond angle of 3 electron pairs
3 electron pairs = 3 electron densities in case of VSEPR
3 BP, 0 LP: Trigonal planar, 120°
2BP, 1 LP: Bent, <120° (e.g. 119°)
Shape & bond angle of 4 electron pairs
4 electron pairs = 4 electron densities in case of VSEPR
4 BP, 0 LP: Tetrahedral, 109.5°
3 BP, 1 LP: Trigonal pyramidal, 107°
2 BP, 2 LP: Bent, 104.5°
Shape & bond angle of 5 electron pairs
5 electron pairs = 5 electron densities in case of VSEPR
5 BP, 0 LP: Trigonal bipyramidal, 120° (equatorial), 90° (axial)
4 BP, 1 LP: See-saw, <120° (equatorial), <90° (axial)
3 BP, 2 LP: T-shaped, <90°
2 BP, 3 LP: Linear, 180°
Shape and bond angle of 6 electron pairs
6 electron pairs = 6 electron densities in case of VSEPR
6 BP, 0 LP: Octahedral, 90° (equatorial & axial)
5 BP, 1 LP: Square pyramidal, <90° (equatorial & axial)
4 BP, 2 LP: Square planar, 90°
What else can determine bond angle other than number of bond pairs and lone pairs?
Electronegativity of central atom or terminal atoms
If central atom has a lower electronegativity / terminal atom has a higher electronegativity, electron density around the central atom decreases, so the repulsion between the neighbouring electron pairs (e.g. bonding pairs) is decreased
Does polar bond mean polar molecule?
No.
Need consider shape, electronegativity, and net dipole moment
What are the molecule shapes that result in cancellation of polar bonds, resulting in non-polar molecules?
ONLY if all terminal atoms are the same:
x bond pair, 0 lone pair
2 bond pair, 3 lone pair: linear
4 bond pair, 2 lone pair: square planar
3 main categories of intermolecular forces of attraction
instantaneous dipole-induced dipole interactions
permanent dipole-permanent dipole interactions
hydrogen bonding
Main ‘characterisitic’ about intermolecular forces of attraction
All are transient (id-id, pd-pd, H-bond)
meaning that the bonds are continuously being formed –> broken –> formed –> broken …
How does instantaneous dipole-induced dipole interactions form?
Temporary fluctuations in the electron distributions within atoms result in the formation of short-lived instantaneous dipole
Instantaneous dipole of the atom can interact with the electrons in an adjacent atom, pulling them towards the positive end of the instantaneous dipole or repelling them from the negative end, forming an induced dipole.
An instant later, the first atom might change its dipole through another movement of electrons within it, so the instantaneous dipole-induced dipole attraction might be destroyed
[Factor] Factors affecting strength of instantaneous dipole-induced dipole interactions
- Number of electrons
* More electrons, bigger and more polarisable electron cloud - Shape of molecule
* Elongated molecules, greater surface area of contact
How does permanent dipole-permanent dipole interactions form?
Uneven distribution of electronic charge within polar molecules result in permanent dipole
Polar molecules always experience both attractive and repulsive dipole-dipole interactions simultaneously
But attractive interactions dominate, hence pd-pd formed
[Factor] Factors affecting strength of permanent dipole-permanent dipole interactions?
Dipole moment
[Definition] Hydrogen Bond
Hydrogen bond is defined as the electrostatic forces of attraction between the partial positive hydrogen atom that is directly bonded to a very electronegative F, O or N atom and a lone pair of electrons of another F, O or N atom.
MUST show when illustrating hydrogen bond
- The hydrogen bond (represented by the dotted lines)
- Labelled hydrogen bond
- The lone pair of electrons
- The partial charges of at least 4 atoms - BOTH atoms involved + the neighbouring atoms which gave the 2 atoms the partial charges
[Factor] Factors affecting strength of hydrogen bond?
- No. of hydrogen bonds (per molecule)
* Smaller value (no. of lone pairs & no. of H atoms) - Electronegativity of the electronegative atom bonded to the H atom
id-id vs pd-pd vs H-bond
For moleules with similar number of electrons ONLY (ie same period)
H-bond > pd-pd > id-id
id-id interactions exists in ALL molecules
[SOP] Comparing compounds with different strengths of intermolecular forces of attraction
- State both are simple molecular structures
- Conditions for the difference in bond
- Difference in bond
- Thus …
E.g. Both compounds are simple molecular structure.
Compound A is non-polar while compound B is polar.
Compound A has id-id while compound B has pd-pd.
So, B higher boiling point / melting point etc….
[Property] Properties of simple molecules
- Low melting/boiling point
* Small amount of energy is needed to overcome the weak id-id / pd-pd / H-bond - Non-conductors of electricity in ANY state
* Electrons are localised in the covalent bonds and are not mobile to conduct electricity - Generally soluble in non-polar solvents & insoluble in polar solvents
* Standard MUST KNOW phrasing:
The energy released in forming _ bonds between the X and Y molecules is sufficient to overcome the _ bond between the X molecules and the _ bond between the Y molecules. Thus, X is soluble in Y.
5 Influences of Hydrogen Bond on the physical properties of molecules
- Unusually high boiling points
* With F/O/N molecules e.g. HF, H₂O, NH₃ - High solubility in water
* For ammonia, small alcohols, organic acids, amines - Anomalous Mr of organic a cids
* Each molecule pairing up to form a cyclic dimer via H bond - Different physical properties of isomers
- Higher density of water compared to ice
Why hydrogen bond result in different physical properties of isomers?
Different shape of isomers result in different types of hydrogen bond
* Intermolecular H bond VS Intramolecular H bond
Intermolecular H bond: More extensive intermolecular H bond, more energy needed to overcome, higher bp, more soluble in water
Intramolecular H bond: Less extensive intermolecular H bond, less energy needed to overcome, lower bp, less soluble in water
Why hydrogen bond result in higher density of water compared to ice?
Hydrogen bond holds ice molecules in an orderly open structure with large amount of empty spaces in between.
In this extended network, each water molecule forms 2 H bond with its neighbouring water molecules and each oxygen atom is surrounded by 4 hydrogen atoms.
[Property] Properties of Giant Molecular Structures
- High melting / boiling point
* Strong covalent bonds between the carbon atoms in a giant molecular structure - Non-conductors of electricity in ALL states
* Electrons are localised in covalent bonds and not mobile to conduct electricity - Insoluble in both water and non-polar solvents
* No favourable solute-solvent interactions between solute and solvent molecules can be formed to break down the giant molecular structure - Hard, strong and non-malleable
* Atoms are held closely by strong covalent bonds in a giant molecular structure
Structure of graphite
Giant molecular structure consisting of atoms held together by strong covalent bonds in a giant extensive planar layer with the adjacent layers held together by weak instantaneous dipole-induced dipole interactions.
To take note when drawing structure of graphite
At least 2 layers with 3 hexagonal rings each
[Property] Properties of graphite
- High melting / boiling point
* Strong covalent bonds between carbon atoms in the giant molecular structure - Good conductors of electricity parallel to the layers but non-conductors of electricity perpendicular to the layers
* 3 of the valence e- are used to form covalent bonds
* Last unused e- is delocalised parallel to the whole layer
* So, presence of delocalised e- which act as mobile charge carriers to conduct electricity along plane - Insoluble in both water and non-polar solvents
* No favourable solute-solvent interactions between solute and solvent molecules can be formed to break down the giant molecular structure - Soft and slippery
* Adjacent layers held together by weak id-id interactions
* Layers can easily slide over one another
[Definition] Ionic Bond
Ionic bonds are strong electrostatic forces of attraction between oppositely charged ions in a giant ionic crystal lattice structure.
[Factor] Factors affecting ionic bond strength
Magnitude of lattice energy
ΔHₗₐₜₜ ∝ |(q⁺ x q⁻) / (r⁺ + r⁻)|
[Definition] Lattice Energy
Lattice energy is the energy released when one mole of the solid ionic compound is formed from its constituent gaseous ions.
Is ionic bond directional?
No. The electrostatic forces of attraction exists between ALL ions of opposite charges. So, ionic bond is non-directional.
[Property] Properties of Ionic Compounds
- High melting / boiling point
* Strong electrostatic forces of attraction between oppositely charged ions - Non electrical conductors in solid state
* Ions can only vibrate about their fixed positions and are not mobile to conduct electricity - Good electrical conductors in molten / aqueous state
* Presence of ions as mobile charge carriers to conduct electricity - Generally soluble in polar solvents (e.g. water)
* Energy released in forming the ion-dipole interactions between ions and water molecules is sufficient to overcome the strong electrostatic forces of attraction between the oppositely charged ions in the giant ionic crystal lattice structure - Insoluble in non-polar solvents
* No favourable ion-dipole interactions between ions and non-polar molecules to break down the giant ionic lattice structure - Strong, hard but brittle
* Hard: Strong EFA between OCI in the GILC
* Brittle: Stress applied results in sliding of ions, resulting in repulsion between ions of similar charges when they go close together, shattering the giant ionic lattice structure
Why is there intermmediate bond character?
Because some covalent bonds have ionic character / ionic bonds have covalent character
How does ionic character in covalent bond occur? / What is it basically?
Polar covalent bond.
When the atoms have a considerable electronegativity difference
Permanent separation of partial charges, each have a partial charge, formation of dipoles and dipole moment, forming polar covalent bond.
[Factor] Extent of ionic character in covalent bond
Difference in electronegativity between the 2 atoms
How is covalent character formed in an ionic bond?
Covalent character arises due to polarisation, which is the distortion of anion electron cloud by a neighbouring cation in an ionic compound.
This results in the electron density lying in the middle of the 2 atoms, resulting in some form of “electron sharing”
[Factor] Extent of covalent character in an ionic bond
- Polarising Power of Cation (depends on charge density)
* Small, highly charged ions have greater charge densities and thus, higher polarising power - Polarisability of Anion (depends on anionic radius)
* Larger the anionic radius, more polarisable it is
Summary, largest extent of covalent character is when both polarising power of cation and polarisability of anion is HUGE
* Higher cationic charge, Small cationic radius, Large anionic radius
Boiling temp of ionic compounds
USUALLY exceed 700°C
[Definition] Metallic Bond
Metallic bonds are strong electrostatic forces of attraction between metal cations and the sea of delocalised valence electrons in a giant metallic lattice structure.
Where do the ‘sea of delocalised valence electrons’ come from in a metallic stucture?
The valence electrons are delocalised from the metal atoms due to the weak electrostatic forces of attraction between them and the nucleus
Neutrality of metal?
Electrically neutral - number of positive and negative charges balances out
[Factor] Factors affecting metallic bond strength
- Number of delocalised valence electrons contributed PER ATOM
- Charge density of the metal (q/r) - only use if atom in same group
[Property] Properties of Metals
- High melting / boiling point
* Strong electrostatic forces of attraction between the metal cations and sea of delocalised valence electrons - Good conductors of electricity in solid AND liquid state
* Presence of sea of delocalised electrons to act as mobile charge carries - Good conductors of heat
* Sea of delocalised electrons transfer heat energy quickly - Malleable (beaten into sheets) and Ductile (pulled into wires)
* Stress applied causes sliding of layers of cations
* WITHOUT breaking down the metallic stucture because the sea of delocalised valence electrons are holding the cations together to prevent repulsion between the cations - Alloy formation
* Presence of cations (metals) / atoms (non-metals) of different sizes in the metallic lattice structure inhibits sliding of the layers, increasing tensile strength of alloy
