Chemical bonds are the attractive forces that hold atoms together, allowing the formation of molecules and compounds. Atoms form these bonds to achieve greater stability by reaching a lower energy state, often resembling the electron configuration of noble gases.
Bonding and Atomic Stability
Atoms are inherently reactive because most do not possess a full outer electron shell. To become more stable, they interact with other atoms through bonding. These interactions involve their outermost, or valence electrons, which are the electrons located in the highest occupied energy level.
Valence Electrons and Their Role in Bonding
Valence electrons are responsible for determining how atoms bond with one another. They are found in the s and p orbitals of the outermost electron shell. Atoms with full valence shells, such as noble gases, are chemically inert. In contrast, atoms with incomplete outer shells tend to be more reactive as they seek to gain, lose, or share electrons.
Key characteristics of valence electrons:
Valence electrons are the only electrons involved in chemical bonding.
They determine the reactivity and bonding behavior of elements.
Elements in the same group of the periodic table have the same number of valence electrons, which explains their similar chemical properties.
Ionization energy data can be used to determine the number of valence electrons in an atom. A large jump between successive ionization energies indicates the removal of a core electron after all valence electrons have been ionized.
Image Courtesy; This table displays the successive ionization energies of elements in period three on the periodic table. Here you can see the jump in ionization energies and how that corresponds to the number of valence electrons an atom has.
Because valence electrons are involved in bonding, they dictate whether atoms will share, donate, or accept electrons, ultimately determining the type of bond that will form.
Electronegativity and Bond Type
Electronegativity is a property that describes an atom’s ability to attract and hold onto shared electrons within a chemical bond. Understanding this concept is key to predicting whether a bond will be ionic, polar covalent, or nonpolar covalent.
Electronegativity trends:
It increases from left to right across a period due to increasing nuclear charge, which pulls the electrons closer to the nucleus.
It decreases down a group because atoms become larger and their valence electrons are farther from the nucleus, weakening the attractive force.
Fluorine is the most electronegative element, with a value of 3.98 on the Pauling scale. Elements like cesium and francium have very low electronegativities, around 0.7.
The difference in electronegativity between two bonding atoms is critical:
A large difference (greater than 1.7) suggests an ionic bond.
A moderate difference (0.4 to 1.7) suggests a polar covalent bond.
A small or zero difference (less than 0.4) suggests a nonpolar covalent bond.
Coulomb’s Law and Electrostatic Attraction
Coulomb’s Law describes the force of attraction or repulsion between two charged particles. It is expressed as:
Force is proportional to (q₁ × q₂) / r²
Where:
q₁ and q₂ are the charges of the particles,
r is the distance between them.
According to Coulomb’s Law:
The greater the magnitude of the charges, the stronger the force of attraction or repulsion.
The smaller the distance between the charges, the stronger the force.
This law helps explain why smaller ions with greater charges form stronger ionic bonds and why these bonds result in compounds with high melting and boiling points.
Ionic Bonding
How Ionic Bonds Form
An ionic bond forms when electrons are transferred from one atom to another. This process creates two oppositely charged ions that are attracted to each other by electrostatic forces. Ionic bonds typically form between metals and nonmetals.
The metal loses one or more electrons, becoming a cation (positively charged ion).
The nonmetal gains one or more electrons, becoming an anion (negatively charged ion).
Example: Sodium and chlorine form an ionic bond as follows:
Na (metal) loses one electron → Na⁺
Cl (nonmetal) gains one electron → Cl⁻
These ions then attract each other → NaCl
This transfer of electrons leads to the formation of a stable ionic compound with a crystal lattice structure.
Structure and Properties of Ionic Compounds
Ionic compounds are characterized by the formation of a crystal lattice—a repeating, three-dimensional structure of alternating positive and negative ions. This highly ordered structure contributes to several key physical properties:
High melting and boiling points: Strong electrostatic forces between ions require significant energy to break.
Brittleness: The rigid lattice structure can fracture when force causes like-charged ions to align and repel each other.
Electrical conductivity: Ionic compounds conduct electricity when melted or dissolved in water because the ions are free to move and carry an electric current.
Solubility in water: Many ionic compounds dissolve in water because water molecules surround and separate the individual ions, overcoming the lattice energy.
Coulomb’s Law and Ionic Strength
The strength of an ionic bond—and by extension, the melting and boiling points of the compound—depends on:
The charges of the ions: Greater charges lead to stronger attractions.
The radii of the ions: Smaller ions allow for closer approach and stronger attractions.
Examples:
Magnesium fluoride (MgF₂) vs. Sodium fluoride (NaF)
Mg²⁺ has a higher charge than Na⁺, leading to a stronger attraction with F⁻. Thus, MgF₂ has a higher melting point.Lithium fluoride (LiF) vs. Sodium bromide (NaBr)
Although both compounds involve ±1 charges, Li⁺ and F⁻ are smaller than Na⁺ and Br⁻, resulting in a stronger attraction and higher melting point for LiF.
Covalent Bonding
Formation and Characteristics
Covalent bonds involve the sharing of electrons between atoms. This usually occurs between nonmetals. The shared electrons are attracted to the nuclei of both atoms, creating a stable bond.
Unlike ionic bonds, covalent bonds do not result in the formation of ions. Instead, the atoms remain neutral but are held together by the shared pair(s) of electrons.
Polar vs. Nonpolar Covalent Bonds
The nature of a covalent bond—whether polar or nonpolar—depends on the difference in electronegativity between the atoms involved.
Nonpolar Covalent Bonds
Electrons are shared equally.
Typically occur between identical atoms or atoms with very similar electronegativity.
There is no separation of charge.
Example: O₂ (oxygen molecule) – both oxygen atoms have the same electronegativity, resulting in equal sharing of electrons.
Image Courtesy of cK-12; In this oxygen molecule, there is a nonpolar covalent bond between the two oxygen atoms!
Polar Covalent Bonds
Electrons are shared unequally.
The more electronegative atom attracts the shared electrons more strongly, creating a partial negative charge (δ⁻).
The less electronegative atom becomes partially positive (δ⁺).
Example: H₂O – oxygen is more electronegative than hydrogen, so it attracts electrons more, leading to two polar bonds.
We could see two O-H polar covalent bonds here! These attractions are very strong and form water which, as you know, is essential for life on Earth.
These bonds create bond dipoles, directional differences in electron density that influence molecular polarity and interactions.
Predicting Bond Type with Electronegativity Difference
Ionic bond: Difference in electronegativity > 1.7
Polar covalent bond: Difference between 0.4 and 1.7
Nonpolar covalent bond: Difference < 0.4
Image Courtesy of Socratic; δ, or lowercase delta, represents a partial positive or partial negative charge.
These thresholds provide a guideline for determining the type of bond likely to form between two atoms.
Examples of Covalent Compounds
Water (H₂O)
Polar covalent bonds exist between hydrogen and oxygen. The molecule is bent, and the oxygen atom carries a partial negative charge.Carbon dioxide (CO₂)
Although the C=O bonds are polar, the molecule is linear and symmetrical, making it nonpolar overall.Methane (CH₄)
The electronegativity difference between C and H is small, making the bonds largely nonpolar.
Comparing Ionic and Covalent Compounds
Ionic Compounds
Typically form between metals and nonmetals.
Have high melting and boiling points.
Conduct electricity in liquid or dissolved form.
Form crystal lattices.
Tend to be soluble in water.
Covalent Compounds
Form between nonmetals.
Can be polar or nonpolar.
Often have lower melting and boiling points.
Do not conduct electricity in any state.
Exist as discrete molecules rather than extended networks.
An additional case is the network covalent solid, such as diamond or quartz. These substances are composed of atoms connected by covalent bonds in a continuous network. They have high melting points and are not conductive. This type of structure is explored in later topics.
Summary of Key Concepts and Vocabulary
Valence Electrons: Outermost electrons responsible for bonding.
Electronegativity: Atom’s ability to attract shared electrons.
Cation: Atom that loses electrons, becoming positively charged.
Anion: Atom that gains electrons, becoming negatively charged.
Ionic Bond: Electrostatic attraction between oppositely charged ions.
Covalent Bond: Sharing of electrons between atoms.
Polar Covalent Bond: Unequal electron sharing resulting in partial charges.
Nonpolar Covalent Bond: Equal sharing of electrons with no charge separation.
Crystal Lattice: Structured arrangement of ions in an ionic compound.
Coulomb’s Law: Describes the strength of force between charged particles.
Bond Dipole: Separation of charge within a polar bond.
Network Solid: Extensive covalently bonded structure with high melting point.
These fundamental concepts are essential for understanding molecular interactions, chemical reactivity, and the physical properties of substances in biological systems.
FAQ
Noble gases are unreactive because they have full valence electron shells, typically containing eight electrons (except helium, which has two). This full outer shell is extremely stable, meaning noble gases are already at a low energy state and have no tendency to gain, lose, or share electrons. As a result:
They have no chemical "drive" to form bonds.
Ionization energy is high, making it difficult to remove electrons.
Electron affinity is low or even positive, meaning they don’t readily accept electrons.
Their filled s and p orbitals make them chemically inert.
This stability is the reason other elements form bonds—to achieve a noble gas electron configuration.
Yes, metallic bonding is a distinct type of chemical bond, though not the focus of this subtopic. While ionic and covalent bonds involve localized electrons (transferred or shared between specific atoms), metallic bonding features a "sea of electrons" that are delocalized and move freely among metal cations. Key characteristics include:
Occurs between metal atoms with low electronegativity.
Valence electrons are not associated with individual atoms.
Electrons form a mobile cloud that allows conductivity and malleability.
Explains high thermal and electrical conductivity in metals.
Although different in structure and behavior, metallic bonding is a true chemical bond due to the electrostatic attraction between positive metal ions and free-moving electrons.
Bond polarity and molecular polarity are related but distinct concepts. Bond polarity refers to the unequal sharing of electrons in a single bond, while molecular polarity involves the overall distribution of charge across an entire molecule. Here's the distinction:
Bond polarity occurs when two atoms have differing electronegativities, leading to partial charges within the bond.
Molecular polarity considers the geometry of the molecule and whether the bond dipoles cancel out.
A molecule like carbon dioxide (CO2) has polar bonds but is nonpolar overall due to its linear shape.
Water (H2O), with a bent geometry, is a polar molecule because dipoles do not cancel.
So, molecular polarity is a sum of both bond polarities and molecular shape.
The number of shared electron pairs in covalent bonds directly influences both bond strength and length:
Single bonds share one pair of electrons and are the longest and weakest among covalent bonds.
Double bonds share two pairs and are shorter and stronger than single bonds.
Triple bonds share three pairs of electrons, making them the shortest and strongest.
This occurs because more shared electrons increase the attractive forces between the nuclei and the bonding electrons. As the number of shared electrons increases, the bonded atoms are pulled closer together, reducing bond length and increasing bond strength. Therefore, bond order correlates with both bond length and bond energy.
Partial ionic character arises when electrons in a covalent bond are shared unequally due to a significant difference in electronegativity between the bonded atoms. Though the bond is covalent, the uneven distribution of electron density causes:
A partial negative charge (δ⁻) on the more electronegative atom.
A partial positive charge (δ⁺) on the less electronegative atom.
This creates a dipole moment, a measurable quantity that indicates the magnitude of polarity in the bond. The dipole moment is:Proportional to both the amount of charge separation and the distance between charges.
Expressed in debye units (D).
A larger dipole moment means greater partial ionic character. For example, the O-H bond in water has a higher dipole moment than the C-H bond in methane, reflecting more ionic character in the water molecule’s bonds.
Practice Questions
Explain how differences in electronegativity determine the type of chemical bond formed between two atoms. Use examples in your response.
Electronegativity differences between atoms determine whether a bond is ionic, polar covalent, or nonpolar covalent. If the difference is large, typically greater than 1.7, one atom completely transfers electrons to another, forming an ionic bond, as seen in sodium chloride (NaCl). If the difference is moderate, between 0.4 and 1.7, electrons are shared unequally, forming a polar covalent bond, such as in water (H2O). When the difference is very small or zero, as with two oxygen atoms in O2, electrons are shared equally, resulting in a nonpolar covalent bond. Thus, electronegativity directly influences bond type and molecular properties.
Compare the physical properties of ionic and covalent compounds, and explain how these differences are related to their bond types.
Ionic compounds, like sodium fluoride (NaF), have high melting and boiling points due to strong electrostatic forces between oppositely charged ions in a crystal lattice. They conduct electricity when melted or dissolved because their ions are mobile. In contrast, covalent compounds like carbon dioxide (CO2) often have lower melting and boiling points due to weaker intermolecular forces. They generally do not conduct electricity in any state because they lack free-moving charged particles. These differences arise because ionic bonds involve electron transfer and strong attractions between ions, while covalent bonds involve electron sharing without creating free ions for conduction.
