# Covalent Bonding & Structure ## Covalent bonds & strength - **Covalent bonding** occurs between two **non-metals**

A covalent bond is the electrostatic attraction between the nuclei of bonding atoms and a shared pair of outer-shell electrons

- Each atom's positive nucleus attracts the shared bonding electrons in the covalent bond - This attraction allows atoms to achieve a stable noble gas configuration without transferring electrons - Different types of covalent bonds can form, depending on the number of electron pairs shared ### Types of covalent bonds | Type of bond | Representation | Number of shared electrons | |--------------|----------------|---------------------------| | single | $\text{C-C}$ | 2 | | double | $\text{C=C}$ | 4 | | triple | $\text{C}\equiv\text{C}$ | 6 | ## Bond energy

Bond energy is the energy required to break one mole of covalent bonds in the gaseous state

- Bond energy has units of $\text{kJ mol}^{-1}$ - The larger the bond energy, the stronger the covalent bond - **Average bond enthalpy** is also used as a measurement of bond strength - It is the average energy needed to break a specific bond type - It is measured across many different compounds ### Bond strength patterns - Multiple covalent bonds (e.g. double or triple) generally have higher bond enthalpies than single bonds - This is due to the greater electron density between the nuclei - In general, **stronger bonds are shorter bonds**: - Triple bonds are shorter and stronger than double bonds - Double bonds are shorter and stronger than single bonds - This is due to increased electron density pulling the bonded atoms closer together

Triple bonds are the shortest and strongest covalent bonds

- For example, comparing carbon-carbon bonds: - $\text{C-C}$: 147 pm, 347 kJ mol$^{-1}$ - $\text{C=C}$: 134 pm, 614 kJ mol$^{-1}$ - $\text{C}\equiv\text{C}$: 120 pm, 839 kJ mol$^{-1}$ The relationship between bond type and bond length can be visualized:

The relationship between bond type and bond strength shows the inverse pattern:

The bond enthalpy given in exam questions may differ from average bond enthalpy values in data tables

## Dot & cross diagrams - **Dot and cross diagrams** are used to represent covalent bonding - They show only the **outer shell electrons** of each atom - To differentiate between the two atoms involved, **dots** for electrons of one atom and **crosses** for electrons of the other atom are used - Electrons are shown in **pairs** on dot-and-cross diagrams ### Single covalent bonding **Chlorine, $\text{Cl}_2$** - Each chlorine atom has 7 outer electrons - Each chlorine atom needs 1 more electron for a full outer shell - Two chlorine atoms share one pair of electrons to form a single covalent bond - Each chlorine atom now has 8 electrons in its outer shell (a stable configuration) **Hydrogen chloride, $\text{HCl}$** - Hydrogen has 1 electron in its outer shell and needs 2 for a full outer shell - Chlorine has 7 electrons in its outer shell and needs 1 more to have a full outer shell - Both atoms share an electron to form a single covalent bond - Each atom now has enough electrons to have a full outer shell **Ammonia, $\text{NH}_3$** - Nitrogen is in Group 15 so it requires 3 more electrons to complete its outer shell - Each hydrogen atom can donate an electron to form a covalent bond - Nitrogen has formed 3 single bonds with each of the hydrogen atoms - Nitrogen in ammonia has a non-bonding pair of electrons (a **lone pair**) ### Double covalent bonding **Oxygen, $\text{O}_2$** - Each oxygen atom has 6 electrons in its outer shell - Two oxygen atoms share 4 electrons to form a double covalent bond - Once bonded, each oxygen atom's octet is complete - Double bonds are stronger than single bonds **Carbon dioxide, $\text{CO}_2$** - Carbon is in Group 14, so it requires 4 more electrons to achieve a full outer shell - Each oxygen atom has 6 outer electrons - 4 electrons are shared between the carbon atom and oxygen in each $\text{C=O}$ double covalent bond - In $\text{CO}_2$, each oxygen atom has 2 non-bonded electron pairs (lone pairs) **Ethene, $\text{C}_2\text{H}_4$** - Each hydrogen atom has 1 electron in its outer shell and 4 hydrogen atoms are required to make ethene - Each carbon atom has 4 electrons in its outer shell - Each carbon atom shares 2 electrons to form a double bond - Carbon's remaining 2 electrons are shared with 2 hydrogen atoms to form 2 single bonds ### Triple covalent bonding **Nitrogen, $\text{N}_2$** - Each nitrogen atom has 5 electrons in its outer shell - Each nitrogen atom shares 3 electrons to form a triple bond - In total, 6 electrons are needed to form a triple bond - Once bonded, each nitrogen atom has a complete octet

Use the periodic table to decide how many electrons are in the outer shell of a nonmetal atom

## Exceptions to the octet rule - In some instances, the central atom of a covalently bonded molecule can accommodate **more** or **less** than 8 electrons in its outer shell ### Expanded octet - Being able to accommodate **more than 8 electrons** in the outer shell is known as **expanding the octet rule** - Some examples of this occurring can be seen with period 3 elements **Sulfur dioxide, $\text{SO}_2$** - Sulfur is in Group 16, so it requires 2 more electrons to achieve a full outer shell - Each oxygen atom has 6 electrons in its outer shell - 4 electrons are shared between the sulfur and oxygen in each $\text{S=O}$ bond - Sulfur has 10 electrons in its outer shell (expanding the octet rule) - Sulfur also has a lone pair of electrons **Phosphorus pentachloride, $\text{PCl}_5$** - Phosphorus is in Group 15, so it requires 3 more electrons to achieve a full outer shell - Each chlorine has 7 electrons in its outer shell - 2 electrons are shared between the phosphorus and chlorine atoms in the $\text{P-Cl}$ bond - Phosphorus has 10 electrons in its outershell (expanding the octet rule) ### Electron deficient - Accommodating **less than 8 electrons** in the outer shell means that the central atom is **electron deficient** **Boron trichloride, $\text{BCl}_3$** - Boron only has 3 valence electrons - Each chlorine has 7 electrons in its outer shell - A shared pair of electrons to form a single covalent bond occurs between boron and each chlorine - Boron is electron deficient with 6 electrons - Each chlorine atom has 8 electrons in its outer shell ## Dative covalent / coordinate bonding - In regular covalent bonds, each atom contributes one electron to the shared pair - In **dative covalent bonding**, both electrons in the bond come from the **same atom** - This can also be called **coordinate bonding**

A dative covalent bond occurs when one atom has a lone pair and the other atom is electron-deficient

- This means that the other atom has an incomplete outer shell - A dative covalent bond is shown using an arrow ($\rightarrow$) pointing from the atom donating the lone pair to the atom accepting it ### Formation of the ammonium ion ($\text{NH}_4^+$) - The nitrogen atom in **ammonia** ($\text{NH}_3$) has a **lone pair** - A **hydrogen ion** ($\text{H}^+$) has **no electrons**, so it can accept a lone pair to form a dative bond - This forms the **ammonium ion** ($\text{NH}_4^+$) - The arrow notation shows the dative bond from nitrogen to $\text{H}^+$ ### Dimer formation in aluminium chloride ($\text{Al}_2\text{Cl}_6$) - At high temperatures, aluminium chloride exists as $\text{AlCl}_3$ - This compound is electron-deficient - At lower temperatures, two $\text{AlCl}_3$ molecules join to form $\text{Al}_2\text{Cl}_6$ - Two chlorine atoms each donate a lone pair to an aluminium atom - This forms two dative covalent bonds - An aluminium chloride dimer is formed using 2 dative covalent bonds, which is more stable than $\text{AlCl}_3$