Introduction to Bonding

How Do Metals And Non-metals React?

The reactions between metals and non-metals are fundamental to understanding chemical bonding. These reactions primarily involve the transfer of electrons from metal atoms to non-metal atoms, leading to the formation of ionic bonds.

The Driving Force: Achieving Stability

Atoms strive to achieve a stable electronic configuration, typically that of a noble gas (like Neon or Argon), which has a full outermost electron shell (usually 8 valence electrons, an octet, except for Helium which has 2, a duet).

• Metals: Typically have few valence electrons (1, 2, or 3). It is energetically easier for them to lose these valence electrons to achieve a stable configuration of the preceding noble gas. When a metal atom loses electrons, it becomes a positively charged ion, called a cation.
• Non-metals: Typically have many valence electrons (5, 6, or 7). It is energetically easier for them to gain electrons to complete their outermost shell and achieve a stable noble gas configuration. When a non-metal atom gains electrons, it becomes a negatively charged ion, called an anion.

Formation of Ionic Bonds:

When a metal reacts with a non-metal, the metal atom transfers one or more of its valence electrons to the non-metal atom. This electron transfer results in the formation of a cation (from the metal) and an anion (from the non-metal). These oppositely charged ions are then held together by strong electrostatic forces of attraction, known as an ionic bond or electrovalent bond.

The compound formed as a result of this ionic bonding is called an ionic compound.

Example: Formation of Sodium Chloride (NaCl)

Sodium (Na) is an alkali metal with atomic number 11. Its electronic configuration is $2, 8, 1$. It has 1 valence electron.

Chlorine (Cl) is a halogen with atomic number 17. Its electronic configuration is $2, 8, 7$. It has 7 valence electrons.

When sodium reacts with chlorine:

1. The sodium atom loses its single valence electron to become a sodium ion ($$Na^{+}$$) with the stable electron configuration $2, 8$.
$$ \text{Na} \rightarrow \text{Na}^{+} + e^{-} $$
2. The chlorine atom gains this electron to become a chloride ion ($$Cl^{-}$$) with the stable electron configuration $2, 8, 8$.
$$ \text{Cl} + e^{-} \rightarrow \text{Cl}^{-} $$
3. The resulting $$Na^{+}$$ cation and $$Cl^{-}$$ anion are attracted to each other by electrostatic forces, forming an ionic bond and the ionic compound Sodium Chloride (NaCl).
$$ \text{Na}^{+} + \text{Cl}^{-} \rightarrow \text{NaCl} $$

This process is often represented by Lewis structures showing electron transfer:

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Properties Of Ionic Compounds

Ionic compounds, formed by the strong electrostatic attraction between cations and anions, exhibit characteristic physical properties:

1. Physical State:

• Ionic compounds are usually solids at room temperature.
• They exist as crystalline solids, forming a regular, three-dimensional arrangement of alternating positive and negative ions called a crystal lattice. The strong ionic bonds throughout the lattice give them a rigid structure.

2. High Melting and Boiling Points:

• Ionic compounds have high melting and boiling points.
• This is because a large amount of energy is required to overcome the strong electrostatic forces of attraction between the ions in the crystal lattice. The strength of these forces depends on the magnitude of the charges on the ions and the distance between them.

3. Solubility:

• Many ionic compounds are soluble in polar solvents like water.
• Polar solvents have molecules with partial positive and negative charges (dipoles). Water molecules can surround the ions in the lattice, separating them and allowing them to dissolve. The positive ends of water molecules orient towards anions, and the negative ends orient towards cations.
• However, some ionic compounds are insoluble or only sparingly soluble, often due to very strong lattice energies that cannot be overcome by the solvation energy.
• Ionic compounds are generally insoluble in non-polar solvents like petrol or kerosene, as these solvents lack the polarity to disrupt the ionic lattice.

4. Electrical Conductivity:

• Ionic compounds do not conduct electricity in the solid state.
• Reason: In the solid state, the ions are fixed in their positions within the crystal lattice and are not free to move. Electrical conductivity requires the movement of charged particles.
• Ionic compounds conduct electricity when molten (in the liquid state) or dissolved in water (in aqueous solution).
• Reason: In the molten or dissolved state, the ions become mobile and are free to move towards oppositely charged electrodes when an electric potential is applied, thus conducting electricity.

5. Hardness and Brittleness:

• Ionic compounds are generally hard due to the strong electrostatic forces holding the ions together.
• However, they are also brittle. If a force is applied that shifts the layers of ions, ions with the same charge will come in proximity. The resulting repulsion between like charges can cause the crystal to shatter or break.

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Bonding In Carbon – The Covalent Bond

Carbon is a unique element central to organic chemistry. Its ability to form a vast array of compounds stems from its unique bonding characteristics, primarily the formation of covalent bonds.

Carbon's Electronic Configuration and Tetravalence:

• Carbon has an atomic number of 6. Its electronic configuration is $1s^2 2s^2 2p^2$.
• It has 4 valence electrons in its outermost shell (the second shell).
• To achieve a stable noble gas configuration (like Neon, $2s^2 2p^6$), carbon needs to either gain 4 electrons or lose 4 electrons.
• Losing 4 electrons would require a very large amount of energy. Gaining 4 electrons is also energetically unfavorable due to the strong repulsion between the incoming electrons and the existing electron cloud, as well as the relatively small nucleus trying to attract 10 electrons.
• Therefore, carbon achieves stability by sharing its 4 valence electrons with other atoms. This sharing of electrons forms covalent bonds.
• Because carbon has 4 valence electrons, it can form a maximum of four covalent bonds. This property is called tetravalence.

Formation of Covalent Bonds by Carbon:

Carbon atoms can form covalent bonds with:

1. Other Carbon Atoms: This allows carbon to form long chains, branched chains, and rings, leading to the vast diversity of organic compounds.
2. Other Elements: Carbon can also form covalent bonds with other non-metals like Hydrogen (H), Oxygen (O), Nitrogen (N), Sulfur (S), and halogens (F, Cl, Br, I).

Types of Covalent Bonds Formed by Carbon:

Carbon can form three types of covalent bonds depending on the number of electron pairs shared:

• Single Bond (C-C or C-H): Formed when one pair of electrons is shared between two atoms. Each carbon atom contributes one electron to the shared pair.
  • Example: In ethane ($$C_2H_6$$), there is a C-C single bond and C-H single bonds.
• Double Bond (C=C or C=O): Formed when two pairs of electrons are shared between two atoms. Each carbon atom contributes two electrons to the shared pairs.
  • Example: In ethene ($$C_2H_4$$), there is a C=C double bond.
• Triple Bond (C≡C or C≡N): Formed when three pairs of electrons are shared between two atoms. Each carbon atom contributes three electrons.
  • Example: In ethyne (acetylene, $$C_2H_2$$), there is a C≡C triple bond.

The presence of single, double, and triple bonds significantly influences the shape, stability, and reactivity of organic molecules. The ability of carbon to form stable bonds with itself and with other elements, forming diverse structures, is the basis of organic chemistry.