Understanding Chemical Bonds

Chemical Bonds

• Chemical bonds hold atoms together, allowing them to behave as a single unit with unique properties.
• Chemical bonds are crucial for determining a substance's physical properties and the course of chemical reactions.
• Chemical reactions involve the breaking of old bonds and the formation of new ones.

Types of Chemical Bonds

• Ionic bonds typically form between a metal and a non-metal.
• Covalent bonds usually form between two non-metals, and can be polar or nonpolar.
• Metallic bonds occur between metallic cations (metal and metal).

Electronegativity

• Electronegativity generally increases from left to right across the periodic table due to increasing effective nuclear charge.
• Electronegativity generally decreases from top to bottom down the periodic table because of increasing energy levels/shells.
• Large electronegativity differences between atoms favor ionic bond formation.
• Small electronegativity differences between atoms favor covalent bond formation.
• Exact electronegativity values are not required to be memorized for AP Chemistry, only the periodic trends

Ionic Bonds

• Ionic bonds form through the transfer of electrons between atoms.
• Electron transfer results in the formation of a cation (positive ion) and an anion (negative ion).
• The electrostatic attraction between cations and anions constitutes an ionic bond.
• Sodium chloride (NaCl) and magnesium oxide (MgO) are examples of ionic compounds.

Ionic Compounds

• Ionic compounds exist as crystal solids with a highly ordered crystal lattice structure.
• Ions are fixed in position within the crystal lattice.
• The crystal lattice structure contributes to the rigidity and brittleness of ionic compounds.
• Ionic solids are rigid due to strong ionic bonds throughout the structure, resisting compression.
• Ionic solids are brittle because applied stress can cause ions of like charge to align, leading to repulsion and fracture.
• Ionic compounds have high melting and boiling points due to the large amount of energy required to break their strong ionic bonds.
• Ionic compounds are poor conductors of electricity in the solid state due to the lack of mobile ions.
• Ionic compounds are good conductors of electricity in the liquid or aqueous phase where ions are free to move.

Coulomb's Law

• Coulomb's Law (F = k * q1 * q2 / r^2) describes the strength of the force between charged particles.
• F is the force between the charges
• k is Coulomb's constant
• q1 and q2 represent the magnitude of the charges
• r is the distance between the centers of the charges
• Greater charges lead to greater attractive or repulsive forces.
• Increasing the distance between charges decreases the force between them.
• Magnesium oxide (MgO) has a stronger ionic bond than sodium chloride (NaCl) due to the higher charges of the ions and smaller ionic radii.

Ionic vs. Covalent Bonds

• Magnesium oxide (MgO) has stronger ionic bonds than sodium chloride (NaCl) due to a greater charge and smaller distance.
• MgO's stronger bonds lead to a higher melting point, better electrical conductivity, and increased brittleness and rigidity compared to NaCl.
• Covalent bonds form through the sharing of electrons between atoms, unlike the transfer of electrons in ionic bonds.
• Nonpolar covalent bonds involve equal sharing of electrons between identical atoms or atoms with similar electronegativity.
• Polar covalent bonds result from unequal sharing of electrons, with the more electronegative atom attracting electrons more strongly.
• In a polar covalent bond, the more electronegative atom gains a partial negative charge (δ-), while the less electronegative atom gains a partial positive charge (δ+).
• Unlike ionic bonds where electrons are completely transferred, covalent bonds involve electrons moving back and forth between atoms.
• In nonpolar covalent bonds, electron clouds are symmetrical, while in polar covalent bonds, the electron cloud is larger around the more electronegative atom.
• Binary compounds composed of elements with the same electronegativity are likely to form nonpolar covalent bonds.

Covalent Bond Length and Strength

• Covalent bond length is determined by the potential energy versus internuclear distance diagram.
• When two atoms are far apart, there is no interaction and potential energy is zero.
• When two nuclei are too close together, there is strong repulsive force and high potential energy, creating an unstable state.
• At a certain distance, attraction between the nucleus and electrons balances the repulsive force between nuclei.
• The balance point represents the most stable state, with the lowest potential energy determining the covalent bond length.
• For hydrogen gas (H2), the balance point is 74 picometers.
• Single, double, and triple bonds are classified by the number of electrons shared between two atoms.
• Single bonds share two electrons, double bonds share four electrons, and triple bonds share six electrons.
• As the number of shared electrons increases from single to triple bonds, bond length decreases and bond strength increases.
• Double bonds are stronger and shorter than single bonds.

Electron Cloud Model and Bonding

• The electron cloud model is a modern atomic structure model where electrons move randomly in a probabilistic manner.
• Electron clouds depict the likelihood of finding electrons around the nucleus, with denser regions indicating higher probability.
• Sharing of electrons in covalent bonds is represented by the overlap of electron clouds.
• Sigma (σ) bonds occur when orbitals overlap head-to-head along the internuclear axis.
• Sigma bonds can form between two s orbitals, two p orbitals, or an s orbital and a p orbital.
• Pi (π) bonds occur only when two p orbitals overlap while parallel to each other.
• Pi bonds have electron cloud overlap above and below the internuclear axis.
• Single bonds consist of one sigma bond, double bonds consist of one sigma bond and one pi bond, and triple bonds consist of one sigma bond and two pi bonds.

Metallic Bonds

• Metals are solid, hard, and shiny, composed of metallic cations and delocalized electrons.
• Metallic cations form a lattice structure surrounded by a "sea" of delocalized electrons, which are free to move.
• Delocalized electrons make metals good conductors of electricity in liquid, gas, or solid phases.
• Metallic bonds are the attraction between metallic cations and delocalized electrons.
• Metals are malleable (can be stretched into thin sheets) and ductile (can be stretched into thin wires) due to the electron sea model.
• Freely moving electrons allow metallic cations to glide around each other without breaking bonds when compressed or stretched.

Alloys

• An alloy is a mixture of different metals, sometimes including nonmetals.
• Alloys are often harder, stronger, less malleable, and less ductile than pure metals.
• Steel is an alloy of iron (Fe) and carbon (C), with carbon increasing its strength.
• Interstitial alloys form when smaller atoms fill spaces between larger atoms in the lattice structure.
• Substitutional alloys form when atoms of similar size replace each other in the lattice structure.
• Carbon atoms in steel fill spaces between iron cations, increasing rigidity by blocking movement.

Lewis Dot Diagram

• It illustrates how atoms are bonded within a molecule, indicating single or double bonds, such as in carbon dioxide (CO2).
• It is important for determining the three-dimensional shape of a molecule.
• In combination with VSEPR theory, Lewis dot diagrams are inportant for understanding Intermolecular forces in unit 2 and unit 3.
• It adheres to rules, including valence electron count, where the diagram cannot exceed the total valence electrons available.
• Achieving a full octet state, with eight valence electrons around each atom (two for hydrogen/helium), is another rule.
• To make sure that each oxygen achieves the full octet state, then each of the oxygen should be able to have 2 times 8 which is 16 valence electrons in total.
• Determine the number of shared electrons and calculate the electrons and the number of lone pair electrons.
• Each oxygen has six valence electrons, and ends up carrying four lone pair electrons, as well as sharing 4 electrons between the two Oxygen atoms.
• One must Check that the Louis dot diagram is correct by counting the electrons.

Nitrogen Gas and Electron Sharing

• Nitrogen atoms share six electrons to achieve a total of 16 electrons in the molecule.
• After sharing six electrons, each nitrogen atom has two lone pair electrons.
• The structure consists of a nitrogen-nitrogen triple bond.
• The nitrogen gas Lewis dot diagram fulfills the octet rule with each nitrogen atom having eight valence electrons.

Formal Charge Calculation

• Formal charge helps determine the most appropriate Lewis dot diagram when multiple possibilities exist.
• Formal charge is calculated as valence electrons minus lone pair electrons minus (bonding pair electrons divided by two).

Phosphate Example (PO43-)

• Two possible Lewis dot diagrams for phosphate include phosphorus bonded to four oxygen atoms via single bonds, and phosphorus bonded to one oxygen via a double bond and three oxygens via single bonds.
• In the first Lewis dot diagram:
  • Phosphorus has a formal charge of +1.
  • Each oxygen has a formal charge of -1.
• In the calculation for phosphorus formal charge:
  • Valence electrons: 5
  • Lone pair electrons: 0
  • Bonding pair electrons: 8, leading to a formal charge calculation of 5 - 0 - (8/2) = +1.
• In the calculation for oxygen formal charge:
  • Valence electrons: 6
  • Lone pair electrons: 6
  • Bonding pair electrons: 2, leading to a formal charge calculation of 6 - 6 - (2/2) = -1.

Formal Charge Calculation

• The sum of formal charges in a polyatomic ion equals the ion's overall charge.
• Phosphorus (P) with 5 valence electrons, 0 lone pair electrons, and 10 bonding electrons has a formal charge of 0.
• Oxygen (O) with 6 valence electrons, 6 lone pair electrons, and 2 bonding electrons has a formal charge of -1.
• Oxygen with 6 valence electrons, 4 lone pair electrons, and 4 bonding electrons has a formal charge of 0.

Evaluating Lewis Structures

• Structures with formal charges closest to zero are more appropriate.
• A formal charge of zero on atoms is generally preferred for stability.
• In cases where formal charges cannot be minimized to zero, electronegativity differences between atoms become important.
• More electronegative atoms are better suited to carry negative formal charges.
• Nitrogen is more electronegative than sulfur and is therefore more likely to carry a negative formal charge.

Lewis Structures and Octet Rule

• Diagram 1 is less appropriate for H2O because carbon does not fulfill the octet rule.
• Non-zero formal charges indicate a less appropriate Lewis structure.

Carbon Dioxide Structures

• Diagram 1 of CO2 is more appropriate because all atoms have a formal charge of zero.
• Diagram 2 shows a less appropriate form with charges of +1 and -1 for oxygen atoms.

Lewis Structures of H3NO

• Option A is the best Lewis dot diagram for H3NO because both nitrogen and oxygen have a formal charge of zero.

Resonance Structures

• Multiple possible Lewis structures exist for a single molecule or ion with the same arrangement of atoms (differing only in the distribution of electrons).
• These structures are called resonance structures.
• Resonance structures don't imply switching between forms, but rather a resonance hybrid.
• The resonance hybrid is an average of all resonance structures.
• In ozone (O3), the two oxygen-oxygen bonds are identical in length in nature, indicating a resonance hybrid with a bond order of 1.5.
• Carbonate (CO32-) has three resonance structures where a double bond can be between carbon and any of the three oxygen atoms.
• The actual structure is a resonance hybrid with each carbon-oxygen bond having a bond order of four-thirds, and each oxygen atom carrying a charge of -2/3.

Octet Exceptions: Underfilled and Overfilled Octets

• Underfilled octets occur when a central atom (typically from groups 2 or 3) has fewer than 8 valence electrons.
• Beryllium dihydride (BeH2) and Boron trichloride (BCl3) are examples of molecules with underfilled octets that are considered stable.
• Overfilled octets occur when a central atom (typically from groups 5, 6, 7, or 8) has more than 8 valence electrons.
• Xenon tetrafluoride (XeF4) and phosphorus pentachloride (PCl5) are examples.
• These compounds are exceptions to the octet rule but are still considered stable.

VSEPR Theory

• VSEPR (Valence Shell Electron Pair Repulsion) theory helps determine the 3D shape of molecules by minimizing electron repulsion.
• Determine the Lewis structure first, then apply VSEPR theory.
• Lewis diagrams only show atom arrangement.
• Molecular geometry refers to the 3D shape of a molecule.
• Electron pair includes both lone pairs and bonding pairs of electrons.

Electron Pair Arrangements & Molecular Geometry

• Central atoms with two electron pairs (MX2) adopt a linear molecular geometry to minimize repulsion. Carbon dioxide (CO2) is an example.
• Three electron pairs (MX3) lead to a trigonal planar electron domain geometry with 120 degree angles. Boron trichloride (BCl3) is an example.
• With one lone pair, the molecular geometry with three electron pairs becomes bent (e.g., ozone).
• Four electron pairs (MX4) result in a tetrahedral 3D structure with bond angles of 109.5 degrees. Methane (CH4) is an example of tetrahedral.
• With one lone pair, the four electron pair arrangement becomes trigonal pyramidal (e.g., ammonia, NH3).
• With two lone pairs, the four electron pair arrangement becomes bent (e.g., water, H2O).
• Five electron pairs (MX5) give a trigonal bipyramidal arrangement.
• Six electron pairs (MX6) form an octahedral molecular geometry where all electron pairs are 90° apart.

Molecular Polarity Criteria

• Examine the molecular geometry and bond polarity.
• Asymmetric molecular geometries mean the molecule is polar.
• If molecular geometry is symmetrical, examine the polarity of the individual bonds.
• If bonds are not balanced, the molecule is polar.

Molecular Polarity Examples

• $BCl_3$ is nonpolar due to its symmetrical trigonal planar shape and identical bonds.
• $CO_2$ is nonpolar because it is a linear shape and each of the bond is the same
• $CH_2Cl_2$ is polar, due to its asymmetrical shape, 2 bonds are chlorine, 2 bonds are hydrogen.
• $N_2$ is nonpolar due to is a symmetrical linear shape.

Hybridization

• Hybridization is the mixing of atomic orbitals to form new hybrid orbitals with equal energy levels.
• Number of electron pairs determines hybridization type.
  • Two electron pairs: sp hybridization
  • Three electron pairs: sp2 hybridization
  • Four electron pairs: sp3 hybridization

Nitrogen Gas and Energy Graphs

• Two nitrogen atoms = 10 valence electrons
• Lewis structures of Nitrogen, $N_2$: $N≡N$
• $N_2$ is nonpolar
• Point on the potential energy graph is where peak potential energy is lowest (or minimum)
• Triple bonds are stronger and shorter than double bonds.
• More energy is required to break a triple bond than a double bond.
• In hydrazine $(N_2H_4)$, the single $N-N$ bond is longer and weaker compared to the bond in $N_2$.
• Each nitrogen atom will have $sp^3$ hybridization.