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2.2 Electrons, Bonding and Structure - Coggle Diagram
2.2 Electrons, Bonding and Structure
Shapes of Molecules and Intermolecular Forces
Shapes of Molecules and Ions
Molecular shape depends on electron pair arrangement
The three-dimensional shape of a molecule or ion depends on the number and arrangement of electron pairs surrounding its central atom. These electron pairs fall into two categories:
Bonding pairs - Involved in forming covalent bonds with other atoms.
Lone pairs - Not involved in bonding and remain on the central atom.
For instance, the ammonia molecule (NH3) has four electron pairs around its central nitrogen atom: three are bonding pairs connected to hydrogen atoms, and one is a lone pair.
Electron pair repulsion theory
A molecule adopts a shape that minimises the repulsion between its electron pairs due to their negative charge:
Electron pairs, whether bonding or lone, repel each other.
Lone pairs cause more repulsion than bonding pairs because they are closer to the nucleus.
Electron pairs arrange themselves as far apart as possible to minimise repulsion.
Repulsion decreases in this order:
Lone pair-lone pair > Lone pair-bonding pair > Bonding pair-bonding pair.
Consequently, lone pairs occupy more space, affecting the molecule's shape by compressing bond angles. This principle underpins the electron pair repulsion theory, which helps predict molecular geometry.
Examples include methane (CH4), ammonia (NH3), and water (H2O), which differ in bond angle and shape despite each having four electron pairs around the central atom due to the varying number of lone pairs.
Shapes of species with multiple bonds
When determining the overall shape of molecules or ions containing multiple bonds, treat each multiple bond as if it were a single electron pair.
Determining the number of electron pairs
To predict the shape of a molecule or ion, first determine the total number of electron pairs (bonding + lone) on the central atom using these steps:
Identify the central atom bonded to all other atoms.
Find the number of outer shell electrons of the central atom using its group number.
Add one electron for each bonded atom.
If the species is an ion, add one electron for each negative charge or subtract one electron for each positive charge.
Divide the total number of electrons by two to get the number of electron pairs.
Subtract the number of bonds from the number of electron pairs to determine the number of lone pairs.
Based on the electron pair arrangement, the molecular shape can be predicted as follows:
2 Electron Pairs
Bonding Pairs- 2
Lone Pairs - 0
Shape- Linear, 180
3 Electron Pairs
Bonding Pairs- 3
Lone Pairs- 0
Shape- Trigonal Planar, 120
4 Electron Pairs
Bonding Pairs- 4
Lone Pairs- 0
Shape- Tetrahedral, 109.5
6 Electron Pairs
Bonding Pairs- 6
Lone Pairs - 0
Shape- Octahedral, 90
4 Electron Pairs
Bonding Pairs- 3
Lone Pairs- 1
Shape- Trigonal Planar, 107
4 Electron Pairs
Bonding Pairs- 2
Lone Pairs- 2
Shape- Bent, 104.5
5 Electron Pairs
Bonding Pairs- 5
Lone Pairs- 5
Shape- Trigonal Bipyramidal, 120 and 90
Electronegativity and Polarity
Electronegativity
Electronegativity is a measure of an atom's ability to attract the shared electron pair in a covalent bond towards itself.
Fluorine is the most electronegative element, followed by oxygen, nitrogen and chlorine.
Electronegativity is measured on the Pauling scale. A higher value on this scale means an atom is more electronegative.
Factors affecting electronegativity
Atomic radius - Smaller atoms are more electronegative because their electrons are closer to the nucleus, resulting in a stronger electrostatic attraction.
Nuclear charge - A higher positive charge in the nucleus increases the strength of the electrostatic attraction between the nucleus and the electrons, making an atom more electronegative.
Shielding - Electrons in inner shells can weaken the electrostatic attraction between the nucleus and the outer shell electrons, reducing an atom's electronegativity.
Trends in electronegativity across the periodic table
Electronegativity increases across a period
This trend occurs because the atomic radius decreases while the nuclear charge increases, resulting in a stronger electrostatic attraction between the nucleus and the electrons in the outer shell.
The effect of an increasing nuclear charge dominates over the decreasing atomic radius, leading to an increase in electronegativity.
Electronegativity decreases down a group
Down a group, although the nuclear charge increases, the atomic radius increases more significantly, and the number of inner shell electrons shielding the outer electrons also increases.
The increased distance between the nucleus and the outer electrons, combined with greater shielding, weakens the electrostatic attraction between them. As a result, electronegativity decreases despite the increase in nuclear charge.
Polar bonds
Differences in electronegativity between bonded atoms can cause polar bonds:
In a covalent bond between atoms with different electronegativities, the bonding electrons are more strongly attracted to the more electronegative atom. This unequal sharing of electrons makes the bond polar.
Polar bonds have a permanent electric dipole. A dipole is a separation of positive and negative charges within a polar covalent bond or polar molecule, resulting from an uneven distribution of electrons.
In a polar bond, the more electronegative atom acquires a slight negative charge (δ-) and the less electronegative atom acquires a slight positive charge (δ+). The bigger the difference in electronegativity, the more polar the bond.
Bonds between atoms of the same element, like H2 and Cl2, are non-polar because the electrons are shared equally.
Bonds between atoms with similar electronegativity, such as carbon and hydrogen, are essentially non-polar due to an even electron distribution.
Charge in polar molecules
The overall polarity of a molecule depends on how its polar bonds are arranged in 3D:
Molecules like CCl4 are non-polar because their symmetric arrangement causes their dipoles to cancel out, even though they have polar bonds. In CCl4, the four polar C-Cl bonds are arranged tetrahedrally, resulting in the dipoles cancelling each other out and no net dipole for the molecule.
Molecules like CHCl3 are polar because their asymmetric arrangement doesn't allow their dipoles to cancel out, leading to a net dipole. In CHCl3, the three polar C-Cl bonds and one C-H bond are arranged tetrahedrally, but the difference in electronegativity between H and Cl causes the dipoles to not fully cancel each other, resulting in a net dipole for the molecule.
How electronegativity differences predict bond type
The type of bond formed between two atoms can be predicted by the difference in their electronegativities:
When the electronegativity difference is zero, as in diatomic molecules like H2 and O2 where the atoms are identical, the bonding electrons are shared equally, resulting in a pure (non-polar) covalent bond.
As the electronegativity difference increases, the bond becomes more polar covalent. For example, in HCl, the chlorine atom attracts the bonding electrons more strongly than the hydrogen atom, leading to a polar covalent bond.
When the electronegativity difference is very large, the bond is considered ionic. In this case, the more electronegative atom effectively takes the bonding electrons from the less electronegative atom. An example is NaCl, where the chlorine atom takes an electron from the sodium atom, forming Na⁺ and Cl⁻ ions.
Most compounds fall somewhere between these extremes, exhibiting a mix of ionic and covalent character depending on the electronegativity difference between the bonded atoms.
Intermolecular Forces
The types of intermolecular force
Intermolecular forces are the attractive forces that exist between molecules. These forces are much weaker than the covalent bonds that hold atoms together within molecules.
There are three main types of intermolecular force, listed in order of increasing strength:
Induced dipole-dipole forces (also known as London dispersion forces or Van der Waals’ forces).
Strength: Weak
Found in: All molecules and noble gases
Permanent dipole-dipole forces.
Strength: Moderate
Found in: Polar molecules
Hydrogen bonding.
Strength: Strong
Found in: Polar molecules with H-F, H-O or H-N bond
Induced dipole-dipole forces
Induced dipole-dipole forces, also known as London dispersion forces or Van der Waals’ forces, are present between all atoms and molecules, even non-polar ones.
How induced dipole-dipole forces arise
Electrons in atoms are constantly moving. At any instant, there may be more electrons on one side of the atom than the other, creating a temporary dipole.
This temporary dipole can induce an opposite dipole in a neighbouring atom, causing a weak electrostatic attraction between the atoms.
This induced dipole can then induce further dipoles in other nearby particles.
Although these dipoles are constantly forming and disappearing as the electrons move, the overall effect is a net attraction between the atoms or molecules.
Factors affecting the strength of induced dipole-dipole forces
The strength of induced dipole-dipole forces increases with the size and surface area of the atoms or molecules:
Size - Larger atoms and molecules have more electrons and a greater volume of electron density that can become polarised, creating stronger temporary dipoles.
Surface area - Molecules with a larger surface area also have stronger induced dipole-dipole forces as more of the electron cloud is exposed for interactions.
Consequently, substances with stronger induced dipole-dipole forces tend to have higher boiling points.
Induced dipole-dipole forces in molecular lattices and noble gases
Induced dipole-dipole forces can also be strong enough to hold molecules together in a lattice structure. For example:
In solid iodine, I2 molecules are held together by strong covalent bonds.
These I2 molecules are then attracted to each other by weak induced dipole-dipole forces, forming a molecular lattice.
Induced dipole-dipole forces also explain the existence of noble gas liquids and solids - even though noble gas atoms have complete outer shells and do not form covalent, ionic or metallic bonds, the weak induced dipole-dipole forces allow them to condense into the liquid and solid states at very low temperatures.
Permanent dipole-dipole forces
Polar molecules have permanent dipoles arising from unequal sharing of electrons in covalent bonds. The partial positive (δ+) and partial negative (δ-) charges on polar molecules enable them to experience permanent dipole-dipole forces.
How permanent dipole-dipole forces arise
Permanent dipole-dipole forces are electrostatic attractions between the partial positive end of one polar molecule and the partial negative end of another.
For example, in gaseous hydrogen chloride (HCl):
The H-Cl bond is polar due to the greater electronegativity of chlorine compared to hydrogen.
The hydrogen atom bears a partial positive charge (δ+) and the chlorine a partial negative charge (δ-).
HCl molecules align so the δ+ hydrogen of one molecule is attracted to the δ- chlorine of a neighbouring molecule.
Polar molecules contain permanent and induced dipole-dipole forces
These permanent dipole-dipole forces are in addition to the induced dipole-dipole forces that exist between all molecules. So polar molecules have stronger overall intermolecular forces than non-polar molecules of similar size.
The difference in boiling points can be explained by the types of intermolecular forces present in each substance:
Methanal is a polar molecule due to the carbonyl group (C=O), so it experiences both permanent dipole-dipole forces and induced dipole-dipole forces. The permanent dipole-dipole forces are stronger than the induced dipole-dipole forces alone.
Ethane is a non-polar molecule, so it only experiences induced dipole-dipole forces.
As a result, more energy is required to overcome the stronger intermolecular forces in methanal, leading to a higher boiling point, even though the molecules are of similar size and have the same molecular mass (Mr = 30).
Hydrogen bonding
Hydrogen bonding is a special type of (permanent) dipole-dipole force that occurs when hydrogen is bonded to the highly electronegative elements fluorine, oxygen or nitrogen.
Requirements for hydrogen bonding
The molecule must contain a hydrogen atom covalently bonded to either fluorine (F), oxygen (O), or nitrogen (N).
There must be a lone pair of electrons on the F, O, or N atom of an adjacent molecule available to interact with the hydrogen.
How hydrogen bonds form
The H-F, H-O, and H-N bonds are highly polar due to the large electronegativity differences between hydrogen and these elements. This leads to a significant partial positive charge (δ+) on the hydrogen atom and a partial negative charge (δ-) on the F, O, or N atom.
The small size of the hydrogen atom allows it to get close to the lone pair of electrons on an adjacent F, O, or N atom.
The lone pairs on F, O, and N atoms are regions of high electron density and therefore high partial negative charge.
The positively charged hydrogen is strongly attracted to the negatively charged lone pair, forming a hydrogen bond between the molecules.
How intermolecular forces affect the properties of substances
Hydrogen bonding has significant effects on the properties of substances:
Greater solubility in water - Substances that can form hydrogen bonds with water (e.g., ethanol) tend to be soluble, while those that cannot (e.g., ethane) are typically insoluble.
Higher melting and boiling points compared to similar-sized molecules that cannot hydrogen bond - Extra energy is needed to overcome the strong hydrogen bonding forces.
Hydrogen bonding explains the anomalous properties of water and ice
Ice is less dense than water - In solid ice, water molecules are arranged in a 3D lattice held together by hydrogen bonds. Upon melting, some of these hydrogen bonds break. Since hydrogen bonds are relatively long compared to covalent bonds, this causes ice to be less dense than liquid water.
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Water and ice have high melting and boiling points - Water has relatively high melting and boiling points compared to other molecules of similar size. This is due to the strong hydrogen bonds between H2O molecules, which require more energy to break
Intermolecular forces explain properties of simple molecular substances. The type and strength of intermolecular forces present in a simple molecular substance influence its physical properties.
Low melting and boiling points - Weak intermolecular forces require little energy to overcome, so covalent compounds often have low melting and boiling points and may be liquid or gaseous at room temperature. Stronger intermolecular forces lead to higher melting and boiling points.
Solubility in water - Polar molecules, especially those capable of hydrogen bonding (e.g., ethanoic acid), can interact favorably with water molecules and are soluble. Non-polar molecules that only have induced dipole-dipole forces (e.g., hexane) are hydrophobic and insoluble in water.
Electrical conductivity - Covalent compounds do not conduct electricity, regardless of polarity. Even though polar molecules have permanent dipoles, they are electrically neutral overall and do not carry charge.
Bonding and Structure
Electron Structure
The arrangement of electrons in shells, sub-shells, and orbitals
In the modern model of the atom, electrons are found in specific energy levels known as shells surrounding the nucleus.
Shells that are further from the nucleus hold electrons with higher energy compared to those closer.
Each shell is defined by a principal quantum number (n = 1, 2, 3...), indicating its relative distance from the nucleus.
These shells are further divided into sub-shells, which have distinct energy levels and are labelled as s, p, d, and f.
Within sub-shells, electrons are located in orbitals, which are regions with a high probability of finding an electron. An orbital is defined as a region around the nucleus that can accommodate up to two electrons with opposite spins.
The capacity of each sub-shell type to hold electrons
s sub-shell
Number of orbitals: 1
Maximum number of electrons: 2
p sub-shell
Number of orbitals: 3
Maximum number of electrons: 6
d sub-shell
Number of orbitals: 5
Maximum number of electrons: 10
f sub-shell
Number of orbitals: 7
Maximum number of electrons: 14
The distribution of sub-shells across the first four shells
Shell 1
Sub-shells: 1s
Total number of electrons: 2
Shell 2
Sub-shells: 2s, 2p
Total number of electrons: 8
Shell 3
Sub-shells: 3s, 3p, 3d
Total number of electrons: 18
Shell 4
Sub-shells: 3s, 3p, 3d
Total number of electrons: 18
The shapes of s and p orbitals
Orbitals, the regions in which electrons are most likely to be found, have unique shapes:
Orbitals within the same sub-shell are of equal energy.
Each orbital can hold two electrons, which must have opposite spins (called spin-pairing).
s orbitals are spherical in shape
p orbitals are dumbbell-shaped. The three p orbitals are oriented at right angles to each other.
Representing electron configurations
The arrangement of electrons in an atom is called its electron configuration. This can be depicted in multiple ways:
Sub- shell notation
This method uses superscripts to indicate the number of electrons within each sub-shell.
For example, the electron configuration of neon (10 electrons) is 1s2 2s2 2p6.
Electrons-in-boxes notation
Orbitals are depicted as boxes, with electrons shown as arrows.
Oppositely directed arrows represent electrons with opposite spins.
Electron pairing within orbitals occurs only with opposite spins.
Electron configurations represent the most stable arrangement
Electron configurations are arranged to minimise the overall energy of the atom or ion.
This lowest energy arrangement corresponds to the most stable electronic structure.
To deduce an atom's electron configuration, follow these guidelines:
Electrons fill the lowest energy orbitals first.
Electrons first occupy orbitals of equal energy singly before pairing up.
When two electrons occupy the same orbital, they must have opposite spins (up and down) to minimise electron-electron repulsion.
For ions in the s and p blocks, electrons are added to or removed from the highest occupied sub-shell
Ionic Bonding and Structure
Ions form to achieve full outer electron shells
Atoms can gain or lose electrons to form ions with full outer shells:
Metal atoms lose electrons to become positive ions called cations. For example, sodium loses one electron to form a sodium ion (Na+).
Sodium atom: 1s2, 2s2, 2p6, 3s1
Sodium ion (Na+): 1s2, 2s2, 2p6
Non-metal atoms gain electrons to become negative ions called anions. For example, chlorine gains one electron to form a chloride ion (Cl-).
Chlorine atom: 1s2, 2s2, 2p6, 3s2, 3p5,
Chloride ion (Cl-): 1s2, 2s2, 2p6, 3s2, 3p6,
In ionic bonding, electrons lost by the metal atom are transferred to the non-metal atom. This transfer allows both atoms to achieve full outer electron shells, resulting in the formation of stable ions.
Compound ions
Compound ions consist of atoms from two or more elements chemically bonded together, resulting in an overall charge.
The important compound ions to know are:
Nitrate ion (NO3-)
Carbonate ion (CO32-)
Sulfate ion (SO42-)
Hydroxide ion (OH-)
Ammonium ion (NH4+)
Ionic bonding occurs between oppositely charged ions
An ionic bond is the electrostatic force of attraction between oppositely charged ions, usually a metal and a non-metal.
These bonds are very strong.
When ions bond this way, an ionic compound is formed.
Structure of ionic compounds
Ionic compounds have giant lattice structures where positive and negative ions pack together.
Key features of a giant ionic lattice:
Each ion is electrostatically attracted to ions of the opposite charge in all directions.
It takes significant energy to overcome these strong electrostatic forces between the ions.
Properties related to ionic structure
The properties of ionic compounds result from their lattice structure:
High melting and boiling points - The strong electrostatic attractions between the positive and negative ions in the giant lattice must be overcome for the lattice to break apart; this requires a lot of energy.
Conduct electricity when molten or in solution - When melted or dissolved, the ions can move freely and carry electric charge through the liquid.
Do not conduct electricity as solids - In the solid lattice structure, the ions are firmly locked in place and unable to move to carry electric charge.
Dissolve in water - Water molecules, which are polar, attract the charged ions in the lattice via ion-dipole forces, pulling them away from the lattice and dissolving the structure.
Covalent Bonding and Simple Molecular Substances
What covalent bonding is
A covalent bond is the strong electrostatic attraction formed between a shared pair of electrons and the nuclei of the bonded atoms.
This bonding creates a stable molecule - a group of two or more atoms held together by covalent bonds.
Examples of covalently bonded molecules include:
Chlorine (Cl2)
Ammonia (NH3)
Water (H2O)
Methane (CH4)
Covalent bond strength
The strength of a covalent bond depends on how much energy is needed to break the bond.
We measure this bond strength using a value called average bond enthalpy, expressed in units of kJ mol-1.
The higher the average bond enthalpy value, the stronger the covalent bond.
For example:
A N≡N triple bond has a very high bond energy of 945 kJ mol-1, indicating it is a very strong bond.
In contrast, a F-F single bond has a lower bond energy of 159 kJ mol-1, indicating it is a weaker bond.
Double and triple covalent bonds
The number of covalent bonds an atom forms depends on how many electrons it needs to fill its outer shell.
Atoms can share multiple electron pairs to fill their outer shell, forming double or triple bonds.
For Example
O≡O, Oxygen has 6 electrons in its outer shell and needs 2 more to complete it, so it forms a double bond by sharing 2 sets of electrons with another oxygen atom.
N≡N, Nitrogen has 5 electrons in its outer shell and needs 3 more to complete it, so it forms a triple bond by sharing 3 sets of electrons with another nitrogen atom.
Dative covalent bonding
In dative covalent bonding, also called coordinate bonding, both shared electrons come from just one of the bonding atoms rather than one electron coming from each atom.
For a dative covalent bond to form between two atoms, the following requirements must be met:
One atom must have a lone pair of electrons to donate.
The other atom must be electron deficient, (i.e., it must have an incomplete electron shell).
This type of bonding is represented by an arrow showing the direction of electron donation from the atom with the lone pair to the electron-deficient atom.
For example, in ammonium NH4+, the nitrogen atom provides both shared electrons to form a dative covalent bond with the hydrogen ion (H+)
Giant Covalent Structures
Giant covalent structures
Some elements can form extensive interconnecting networks of covalently bonded atoms known as giant covalent structures.
These structures involve huge lattices extending in three dimensions.
In carbon, the small atomic size and ability to form 4 covalent bonds per atom allow the formation of giant covalent structures.
The different structural forms of an element in the same state are called allotropes.
The 3 allotropes of carbon with giant lattice structures that you need to know about are diamond, graphite, and graphene.
Diamond - bonding, structure, properties
Bonding and structure:
Each carbon atom forms 4 very strong covalent bonds with others in a tetrahedral arrangement.
Properties:
Extremely hard - Extensive network of strong covalent bonds not easily broken.
Very high melting point - Huge amount of energy needed to break enough bonds to melt diamond.
Good thermal conductor - Strong interatomic bonds transmit heat through vibrations.
Electrical insulator - All outer electrons tied up in localised bonds so no free electrons to carry charge.
Insoluble - Covalent bonds too strong to be broken by solvation.
Silicon, another group 4 element, also forms a similar giant covalent structure where each silicon atom forms 4 strong covalent bonds in a tetrahedral arrangement. Like diamond, silicon has a high melting point, is very hard, and is insoluble due to its network of strong covalent bonds between atoms.
Graphite - bonding, structure, properties
Bonding and structure:
Each carbon atom forms 3 strong covalent bonds in a planar hexagonal pattern, with each carbon contributing 1 delocalised electron.
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Multiple stacked layers of hexagonal carbon arrays with weak intermolecular forces between layers
Properties:
Soft and slippery - Weak intermolecular forces let sheets slide over each other.
Conducts electricity along layers - Delocalised electrons move through the 2D lattice carrying electrical charge.
Lower density than diamond - Weak intermolecular forces lead to increased separation between layers.
High sublimation temperature but lower melting point than diamond - Covalent bonds within each layer are very strong but the weaker intermolecular forces between layers means graphite melts at a lower temperature.
Graphene - bonding, structure, properties
Bonding and structure:
Graphene consists of a single layer of carbon atoms interconnected through strong planar covalent bonds in a hexagonal pattern, with each carbon contributing 1 delocalised electron. This essentially forms a one-atom thick slice of graphite.
Properties:
Excellent electrical and thermal conductivity - Delocalised electrons move through the 2D lattice transporting heat and charge.
Very strong - Extensive network of covalent bonds not easily broken.
Transparent and extremely lightweight - A single layer of atoms light and thin enough to transmit visible light.