Molecular and Intermolecular Forces

• Molecular forces are the bonding forces within or between molecules.

Intramolecular forces

• Intramolecular forces are the bonds that hold atoms or ions together within a substance.
• They are also called chemical bonds.
• Intramolecular forces determine the chemical properties of substances.
• They encompass strong forces, including ionic, covalent, and metallic bonds.

Intermolecular forces

• Intermolecular forces are forces or bonds between the particles (molecules or ions) of a substance.
• They hold particles together in liquids and solids.
• Intermolecular forces determine the physical properties of substances.
• They are weak relative to intramolecular forces.
• Types of intermolecular forces are dipole-dipole forces, ion-dipole forces, hydrogen bonding, and London dispersion forces.

Chemical bonds

• Attractive forces that hold atoms together in compounds are called chemical bonds.
• Atoms bond together to achieve a more stable, lower-energy state than existing separately.
• Bond formation is an exothermic reaction, releasing energy.
• Breaking bonds is an endothermic reaction, absorbing energy.
• Valence electrons play a crucial role in chemical bonding.

Octet and Duet rules

• Representative elements tend to follow the octet or duet rule when forming chemical bonds.
• The octet rule states that atoms lose, gain, or share valence electrons to achieve 8 electrons in their outer shell, resembling the nearest noble gas configuration (ns²np⁶).
• The duet rule applies to some atoms, leading them to lose, gain, or share valence electrons until they have 2 electrons in their outer shell, similar to the electron configuration of helium (He: 1s²); examples include Hydrogen (H), Lithium (Li), and Beryllium (Be).

Lewis structures

• Lewis structures are a way to keep track of valence electrons in an atom.
• In a Lewis symbol, the element's chemical symbol represents the nucleus and core electrons.
• Valence electrons are shown as dots around the symbol, with each dot representing one valence electron.
• When writing Lewis symbols, the first four dots are placed singly on each side of the symbol.
• Dots are paired as the next four are added.

Ionic Bonds

• Ionic, covalent, and metallic bonds are several types of chemical bonds that hold atoms together.
• Atoms tend to lose or gain electrons to acquire the electron configuration of the nearest noble gas.
• Metals and nonmetals react together, with valence electrons transferring from metal to nonmetal atoms, forming positive and negative ions.
• Electrostatic attraction between oppositely charged ions leads to ionic bonds.
• Ionic bonds mostly form ionic compounds.
• The formation of sodium chloride salt from sodium metal and chlorine gas is an example.
• Ionic compounds are polar.
• They have high melting and boiling points.
• Ionic compounds are often soluble in water (hydrophilic).
• They are insoluble in nonpolar solvents like chloroform and benzene.
• Ionic compounds do not conduct electricity as solids, but do when molten or dissolved in water.

Covalent Bonds

• Covalent bonds link nonmetal atoms to form covalent compounds.
• Covalent bond forms by sharing valence electrons between bonded atoms, so each atom acquires the electron configuration of the nearest noble gas.
• The simplest way of indicating covalent bonds in molecules is the use of line-bond structures (Kekulé structures) with a line drawn between atoms.
• The number of covalent bonds an atom forms depends on how many additional valence electrons it needs to reach a noble-gas configuration.
• Hydrogen has one valence electron (1s¹) and needs one more to reach the helium configuration (1s²), so Hydrogen forms one bond.
• Carbon has four valence electrons (2s²2p²) and needs four more to reach the neon configuration (2s²2p⁶), so Carbon forms four bonds.
• Nitrogen has five valence electrons (2s²2p³), needs three more, and forms three bonds.
• Oxygen has six valence electrons (2s²2p⁴), needs two more, and forms two bonds.
• The halogens have seven valence electrons, need one more, and form one bond.
• Covalent bonding can be single, double, or triple, depending on the number of electron pairs shared.
• Valence electrons that are not used for bonding are called lone-pair electrons or nonbonding electrons.
• A bond pair is a pair of electrons involved in covalent bonding.
• A lone pair is a pair of electrons not involved in covalent bonding.

Electronegativity and bond type

• Electronegativity difference (ΔEN) between two atoms determines the bond type: ionic or covalent.
• Covalent bonds can be either polar or nonpolar.
• Nonmetal atoms have little or no difference in electronegativity (EN) in nonpolar covalent bonds, with electron pairs shared equally, ΔEN < 0.5 (e.g., H₂, Cl₂, CH₄).
• Nonmetal atoms have a significant difference in electronegativity (EN) in polar covalent bonds, with electron pairs shared unequally, drawn closer to the atom with higher EN, 0.5 < ΔEN < 1.7 (e.g., H₂O, HCl).
• Metal and nonmetal atoms have significant differences in electronegativity (EN) in ionic bonds, with electrons completely transferred from metal to nonmetal atoms, ΔEN ≥ 1.7 (e.g., NaCl, CaCl₂, KBr).

Metallic Bonds

• Metallic bonds occur among metal atoms.
• Atoms are packed closely together in a regular arrangement called a lattice.
• Each atom loses all valence electrons, forming a pool of electrons (sea of electrons).
• Positively charged metal ions are held together by electrostatic attraction to the pool of electrons.
• Outer-shell electrons are free to move (delocalized or mobile electrons), not associated with any one atom or bond.
• Metallic bonding is strong, due to the electrostatic attraction between positive ions and negative delocalized electrons, acting in all directions.
• Metallic bonding strength increases with more mobile electrons per atom.
• It also increases with increasing the positive charge on metal ions in the lattice.
• The strength decreases with decreasing size of metal ions in the lattice.
• Ionic compounds or salts are formed through ionic bonds between metal and nonmetal atoms, with electrons transferred from metal to nonmetal atoms.
• Covalent compounds (molecules) occur between nonmetals with electrons shared between atoms.
• Metallic bonds occur in metals and alloys with a sea of electrons.
• Full charges result from the full transfer of electrons in ionic bonds (e.g., NaCl, KF).
• Partial charges result from unequal sharing of bond electron pairs in polar molecules (positive and negative ends) (e.g., HCl, H₂O).

Intermolecular Forces Types

• Types of intermolecular forces include dipole-dipole forces, ion-dipole forces, dispersion forces, and hydrogen bonding.
• The greater the intermolecular force between particles, the higher the substance's boiling point.
• More heat energy is required to overcome the intermolecular forces to change a liquid into a gas.

Dipole-Dipole Interactions

• Permanent dipole-dipole interactions occur between polar covalent molecules.
• A polar covalent molecule has a partially positive end (δ+) and a partially negative end (δ−).
• When polar covalent molecules come close, opposite charges of the dipoles attract one another.
• The polarity of molecules depends on electronegativity differences between bonded atoms.
• More polar molecules have stronger dipole-dipole forces.
• Polar covalent molecules tend to have higher boiling points than non-polar substances of similar molecular weight.
• As molecule polarity increases, the boiling point temperature increases.

Ion-Dipole Forces

• Ion-dipole forces occur between ions and polar solvents.
• For example, when an ionic compound like table salt (NaCl) dissolves in a polar solvent like water.
• Ion-dipole force results from the electrostatic attraction between an ion and a neutral molecule with a dipole (polar covalent compound).
• A positive ion (cation) attracts the partially negative end of a neutral polar molecule.
• A negative ion (anion) attracts the partially positive end of a neutral polar molecule.

Hydrogen Bonds

• A hydrogen bond occurs between partially positive hydrogen (H) of one molecule and a partially electronegative atom of another molecule, mainly Nitrogen, Oxygen, and Fluorine atoms (N, O, F).
• Electronegativities of Nitrogen, Oxygen, and Fluorine are 3, 3.5, and 4, respectively.
• Hydrogen bonds are of two types: Intermolecular and Intramolecular.
• Intermolecular hydrogen bonds (Between different molecules).
• The classic example of intermolecular hydrogen bonds is the hydrogen bonds between water molecules.
• Each water molecule can form four hydrogen bonds with the other four water molecules.
• Water has a high boiling point temperature, with respect to its tiny mass, because of the intermolecular hydrogen bonds between its molecules.
• Intramolecular hydrogen bonds (Inside the same molecule).
• The intramolecular hydrogen bond is formed between a hydrogen atom and a highly electronegative atom (N, O, or F) inside the same molecule.

Dispersion Forces

• Dispersion forces are also known as "London forces", Dipole-induced dipole forces, or Van der Waals forces.
• Dispersion forces are the weakest intermolecular attractive forces that are important only over extremely short distances.
• Dispersion forces are the only type of forces present in non-polar molecules such as CO₂, O₂, N₂, Cl₂, Br₂, H₂, and monatomic species such as noble gases.
• London forces are also responsible for the condensation of these substances.
• Dispersion forces result from the attraction of the positively charged nucleus of one atom for the electron cloud of an atom in the nearby molecules.
• This induces a temporary dipole between the atoms or molecules.
• The temporary (momentary) dipole in one molecule INDUCES a similar dipole in a neighboring molecule.
• The resulting temporary dipoles cause weak attraction among the molecules.
• London dispersion force is the weakest intermolecular force.
• Its strength depends on molecular polarizability.
• The greater the polarizability of molecules, the stronger the dispersion forces between them.
• Polarizability increases with increasing numbers of electrons and sizes of molecules.
• Polarizability depends on molecular shape or surface area.
• The elongated molecule is more easily polarized than the compact molecule.
• Polarizability depends on Molecular Size or Weight:
• Greater molecular size or weight causes larger, more polarizable electron clouds.
• The strength order of intermolecular forces is ion-dipole forces > hydrogen bonds > dipole-dipole forces > london forces.