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Types of Chemical Bonds study notes -AP Chemistry - New Syllabus 2024-2025

Types of Chemical Bonds Study Notes -AP Chemistry

AP Chemistry: Types of Chemical Bonds Study Notes. Comprehensive coverage of topics. Prepare for the AP Chemistry Exam

LEARNING OBJECTIVE

  • Explain the relationship between the type of bonding and the properties of the elements participating in the bond.

Key Concepts:

  • Electronegativity Values
  • Covalent Bonds
  • Ionic vs Covalent
  • Valence Electrons in a Metallic Solid
  • Intramolecular Force & Potential Energy
  • Coulomb’s Law & Attractive Forces

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 2.1.A.1 Electronegativity

Definition:

Electronegativity is the ability of an atom to attract electrons toward itself when forming a chemical bond. The higher the electronegativity, the stronger the atom pulls bonding electrons.

Electronegativity determines how strongly an atom attracts bonding electrons. Atoms with higher electronegativity pull electrons more strongly, while those with lower electronegativity attract electrons less, influencing whether the bond is ionic, covalent, or polar covalent.

Why It Is Important To Understand Electronegativity:

Electronegativity plays a key role in determining the type of bond formed between atoms:

  1. Ionic Bond: Formed when there is a large difference in electronegativity between two atoms (usually between a metal and a nonmetal). The more electronegative atom gains electrons from the less electronegative atom, resulting in the formation of oppositely charged ions.

  2. Covalent Bond: Formed when two atoms have similar electronegativities. Electrons are shared equally between the atoms (e.g., in molecules like H₂ or O₂).

  3. Polar Covalent Bond: Formed when there is a moderate difference in electronegativity. The electrons are shared unequally, with the more electronegative atom pulling the electrons closer to itself, creating a partial negative charge on one atom and a partial positive charge on the other (e.g., in H₂O).

Periodic trends of Electronegativity:

Across a Period (Left to Right):Electronegativity increases as you move across a period.

Down a Group (Top to Bottom):Electronegativity decreases as you move down a group.

Correlation with Atomic Size:

    • As atomic size decreases across a period, the nucleus attracts electrons more strongly, leading to higher electronegativity.
    • As atomic size increases down a group, the outer electrons are farther from the nucleus, and the attraction weakens, resulting in lower electronegativity.
    • Elements with electron configurations that are closer to achieving a full outer shell (an octet) have higher electronegativity because they are more likely to accept electrons to complete their valence shell and achieve stability.

How Electronic Structure of  an Atoms affects the Electronegativity:

1. Number of Valence Electrons:

  • Atoms with more valence electrons tend to have higher electronegativity because they are closer to completing their outer electron shell (which follows the octet rule for most elements). These elements are more eager to attract electrons to achieve a stable configuration.
  • Example: Oxygen (O) has a configuration of
    1s22s22p41s^2 2s^2 2p^4

    , and it is only two electrons away from filling its outer shell, making it highly electronegative.

2. Effective Nuclear Charge (Z_eff):

  • The effective nuclear charge is the net positive charge experienced by an electron in an atom, which is the result of the protons in the nucleus attracting the electrons, minus the shielding effect from inner electrons.
  • Electron configuration influences the effective nuclear charge. Atoms with more protons and fewer inner electrons (i.e., a higher Z_eff) have a greater pull on their valence electrons, increasing their electronegativity.
  • Example: Fluorine (F) has a high electronegativity because its configuration is
    1s22s22p51s^2 2s^2 2p^5

    , meaning it has a high Z_eff and is very effective at attracting electrons.

3. Shielding Effect:

  • Electron shielding occurs when inner electrons “shield” outer electrons from the full charge of the nucleus. This reduces the attraction between the nucleus and valence electrons.
  • Elements with more electron shells (like those further down the periodic table) have more shielding, reducing their ability to attract electrons. This leads to lower electronegativity.
  • Example: Sodium (Na) with electron configuration
    1s22s22p63s11s^2 2s^2 2p^6 3s^1

    has a lower electronegativity because the outer electron is farther from the nucleus and is shielded by more inner electrons.

4. Atomic Size:

  • Electron configuration influences the atomic size (or radius). Elements with more electron shells (found lower in the periodic table) have larger atoms. Larger atoms have their valence electrons farther from the nucleus, leading to a weaker attraction between the nucleus and bonding electrons.
  • As atomic size increases, electronegativity decreases because the attraction for electrons becomes weaker due to greater distance and more shielding.
  • Example: Chlorine (Cl) with a configuration of
    1s22s22p63s23p51s^2 2s^2 2p^6 3s^2 3p^5

    is smaller than sodium, giving it a stronger ability to attract electrons.

5. Octet Rule and Stability:

      • Elements with electron configurations that are closer to achieving a full outer shell (an octet) have higher electronegativity because they are more likely to accept electrons to complete their valence shell and achieve stability.
      • Example: The noble gases (e.g., Neon, Argon) have full outer electron shells and very low electronegativity because they are already stable and don’t tend to attract electrons.
 The Shell Model:
                 1.Distance from the Nucleus: The closer the valence electrons are to the nucleus, the stronger the attraction, leading to higher electronegativity. Conversely, the farther the                         valence electrons are, the weaker the attraction, resulting in lower electronegativity.

                 2.Electron Shielding: Electron shielding occurs when inner electrons reduce the effective attraction between the nucleus and valence electrons. This occurs because inner electrons “block” or “shield” the outer electrons from the full positive charge of the nucleus.

Effect on Electronegativity:

  • More shielding (due to more inner electrons) decreases the nucleus’s ability to attract valence electrons, resulting in lower electronegativity.
  • Less shielding (fewer inner electrons) allows the nucleus to attract valence electrons more strongly, leading to higher electronegativity.
 

{\displaystyle S(\mathrm {s} )>S(\mathrm {p} )>S(\mathrm {d} )>S(\mathrm {f} ).}

            3.Periodicity of electron configuration:

                   Across a Period: Electronegativity increases because the number of protons increases, pulling electrons more strongly as the shielding effect remains constant.

                   Down a Group: Electronegativity decreases because the atomic size increases, leading to more shielding, which reduces the nucleus’s ability to attract electrons.

 

Coulomb’s Law:

Coulomb’s Law describes the force between two charged objects. It states that the electrostatic force (F) between two point charges is directly proportional to the product of the magnitudes of the charges and inversely proportional to the square of the distance between them.

Formula:

F=ke⋅∣q1⋅q2∣r2F = k_e \cdot \frac{|q_1 \cdot q_2|}{r^2}

Where:

is the electrostatic force between the two charges.

ke is Coulomb’s constant 

(ke≈8.99×109N⋅m2/C2k_e \approx 8.99 \times 10^9 \, \text{N} \cdot \text{m}^2 / \text{C}^2

and are the magnitudes of the two charges.

is the distance between the charges.

  • The force is attractive if the charges are of opposite signs (negative and positive) and repulsive if they are of the same sign (both positive or both negative).

  • This law helps explain how the distance and magnitude of charges affect the force, which is relevant in understanding the attraction between the nucleus and electrons, influencing electronegativity.   

Role of Electronegativity in Chemical Bonding:

1. Ionic Bonding:

  • Electronegativity Difference: A large difference (typically greater than 1.7) in electronegativity between two atoms results in ionic bonding.
  • How It Happens: The atom with the higher electronegativity (more “electron-hungry”) will attract electrons more strongly, pulling one or more electrons from the other atom. This creates a positive ion (cation) and a negative ion (anion), which are held together by electrostatic forces.
  • Example: In NaCl (sodium chloride), sodium (Na) has a low electronegativity and loses an electron to chlorine (Cl), which has a high electronegativity and gains the electron, forming an ionic bond.

2. Covalent Bonding:

  • Electronegativity Difference: A small or moderate difference (typically less than 1.7) in electronegativity results in covalent bonding.
  • How It Happens: Atoms with similar electronegativities share electrons to achieve a stable electron configuration. Both atoms exert similar attractive forces on the shared electrons.
  • Example: In H₂ (hydrogen gas), both hydrogen atoms have the same electronegativity, so they share electrons equally.3. Polar Covalent Bonding:
    • Electronegativity Difference: A moderate difference (between 0.4 and 1.7) in electronegativity results in a polar covalent bond.
    • How It Happens: One atom attracts the shared electrons more strongly than the other, creating a partial negative charge (δ-) on the more electronegative atom and a partial positive charge (δ+) on the less electronegative atom. The electron pair is unequally shared.
    • Example: In H₂O (water), oxygen is more electronegative than hydrogen, so the electrons are pulled toward oxygen, making it partially negative and hydrogen partially positive.

Effect on Molecular Polarity:

        • Polarity of Molecule: The electronegativity difference within the molecule affects its polarity. If a molecule has polar covalent bonds and the molecule itself is asymmetrical, the molecule will be polar.
        • Example: In H₂O (water), the polar covalent bonds lead to a polar molecule, where the oxygen atom has a partial negative charge, and the hydrogen atoms have partial positive charges.

Exceptions to Electronegativity Trends :

          1. Noble Gases:

            • Reason: Do not have a defined electronegativity because they are chemically inert, with a full valence shell and rarely form bonds.

          2. Transition Metals:

            • Reason: Electronegativity trends are irregular due to the involvement of d-orbitals in bonding, which behave differently from s- and p-orbitals.

          3. Lanthanides and Actinides:

            • Reason: These elements have f-orbitals, which complicate electron sharing and cause deviations from the general trend.

          4. Beryllium (Be):

            • Reason: Beryllium has a higher electronegativity than magnesium due to its smaller atomic size and higher effective nuclear charge, allowing it to attract electrons more strongly.

          5. Fluorine (F):

            Reason: While the most electronegative, fluorine’s small size and high electron density cause polarizability effects, which can influence its electron-accepting behavior in certain bonds.

Other Factors Affecting Electronegativity:

            • Atomic Size: Smaller atoms (e.g., Fluorine) are more electronegative than larger atoms (e.g., Iodine).
            • Effective Nuclear Charge: Atoms with more protons (e.g., Oxygen) have higher electronegativity.
            • Electron Shielding: More shielding (e.g., Sodium) lowers electronegativity.
            • Electron Configuration: Atoms with nearly full valence shells (e.g., Chlorine) are more electronegative.
            • Ionization Energy: High ionization energy (e.g., Helium) correlates with high electronegativity.
            • Electron Affinity: High electron affinity (e.g., Chlorine) leads to high electronegativity.
                •  

2.1.A.2  Nonpolar Covalent Bonds: Shared Electrons Between Atoms of Similar Electronegativity:

1.Electronegativity:

1.Definition:

  • Electronegativity is the ability of an atom to attract electrons in a chemical bond.

2.Trends Across the Periodic Table:

  1. Across a Period (Left to Right):

    • Increases as you move from left to right.
    • Reason: Increased nuclear charge (more protons) and constant electron shielding attract bonding electrons more strongly.

  2. Down a Group (Top to Bottom):

    • Decreases as you move down.
    • Reason: More electron shells (greater distance from the nucleus) and increased shielding reduce the nucleus’ pull on bonding electrons.

3.Points to Keep in Mind:

  • Most electronegative element: Fluorine (F) with a value of 4.0.
  • Least electronegative element: Cesium (Cs) and Francium (Fr), with values around 0.7.
  • Electronegativity scale: Ranges from 0.7 (Cs) to 4.0 (F).

4.Electronegativity and Bond Types:

  • Large difference (≥ 1.7): Ionic bond (electrons transferred).
  • Moderate difference (0.4 to 1.7): Polar covalent bond (uneven sharing of electrons).
  • Small/no difference (0 to 0.4): Nonpolar covalent bond (even sharing of electrons).

5.Electronegativity and Reactivity:

  • High electronegativity (e.g., F, O, N): Highly reactive (strong attraction for electrons).
  • Low electronegativity (e.g., alkali metals): Highly reactive (tend to lose electrons).

2.Covalent Bonding:

  1. Definition:

    • Covalent bonding is the sharing of electrons between two atoms to achieve a stable electron configuration (often an octet).

  2. Electron Sharing:

    • Atoms share electrons to fill their outer electron shell.
    • This sharing allows atoms to achieve a stable configuration (usually 8 electrons in the outer shell, except for hydrogen, which aims for 2).

  3. Types of Covalent Bonds:

    • Single Bond: One pair of electrons is shared (e.g., H₂).
    • Double Bond: Two pairs of electrons are shared (e.g., O₂).
    • Triple Bond: Three pairs of electrons are shared (e.g., N₂).

  4. Bond Strength and Length:

    • Single bonds: Longest and weakest.
    • Double bonds: Shorter and stronger than single bonds.
    • Triple bonds: Shortest and strongest.

  5. Polarity of Covalent Bonds:

    • Nonpolar Covalent Bond: Equal sharing of electrons (e.g., Cl₂).
    • Polar Covalent Bond: Unequal sharing of electrons due to different electronegativities (e.g., H₂O).

  6. Bonding in Simple Molecules:

    • H₂ (Hydrogen): Two hydrogen atoms share one pair of electrons.
    • O₂ (Oxygen): Two oxygen atoms share two pairs of electrons.
    • N₂ (Nitrogen): Two nitrogen atoms share three pairs of electrons.

3.Bond Polarity: Difference Between Nonpolar and Polar Covalent Bonds:

AspectNonpolar Covalent BondPolar Covalent Bond
DefinitionA bond formed when two atoms share electrons equally.A bond formed when two atoms share electrons unequally.
Electronegativity Difference< 0.40.4 to 1.7
Electron SharingElectrons are shared equally between atoms.Electrons are shared unequally; the more electronegative atom attracts electrons more.
Charge DistributionNo separation of charge (no dipole). Both atoms have an equal pull on the electrons.Partial charges form on atoms (δ+ and δ-), creating a dipole moment.
Partial ChargesNo partial charges on the atoms.One atom becomes slightly negative (δ-), the other becomes slightly positive (δ+).
ExamplesH₂, O₂, Cl₂H₂O, HF, NH₃
Molecule SymmetryMolecule is typically symmetrical, with no dipole moment.Molecule is usually asymmetrical, resulting in a dipole moment.
Bond StrengthGenerally weaker, as the electron sharing is equal.Stronger due to unequal electron sharing, creating a more polarized bond.
Polarity of MoleculeThe entire molecule is nonpolar (no net dipole).The molecule is polar due to the dipole moment created by unequal electron sharing.

4.Carbon-Hydrogen Bonding:

1. Electronegativity Difference in C-H Bond:

  • Carbon (C): Electronegativity = 2.55
  • Hydrogen (H): Electronegativity = 2.20
  • Electronegativity Difference: 0.35 (small difference)

This small difference results in a nonpolar covalent bond, where electrons are shared nearly equally between carbon and hydrogen.

2. Electron Sharing in C-H Bond: 

  • Almost Equal Sharing: The small electronegativity difference (0.35) leads to nearly equal electron sharing between C and H.
  • Nonpolar Bond: This equal sharing results in a nonpolar covalent bond with no significant dipole.
  • Minimal Electron Pull: Carbon pulls slightly more, but the difference is too small to create a noticeable polarity.

3. Bond Behavior in C-H Bond:

  • Nonpolar Behavior: Despite a slight electronegativity difference, the bond acts nonpolar because electron sharing is nearly equal.
  • Negligible Dipole Moment: The small dipole moment created is insignificant and effectively negligible.
  •  

  • Equal Electron Sharing: The near-equal sharing of electrons means no significant charge separation, keeping the bond nonpolar.

5. Effect on Molecule:

  • No Significant Dipole Moment: C-H bonds share electrons almost equally, creating no significant dipole in the bond.
  • Nonpolar Organic Molecules: Molecules with many C-H bonds, like hydrocarbons (e.g., methane, ethane), are typically nonpolar because individual C-H bonds don’t create an overall dipole.
  • Cancellation of Polarity: Multiple C-H bonds don’t contribute to polarity as their effects cancel out, keeping the molecule nonpolar.
  • Impact on Properties: Nonpolar molecules tend to dissolve in nonpolar solvents and are poor in polar solvents (like water), affecting their solubility and intermolecular interactions.
  • Exceptions in Larger Molecules: Larger organic molecules with functional groups (e.g., -OH, -NH₂) can be polar, even if C-H bonds remain nonpolar.
  •  

 

 

2.1.A.3 Polar Covalent Bonds: Formation, Charge Distribution, and the Continuum Between Ionic and Covalent Bond:

  •  Polar Covalent Bonds: Unequal Electronegativity and Shared Valence Electrons:

    1. Electronegativity:

    • Electronegativity: The ability of an atom to attract electrons in a bond.
    • Trend: Increases across a period, decreases down a group.

    2. Polar Covalent Bond Formation:

    • Occurs when atoms with different electronegativities share electrons unequally.
    • Creates partial charges: δ+ (positive) and δ− (negative).

    3. Dipole Moment:

    • Dipole: A molecule with a positive and a negative end due to unequal electron sharing.
    • A polar molecule has a net dipole.

    4. Molecular Geometry:

    • Asymmetrical shapes (e.g., H₂O) result in polar molecules.
    • Symmetrical shapes (e.g., CO₂) lead to nonpolar molecules despite polar bonds.

    5. Properties of Polar Molecules:

    • Solubility: Polar molecules dissolve in polar solvents (e.g., water).
    • Boiling/Freezing Points: Higher due to stronger intermolecular forces.
    • Conductivity: Polar molecules can conduct when dissolved (e.g., NaCl in water).

    6. Examples:

    • Polar: H₂O, HF, NH₃.
    • Nonpolar: O₂, CO₂, CH₄.

i. The atom with a higher electronegativity will develop a partial negative charge relative to the other atom in the bond:

In a covalent bond, the more electronegative atom attracts electrons, creating a partial negative (δ−) charge, while the less electronegative atom has a partial positive (δ+) charge. For example, in water (H₂O), oxygen is δ− and hydrogen is δ+.

1.Electronegativity: It’s the ability of an atom to attract electrons in a bond. Higher electronegativity means stronger attraction.

2.Unequal Electron Sharing:

In polar covalent bonds, electrons are shared between two atoms, but the sharing is unequal due to differences in electronegativity. Electronegativity refers to an atom’s ability to attract and hold onto electrons when it forms a bond with another atom.

  • Higher electronegativity means an atom has a stronger attraction for electrons.
  • Lower electronegativity means an atom has a weaker attraction for electrons.

    3.Unequal Electron Sharing:

    In a bond between atoms with different electronegativities, the more electronegative atom pulls the shared electrons closer, creating a partial negative charge (δ−).

    4.Partial Charges:

    The less electronegative atom develops a partial positive charge (δ+) due to having fewer electrons.

    5.Dipole Formation:

    The separation of charges (δ+ and δ−) creates a dipole, giving the molecule a positive and a negative end, which affects its properties (e.g., polarity, solubility).

    ii.In single bonds, greater differences in electronegativity lead to greater bond dipoles:

    The bond dipole moment measures bond polarity and depends on the electronegativity difference between atoms. A larger difference results in a stronger dipole. In HCl, chlorine (Cl) is more electronegative than hydrogen (H), so the bond is polarized, with chlorine being δ− and hydrogen being δ+.

    1.Covalent Bonding:

    A covalent bond forms when two atoms share electrons to achieve a stable electron configuration, usually following the octet rule. This typically occurs between nonmetals.

    • Atoms share electrons to reach a stable, low-energy state (like noble gases).
    • Single bonds share one pair of electrons, while double and triple bonds share two or three pairs, respectively. A Lewis dot structure is a diagram that shows how atoms in a molecule are bonded and the arrangement of valence electrons, using dots for electrons and lines for bonds.
      Covalent BondsNonpolar Covalent BondsPolar Covalent Bonds
      Electron SharingElectrons shared equally or nearly equallyElectrons shared unequally
      Electronegativity Difference0 to 0.40.4 to 1.7
      Partial ChargesNo partial charges (δ+ or δ−)Partial charges (δ+ and δ−)
      Electron DistributionSymmetrical electron distributionAsymmetrical electron distribution
      Molecule PolarityMolecule is nonpolarMolecule is polar
      Examples– H₂ (Hydrogen)– H₂O (Water)
       – Cl₂ (Chlorine)– HCl (Hydrogen chloride)

      iii. All polar bonds have some ionic character, and the difference between ionic and covalent bonding is not distinct but rather a continuum:

    The distinction between ionic and covalent bonds is a spectrum. In ionic bonds, electrons are fully transferred (e.g., NaCl), while in covalent bonds, electrons are shared. Polar covalent bonds involve unequal sharing due to electronegativity differences. As the electronegativity gap increases, the bond becomes more ionic, but there’s no sharp boundary. For example, Na-Cl is more ionic than H-Cl, yet both share electrons to some degree.

    1.Ionic vs. Covalent Bonds:

    PropertyIonic BondsCovalent Bonds
    Electron Transfer/SharingComplete electron transferElectrons are shared
    ElectronegativityLarge difference (>1.7)Small difference (<1.7)
    Bond TypeBetween metals and nonmetals (e.g., NaCl)Between nonmetals (e.g., H₂O)
    Ion FormationForms positive and negative ionsNo ions, partial charges in polar bonds
    Bond StrengthStrong due to electrostatic attractionWeaker, varies with atoms
    ConductivityConducts in molten or aqueous formDoes not conduct
    Melting/Boiling PointsHigh (e.g., NaCl: 801°C)Lower (e.g., H₂O: 100°C)
    SolubilitySoluble in waterVaries; nonpolar in nonpolar solvents
    ExamplesNaCl, MgOH₂O, CO₂, HCl

    2.Bond Character Continuum:

    The bond character continuum refers to the idea that bonding is not strictly ionic or covalent, but exists on a spectrum. As the electronegativity difference between atoms increases, the bond becomes more ionic in nature.

    • Ionic bonds occur when there is a large electronegativity difference (typically >1.7), where one atom transfers electrons to the other, forming ions (e.g., NaCl).
    • Polar covalent bonds occur with a moderate electronegativity difference (0.4 to 1.7), where electrons are unequally shared, creating partial charges (e.g., H₂O).
    • Nonpolar covalent bonds occur with a small electronegativity difference (<0.4), where electrons are shared equally (e.g., Cl₂).
    •  

2.1.A.4   Factors Influencing Bond Type: Electronegativity and Properties:

The difference in electronegativity is not the only factor in determining whether a bond is ionic or covalent. Generally, bonds between metals and nonmetals are ionic, while bonds between two nonmetals are covalent.

  • Ionic Bonds and Covalent Bonds:
    AspectIonic BondsCovalent Bonds
    Characteristics– Electron transfer from metal to nonmetal– Electron sharing between nonmetals
     – Formation of ions (cation and anion)– Formation of a stable electron configuration
    Metal-Nonmetal Bonding– Metals lose electrons to form cations (positive ions)– Nonmetals share electrons to form bonds
     – Nonmetals gain electrons to form anions (negative ions)– Bonding typically occurs between two nonmetals
    Examples of Compounds– Sodium chloride (NaCl)– Water (H₂O), Carbon dioxide (CO₂)
     – Magnesium oxide (MgO)– Nitrogen gas (N₂), Oxygen gas (O₂)
    Properties– High melting and boiling points– Lower melting and boiling points
     – Hard, brittle crystals– Soft, flexible structures
     – Good electrical conductivity when molten or dissolved– Poor electrical conductivity (non-electrolytes)
     – Soluble in polar solvents (e.g., water)– Poor solubility in water
    Bond Types– One type: Ionic bonds (electrostatic attraction between ions)– Single, double, and triple bonds (based on number of electron pairs shared)
    Bonding Between Elements– Metal and nonmetal– Nonmetal and nonmetal
  • Electronegativity Difference and Bond Character:  The electronegativity difference between two atoms plays a crucial role in determining the type of bond they form. It influences how electrons are shared or transferred, affecting whether a bond is ionic, polar covalent, or nonpolar covalent.

    Electronegativity Difference Ranges:

    1. Nonpolar Covalent Bond (0 to 0.4)

      • Electrons Shared Equally: Atoms have similar electronegativities, so electrons are shared equally.Chlorine - Wikipedia
      • Example: Cl₂ (Chlorine).

    2. Polar Covalent Bond (0.4 to 1.7)

      • Unequal Sharing: The more electronegative atom pulls electrons more strongly, creating partial charges.Chemical polarity - Wikipedia
      • Example: H₂O (Water).

    3. Ionic Bond (> 1.7)

      • Electron Transfer: One atom transfers electrons to the other, forming ions.File:NaCl Formation Lewis.svg - Wikimedia Commons
      • Example: NaCl (Sodium chloride).
         

        Electronegativity Difference and Its Effect on Bond Character:

        The electronegativity difference between two atoms plays a crucial role in determining the type of bond they form. It influences how electrons are shared or transferred, affecting whether a bond is ionic, polar covalent, or nonpolar covalent.

        Electronegativity Difference Ranges:

        1. Nonpolar Covalent Bond (Difference: 0 to 0.4)

          • Electrons Shared Equally: When two atoms have a very similar or equal electronegativity, electrons are shared equally, resulting in a nonpolar covalent bond.
          • Example: Cl₂ (Chlorine molecule). Both chlorine atoms have the same electronegativity, so electrons are shared equally.

        2. Polar Covalent Bond (Difference: 0.4 to 1.7)

          • Unequal Sharing of Electrons: If the electronegativity difference is moderate, the more electronegative atom will attract the shared electrons more strongly, creating partial charges (δ+ and δ−).
          • Example: H₂O (Water). Oxygen is more electronegative than hydrogen, so the shared electrons are pulled more toward oxygen, making the bond polar.

        3. Ionic Bond (Difference: > 1.7)

          • Electron Transfer: A large electronegativity difference leads to complete electron transfer, with one atom becoming negatively charged (anion) and the other positively charged (cation). This creates an ionic bond.
          • Example: NaCl (Sodium chloride). Sodium (Na) has a low electronegativity and transfers its electron to chlorine (Cl), forming Na⁺ and Cl⁻ ions.

        Effect of Electronegativity Difference:

        Electronegativity DifferenceBond Character
        0Nonpolar Covalent
        0.4Polar Covalent
        1.7+Ionic
      • Macroscopic properties relate to bond types:
        PropertyIonic BondsPolar Covalent BondsNonpolar Covalent Bonds
        Melting/Boiling PointsHigh (strong electrostatic forces between ions)Moderate (dipole-dipole or hydrogen bonding)Low (weak London dispersion forces)
        SolubilitySoluble in polar solvents (e.g., water)Soluble in polar solvents (e.g., water)Soluble in nonpolar solvents (e.g., oils)
        ConductivityConductive in molten state or aqueous solutionNon-conductive (no free ions or electrons)Non-conductive (no free ions or electrons)
  • Ionic bonds lead to high melting/boiling points, solubility in polar solvents, and conductivity in molten or aqueous form.
  • Polar covalent bonds lead to moderate melting/boiling points, solubility in polar solvents, and non-conductivity.
  • Nonpolar covalent bonds lead to low melting/boiling points, solubility in nonpolar solvents, and non-conductivity.
    •  

2.1.A.5   Delocalized Valence Electrons in Metallic Solids:

1. Metallic Bonding: Metallic bonding occurs in metals when atoms lose their valence electrons, which become delocalized and move freely around a lattice of metal cations. This creates a “sea” of electrons that holds the metal atoms together.

2.Electron sea model:

The electron sea model explains the structure and behavior of metals:File:Sea of electrons.jpg - Wikipedia

  • Delocalized Electrons: In metals, the valence electrons are not attached to specific atoms. Instead, they form a “sea” of free-moving electrons that flow through the metal lattice.

  • Metal Lattice: The metal consists of a regular arrangement of positive metal ions (cations) surrounded by this “sea” of delocalized electrons.

  • Properties:

    • Electrical Conductivity: The free electrons can move easily, allowing metals to conduct electricity.
    • Thermal Conductivity: The mobile electrons transfer heat quickly through the metal.
    • Malleability: The “sea” of electrons allows metal atoms to slide past each other without breaking the bond, making metals flexible and able to be shaped.

3.Structure of metallic solids: The arrangement of metal atoms in a regular, closely packed structure, often in the form of body-centered cubic (BCC), face-centered cubic (FCC), or hexagonal close-packed (HCP) lattices.

                                                                                       

BCC                                                                    FCC                                                                  HCP

2. Delocalized Electrons:

  Valence Electrons in Metals:

  • Electron Loss: In metals, the valence electrons (outermost electrons) are loosely held by individual atoms due to the low ionization energy.

  • Delocalization: These valence electrons are freed from their parent atoms and become delocalized, meaning they can move freely throughout the metal’s structure.

  • “Electron Sea”: The free-moving electrons form a “sea” of electrons that surrounds the positively charged metal ions (cations) in the metal lattice.

  • Electron mobility:

    • Electrical Conductivity: Delocalized electrons move freely, allowing metals to conduct electricity.
    • Thermal Conductivity: Electrons transfer heat efficiently through the metal.
    • Malleability and Ductility: Electrons enable atoms to slide past each other, allowing metals to be shaped without breaking.

    In short, electron mobility gives metals their ability to conduct electricity, heat, and remain flexible.

  • Electrical Conductivity in Metals:

    • Delocalized electrons move freely in the metal lattice.
    • When an electric field is applied, these electrons flow, creating electric current.
    • This movement of electrons allows metals to conduct electricity efficiently.

    In short, delocalized electrons enable metals to conduct electricity by easily flowing when a voltage is applied.

    3. Properties of Metals:

  • Electrical Conductivity:

    • Metals have delocalized electrons that can move freely within the lattice, allowing them to conduct electricity efficiently.

  • Malleability and Ductility:

    • Due to the flexible electron sea, metal atoms can slide past one another without breaking the structure, making metals malleable (able to be hammered into sheets) and ductile (able to be stretched into wires).

  • Luster:

    • The shiny appearance of metals is caused by the reflection of light off the surface, due to the delocalized electrons interacting with light.

4. Metallic vs. Ionic and Covalent Bonding:

  • Metallic Bonding

    • Electron Behavior: Electrons are delocalized and move freely in the metal.
    • Result: The “sea” of electrons holds metal ions together, creating conductivity, flexibility, and luster.
    • Metallic bonds are found in metals like zinc.

  • Ionic Bonding

    • Electron Behavior: Electrons are transferred from one atom to another.
    • Result: Charged ions (cations and anions) are formed, held together by electrostatic forces.

  • Covalent Bonding

    • Electron Behavior: Electrons are shared between atoms.
    • Result: Molecules are formed, with a more localized electron distribution.

5. Energy and Strength of Metallic Bonds:

AspectExplanation
Bond StrengthThe number of delocalized electrons influences the strength and melting points of metals. More delocalized electrons lead to stronger bonds.
Alloy FormationAlloys are created by mixing different metals, which changes the electron sea and modifies the properties of the resulting material.
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