Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry → Chemical bond ↓


Bond polarity and dipole moment


Bond polarity and dipole moment are fundamental concepts in chemistry, especially when it comes to understanding molecular interactions and properties. These concepts revolve around the idea that atoms within a molecule do not always share electrons equally in a covalent bond. This unequal sharing leads to the formation of polar bonds and impacts the overall behavior of the molecule, affecting everything from melting and boiling points to solubility and reactivity. In this comprehensive guide, we will break down these concepts, providing detailed explanations, examples, and diagrams to ensure a deep understanding of bond polarity and dipole moments.

1. Understanding electronegativities

To understand the concept of bond polarity, one must first understand electronegativity. Electronegativity is a measure of an atom's ability to attract and hold electrons within a bond. It varies across the periodic table, generally increasing from left to right across a period and decreasing down a group. Elements with high electronegativities, such as fluorine, oxygen, and nitrogen, are more likely to attract electrons, while elements with low electronegativities, such as sodium and cesium, are less likely.

        Electronegativity trends in the Periodic Table: - Increases across a period from left to right. - Decreases down a group from top to bottom.

2. Bond polarity

A bond is considered polar when there is a significant difference in electronegativities between the two atoms involved. When one atom's electronegativities are higher than those of the other, it pulls the shared electrons closer to itself, forming a dipole. A dipole is essentially a separation of charge within a molecule.

For example, in the molecule of hydrogen fluoride (HF), fluorine is more electronegative than hydrogen. Thus, the shared electrons are more closely bound to the fluorine atom, resulting in a partial negative charge (δ-) on fluorine and a partial positive charge (δ+) on hydrogen.

H F δ+ δ-

3. Nonpolar bonds

In contrast, a nonpolar bond occurs when the electronegativities of the atoms are equal or very similar. This means that the electrons are shared equally, and there is no charge separation. A classic example can be found in diatomic molecules such as nitrogen (N2) or oxygen (O2), where the two atoms involved are identical, and thus, their electronegativities cancel each other out.

Hey Hey Nonpolar

4. Dipole moment

When a molecule has polar bonds, it is possible that these dipoles do not cancel out if they are asymmetrically oriented. In such cases, the molecule as a whole is said to have a dipole moment. It is a vector quantity, meaning it has both magnitude and direction, usually measured in Debye units (D).

The dipole moment depends on both the polarity of the individual bonds and the geometry of the molecule. For example, although carbon dioxide (CO2) has polar bonds, the dipole moments cancel out due to the linear geometry, resulting in a net dipole moment of zero, making the molecule nonpolar.

Hey Hey C no net dipole moment

5. Polar and nonpolar molecules

Identifying whether a molecule is polar or nonpolar involves evaluating both bond polarity and molecular geometry. Some molecules, such as water (H2O), have bent shapes that cause the dipoles from the polar OH bonds to not cancel, resulting in a net dipole moment. Thus, water is a polar molecule.

Hey H H Net dipole moment

6. Applications of bond polarity and dipole moment

Understanding bond polarity and dipole moments is important for predicting the behavior of substances. These concepts play an important role in determining solubility, boiling and melting points, and intermolecular forces. For example, polar solvents such as water are excellent at dissolving polar solutes due to the attraction between dipoles. Similarly, the high boiling point of water can be attributed to its strong hydrogen bonding, which arises from its polarity.

Dipole moments also affect reactivity and interactions in chemical reactions. Polar molecules have different reaction pathways than nonpolar molecules and can interact with electric fields, which is essential in spectroscopy and other analytical techniques.

7. Calculation of dipole moment

The dipole moment (μ) of a molecule can be calculated using the following equation:

        μ = Q * r

where Q is the magnitude of the charge difference, and r is the distance between the charges. While accurate measurements require quantum mechanical calculations, this equation provides a basic understanding of the factors affecting dipole moments.

8. Visualization of dipole moments

To visualize dipole moments, consider a water molecule in which the dipole is represented as an arrow pointing from the positive pole to the negative pole. This can help predict molecular interactions:

δ-O H δ+ H δ+ Dipole

9. Summary

Bond polarity and dipole moment are important for understanding how molecules interact with each other and their environment. The polarity of a bond arises from differences in electronegativities, creating dipole moments, which affect the physical and chemical properties of the molecule. A solid understanding of these phenomena helps predict behaviors such as solubility, reactivity, and interactions with electric fields.

By understanding both molecular geometry and the nature of bonds, chemists can predict molecular behavior and design molecules with specific properties for industrial, pharmaceutical, and technological applications. This understanding is important not only for pure chemistry but also for fields such as biology, environmental science, and materials science.


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Bond polarity and dipole moment

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