Chemical Bonding notes

MDCAT Chemistry

Chemical bonding explains how atoms combine through transfer, sharing or delocalisation of electrons. These notes cover VSEPR theory, sigma and pi bonds, hybridization, molecular shape, polarity, dipole moment, bond energy, bond length and hydrogen bonding.

Basic Ideas of Chemical Bonding

Atoms form chemical bonds to attain greater stability. Many atoms achieve a noble gas configuration by completing their valence shell. This is commonly explained by the octet rule, although some elements show exceptions.

Electropositive elements, especially Group 1A elements, lose electrons easily and form positive ions. For example, aluminium has three valence electrons and commonly forms Al3+, so it is called a trivalent element. Electronegativity, ionization energy and electron affinity help explain the tendency of atoms to lose or gain electrons.

Elements of Group 1A are good electron losers because they have one valence electron and low ionization energy.
Aluminium has three valence electrons and forms the trivalent ion Al3+.
A low ionization energy indicates that an atom can lose an electron easily, but it does not prove that the element belongs to the p-block.
Electron affinity is the energy change when an isolated gaseous atom accepts an electron. Its unit is kJ mol-1.
When energy is released during electron addition, electron affinity is assigned a negative sign.
The octet rule has exceptions. Phosphorus can form compounds in which its valence shell contains more than eight electrons.

VSEPR Theory and Molecular Shapes

VSEPR stands for Valence Shell Electron Pair Repulsion theory. It states that electron pairs around a central atom repel one another and arrange themselves as far apart as possible. The shape of a molecule depends on the number of bond pairs and lone pairs around the central atom.

Lone pair repulsion is stronger than bond pair repulsion because a lone pair is attracted by only one nucleus and occupies more space. The order of repulsion is lone pair-lone pair greater than lone pair-bond pair greater than bond pair-bond pair.

The general formula ABn describes a central atom A bonded to n surrounding atoms B.
Methane, CH4, is an AB4 molecule with four bond pairs and no lone pair. Its geometry is tetrahedral.
Molecules with four bond pairs and no lone pair have a tetrahedral geometry with a bond angle of about 109.5 degrees.
NH2- has two bond pairs and two lone pairs around nitrogen. Its molecular shape is angular or bent, although its electron-pair geometry is tetrahedral.
Methyl chloride, CH3Cl, has four bond pairs around carbon and no lone pair. Its geometry is tetrahedral.
VSEPR predicts molecular shape by considering both shared electron pairs and unshared electron pairs.
A double or triple bond is counted as one electron domain in VSEPR, although it repels slightly more strongly than a single bond.

Hybridization and Overlap

Hybridization is the mixing of atomic orbitals of comparable energy on the same atom to form equivalent hybrid orbitals. Hybrid orbitals have different shapes and geometry from the parent atomic orbitals. Their number is equal to the number of atomic orbitals mixed.

According to valence bond theory, a covalent bond forms when half-filled orbitals of two atoms overlap. Greater effective overlap generally produces a stronger bond. The overlapping orbitals must have suitable orientation and comparable energy.

The essential condition for covalent bond formation according to valence bond theory is overlap of participating atomic orbitals.
sp hybridization produces two sp hybrid orbitals and gives a linear arrangement with an angle of 180 degrees.
sp2 hybridization produces three sp2 hybrid orbitals and gives a trigonal planar arrangement with angles of about 120 degrees.
sp3 hybridization produces four sp3 hybrid orbitals and gives a tetrahedral arrangement with angles of about 109.5 degrees.
In methane, carbon is sp3 hybridized. Each C-H bond is formed by overlap of a carbon sp3 orbital with a hydrogen 1s orbital, written as sp3-s overlap.
In HF, the covalent bond is formed by overlap of the hydrogen 1s orbital with a suitable fluorine p orbital, commonly represented as s-pz overlap.
Hybrid orbitals are directional, so they form bonds in definite geometrical directions.

Sigma and Pi Bonds

A sigma bond is formed by head-on or end-to-end overlap along the internuclear axis. It has cylindrical symmetry around the bond axis. A pi bond is formed by sideways overlap of parallel p orbitals, with electron density above and below the internuclear axis.

A single covalent bond always contains one sigma bond. Multiple bonds contain one sigma bond and one or more pi bonds. The sigma bond is formed first, and pi bonding occurs after the atoms are held together by the sigma bond.

Single bond: one sigma bond, for example the C-H bond in methane.
Double bond: one sigma bond and one pi bond, for example the C=C bond in ethene.
Triple bond: one sigma bond and two pi bonds, for example the C≡C bond in ethyne.
Sigma bonds are formed by s-s, s-p, p-p or hybrid orbital overlap along the internuclear axis.
Pi bonds are formed by lateral overlap of parallel unhybridized p orbitals.
Sigma bonds generally allow rotation around the bond axis, whereas pi bonds restrict rotation because sideways overlap must be maintained.
A sigma bond is usually stronger than an individual pi bond because head-on overlap is more effective than sideways overlap.

Dipole Moment and Polarity

A covalent bond is polar when the bonded atoms have different electronegativities. The more electronegative atom develops a partial negative charge, written δ-, while the other atom develops a partial positive charge, written δ+. Bond polarity is represented by a dipole moment.

The polarity of a molecule depends on both bond dipoles and molecular shape. Equal bond dipoles can cancel in a symmetrical molecule, making the whole molecule non-polar. If the dipoles do not cancel, the molecule is polar.

Dipole moment is represented by μ and is related to charge separation and distance: μ = q × r.
The usual unit of dipole moment is Debye, represented by D.
Carbon disulphide, CS2, has a linear shape. Its two C-S bond dipoles cancel, so the molecule is non-polar.
CO2 is also linear and non-polar because its two equal C=O bond dipoles act in opposite directions.
H2O is bent, so its O-H bond dipoles do not cancel. Water is polar.
BF3 is trigonal planar and symmetrical, so its bond dipoles cancel and the molecule is non-polar.
CH3Cl is tetrahedral but not symmetrical because chlorine differs from hydrogen. Its bond dipoles do not cancel completely, so it is polar.
A polar bond does not always make the whole molecule polar. Molecular geometry must also be considered.

Bond Length and Bond Energy

Bond length is the average distance between the nuclei of two atoms joined by a covalent bond. It depends on atomic size, bond order and hybridization. A shorter bond generally has greater electron density between the nuclei and is stronger.

Bond energy, also called bond dissociation energy in suitable contexts, is the energy required to break one mole of a particular bond in gaseous molecules. Bond formation releases approximately the same amount of energy that bond breaking absorbs.

Bond length is affected by atomic radii, bond order, hybridization and the nature of the bonded atoms.
As bond order increases, bond length generally decreases: C-C is longer than C=C, and C=C is longer than C≡C.
As bond order increases, bond energy generally increases: a triple bond is stronger than a double bond, and a double bond is stronger than a single bond.
Larger atoms usually form longer bonds because their valence orbitals are farther from the nuclei.
sp hybrid orbitals have more s-character than sp2 and sp3 orbitals. Greater s-character holds bonding electrons closer to the nucleus and can produce shorter, stronger bonds.
Bond breaking is endothermic because energy is absorbed.
Bond formation is exothermic because energy is released.
The sum of the cationic radius and anionic radius in an ionic crystal is equal to the interionic distance, R.

Hydrogen Bonding and Intermolecular Forces

Hydrogen bonding is a relatively strong intermolecular force. It occurs when hydrogen is covalently bonded to a highly electronegative atom, usually fluorine, oxygen or nitrogen, and is attracted to a lone pair on an electronegative atom of another molecule or another part of the same molecule.

Hydrogen bonding strongly affects boiling point, solubility and molecular structure. It is responsible for important biological structures, including the alpha-helix and beta-pleated sheets of proteins. It also helps stabilise the double-helical structure of DNA.

A hydrogen bond is represented by a dotted line, such as O-H···O or N-H···N.
The usual order of hydrogen bond strength is H-F greater than H-O greater than H-N.
Hydrogen bonding may be intermolecular, between separate molecules, or intramolecular, within the same molecule.
Water has a relatively high boiling point because its molecules form extensive hydrogen bonds.
Hydrogen bonds stabilise the alpha-helix and beta-pleated sheet structures of proteins.
Hydrogen bonding contributes to the double-helical structure of DNA.
Hydrogen bonding is stronger than ordinary dipole-dipole attraction but weaker than a covalent or ionic bond.

Ionic, Covalent and Metallic Bonding

In ionic bonding, electrons are transferred from one atom to another. The resulting oppositely charged ions are held by electrostatic attraction. In covalent bonding, atoms share electron pairs. In metallic bonding, positive metal ions are held together by a mobile sea of delocalised electrons.

The reaction X(s) + 2HCl(aq) → XCl2(aq) + H2(g) indicates that X forms X2+ ions in XCl2. If X is a metal, metallic bonding is present in elemental X, while XCl2 contains ionic bonding between X2+ and Cl- ions.

Ionic compounds usually have high melting and boiling points because of strong electrostatic forces in their lattices.
Covalent compounds generally have low melting and boiling points when they consist of separate molecules.
Many covalent compounds are poor conductors of electricity because they do not contain freely moving ions or electrons.
The statement that all covalent compounds are more reactive than ionic compounds is not a general property of covalent compounds.
Metallic substances conduct electricity because their electrons are delocalised and mobile.
Ionic bonding commonly occurs between metals and non-metals, while covalent bonding commonly occurs between non-metals.
A compound may contain more than one type of bonding. For example, a polyatomic ion can have covalent bonds internally and ionic attraction with another ion.

Key terms

Chemical bond: A force of attraction that holds atoms or ions together in a substance.
VSEPR theory: A theory that explains molecular shape by the repulsion between electron pairs around a central atom.
Bond pair: A pair of electrons shared between two bonded atoms.
Lone pair: A pair of valence electrons not involved in bonding.
Hybridization: The mixing of atomic orbitals to form new equivalent hybrid orbitals.
Sigma bond: A covalent bond formed by head-on overlap along the internuclear axis.
Pi bond: A covalent bond formed by sideways overlap of parallel p orbitals.
Dipole moment: A measure of charge separation in a bond or molecule, represented by μ.
Bond length: The average distance between the nuclei of two bonded atoms.
Bond energy: The energy required to break one mole of a particular bond in gaseous molecules.
Electronegativity: The tendency of an atom in a molecule to attract the shared electron pair towards itself.
Electron affinity: The energy change when an isolated gaseous atom accepts an electron.
Hydrogen bonding: Attraction involving hydrogen bonded to fluorine, oxygen or nitrogen and a lone pair on an electronegative atom.
Interionic distance: The distance between the centres of a cation and an anion in an ionic lattice.

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Chemistry shortcuts

Finding the limiting reactant and percentage composition

Convert every given mass or volume into moles first. The reactant that produces the least amount of the required product is the limiting reactant.

Write the balanced equation and calculate moles using n = mass/Mr.
Use the mole ratio to calculate the product. For percentage composition, use percentage = mass of element in one mole of compound divided by molar mass, multiplied by 100.
Example: Percentage of nitrogen in KNO3 = 14/101 × 100 = 13.86%.
Answer: 13.86% nitrogen.

Use gas volume at molar volume only when the gas conditions are stated or are standard conditions.

Using gas volume, pressure and temperature relations

At the same temperature and pressure, gas volume is directly proportional to the number of molecules. For changing conditions, use P1V1/T1 = P2V2/T2.

At constant temperature and pressure, divide or multiply the volume in the same ratio as the number of molecules.
Example: 10 mL H2 contains 2 × 10^3 molecules. Oxygen in 200 mL contains 20 × 2 × 10^3 = 4 × 10^4 molecules.
Answer: 4 × 10^4 molecules.
For a rigid container, increasing temperature increases molecular speed and mean free path if the gas remains in the same phase.

The direct volume to molecule ratio does not apply when temperature or pressure changes.

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