Types of Bonding and Properties of Ionic Compounds

BONDS

WHAT ARE THE TYPES OF BONDING?

  • Bonding Types: Different types of bonds are formed depending on the types of atoms involved. All bonds involve electrons, and bonding includes changes to the outer shell electron numbers of atoms.

    • Ionic Bonding

    • Covalent Bonding

    • Metallic Bonding

WHY DO ATOMS FORM BONDS?

  • Atoms and Stability: Noble gases have completely full outer shells, making them stable and unreactive, while other element atoms have incomplete outer electron shells, rendering them unstable.

  • Formation of Bonds: Atoms bond to fill their outer shells and achieve stability. Noble gases typically do not form bonds due to their stable nature.

IONIC BONDING

  • Reactive Metal Atoms: Reactive metal atoms become stable positive ions by losing electrons.

  • Ionic Bond Example:

    • Sodium

    • Na loses one electron to form a positive ion: Na+Na^+

    • Chlorine gains one electron to form a negative ion: Cl−Cl^-

  • Electrostatic Attraction: The bond formed between a cation (positively charged ion) and an anion (negatively charged ion) due to the strong electrostatic attraction is termed ionic bonding.

IONS AND ATTRACTION

  • Ion Attraction and Repulsion: Some ions attract each other (e.g., Sodium
    Na and Chlorine Cl) due to opposite charges, while like charges repel.

  • Example Phrase:

    • “Will you accept my electron forever and make an ionic bond with me?”

HOW IS A SODIUM ION FORMED?

  • Electron Configuration of Sodium:

    • Sodium (Na) has:

    • Atomic Structure: 2.8.12.8.1

    • 11 Protons: +11

    • 11 Electrons: -11

    • Total Charge: 0 (neutral sodium atom)

  • Ion Formation: After losing 1 electron, sodium ion structure changes:

    • Na+:2.8Na^+: 2.8 (full outer shell)

HOW IS A FLUORIDE ION FORMED?

  • Electron Configuration of Fluorine:

    • Fluorine (F) has:

    • Atomic Structure: 2.72.7

    • 9 Protons: +9

    • 9 Electrons: -9

    • Total Charge: 0 (neutral fluorine atom)

  • Ion Formation: After gaining 1 electron, fluoride ion structure changes:

    • F−:2.8F^-: 2.8 (full outer shell)

WHAT IS IONIC BONDING?

  • Ionic Compounds: Compounds containing ions are referred to as ionic compounds, formed when metal atoms transfer electrons to non-metal atoms, resulting in the formation of ions with filled outer shells.

  • Formation Overview:

    • Metal atoms lose electrons and non-metal atoms gain electrons.

    • The electrostatic attraction between cations and anions results in ionic bonding.

COVALENT BONDING

HOW ARE COVALENT BONDS FORMED?

  • Non-Metal Bonds: Non-metals tend to need just one or two electrons to fill their outer shells.

  • Covalent Bond Definition: A bond formed when atoms share electrons is called a covalent bond, resulting in each atom having a full, stable outer shell.

IONIC VERSUS COVALENT BONDING

  • Distinction in Bonding:

    • Electronegativity: The atom's ability to attract the shared pair of electrons.

    • Non-polar Covalent Bond: Equally shared electrons leading to a balanced electrical charge.

    • Polar Covalent Bond: Unequally shared electrons leading to an uneven distribution of electrical charge.

ELECTRONEGATIVITY AND BONDING CHARACTER

  • Difference usage in Bonding Classification:

    • Use electronegativity differences to identify bonding types between sulfur (S) and elements like hydrogen (H), cesium (Cs), and chlorine (Cl).

REVIEWING MAIN IDEAS

  1. Ionic vs. Covalent: What is the main distinction between ionic and covalent bonding?

  2. Electronegativity in Bonding: How does electronegativity determine bonding character?

  3. Predicting Bond Types: Expected bonding types between given pairs of atoms:

    • a. Li and F

    • b. Cu and S

    • c. I and Br

  4. Increasing Ionic Character: List pairs from question 3 in order of increasing ionic character.

CRITICAL THINKING

  1. Comparing Atoms: i. Cu and Cl; ii. I and Cl.

    • a. Which pair has a greater percent ionic character?

    • b. Which pair does Cl have a greater negative charge?

  2. Inferring Ionic Relationships:

    • a. What type of bond is expected between K and Br?

    • b. Which ion is larger in KBr?

PROPERTIES AND STRUCTURE OF IONIC COMPOUNDS

IONIC COMPOUNDS STRUCTURE

  • Ionic Lattice: Millions of ions in a regular cubic arrangement form a giant 3D structure called an ionic lattice, impacting ionic properties.

IONIC COMPOUNDS PROPERTIES

  • **Physical Properties: **

    • High melting and boiling points.

    • Brittleness when fractured.

    • Ability to conduct electricity when dissolved in water or melted.

WHAT IS THE STRUCTURE OF METALS?

  • Metallic Bonding: The attraction between positively charged metal ions and a delocalized sea of electrons is described as metallic bonding.

  • Typical Metal Characteristics:

    • Solid at room temperature

    • High melting points

    • Conducts heat and electricity well

    • Malleable (shapeable)

    • Ductile (drawn into wires)

    • Generally strong and dense.

WRITING FORMULAS FOR IONIC COMPOUNDS

  1. **Steps for Writing Ionic Formulas: **

    • Write down symbols for each element (metal first).

    • Calculate charge for each ion type.

    • Balance the number of ions to achieve a neutral compound.

    • Use the balanced ratio to write the formula.

EXAMPLES OF WRITING FORMULAS

  • Sodium chloride: Na

    • Na++Cl−Na^+ + Cl^- → NaCl

  • Magnesium oxide:

    • Mg2++O2−Mg^{2+} + O^{2-} → MgO

  • Sodium oxide:

    • Sodium needs to lose one electron and oxygen must gain two electrons, yielding the formula Na2ONa_2O .

ELECTRON-DOT STRUCTURE

  • Electron-Dot Notation: Uses dots to represent the valence electrons of an element shown around the element’s symbol.

  • Example:

    • Hydrogen: H:</p></li><li><p>Nitrogen:</p></li><li><p>N..H:</p></li><li><p>Nitrogen:</p></li><li><p>N..

LEWIS STRUCTURES

  1. How to Draw Lewis Structures:

    • Determine the type and number of atoms in the molecule.

    • Assess total valence electrons available.

    • Arrange atoms for skeleton structure (central atom rules).

    • Connect atoms with electron-pair bonds.

    • Add unshared electron pairs to nonmetals for octet formation.

    • Verify electron counts match valence electrons available.

EXCEPTION TO OCTET RULE

  • Specific Cases:

    • BF3 (electron deficient)

    • PCl5 (expanded octet)

FORMULA OF IONIC COMPOUNDS

  • Example Questions:

    • What is the formula of Aluminium Bromide?

    • What is the formula of Aluminium Oxide?

SUMMARY OF METAL PROPERTIES

  • High melting point.

  • Good electrical and heat conductors.

  • Malleable and ductile.

  • Strong and dense.

SHAPES OF MOLECULES

VALENCE SHELL ELECTRON PAIR REPULSION THEORY (VSEPR)

  • Definition: VSEPR theory explains how the geometrical arrangement of atoms in a molecule is determined by the repulsions between electron pairs in the valence shell of the central atom.

    • Covalent Bonds: Molecules consist of covalent bonds, which are formed by pairs of electrons.

    • Repulsive Forces:

    • Bonds will repel each other to minimize repulsive forces.

    • Distance:

      • When bonds are further apart, repulsive forces between them are lesser.

      • When bonds are closer, repulsive forces are greater.

    • Bonds are arranged to be as far apart as possible to reduce these forces.

    • This arrangement results in equally spaced bonds to minimize repulsion.

BOND ANGLES AND MOLECULAR SHAPES

  • Because repulsions between bond pairs are equal, simple molecules (having a central atom with other atoms bonded) typically exhibit standard shapes with equal bond angles.

  • The presence of lone pairs on the central atom alters the angles between bonds and, consequently, the molecular shape.

MOLECULAR GEOMETRY BY ELECTRON DOMAINS

Two Electron Domain Molecules

  • Molecule Example:

    • Carbon Dioxide (=C=O)

    • Structure: Two double bonds around the central atom, Carbon.

    • Shape: Linear

    • Bond Angle: 180.0°

Three Electron Domain Molecules

  • Structure:

    • Example: :

  • Shape: Trigonal planar

    • Bond Angles: 120°

    • Alternate example:

    • Sulfur Dioxide (O-S=O)

    • Shape: Bent (although classified as trigonal planar due to lone pair effects).

Four Electron Domain Molecules

  • General Structure:

    • Example: Methane (H-C-H)

  • Geometry: Tetrahedral

    • Angle: 109.5°

  • Specific Cases:

    • Trigonal Pyramidal:

    • Geometry when 1 lone pair is present.

    • Bond Angles: Slightly less than 109.5° due to lone pair repulsion.

Five Electron Domain Molecules

  • Shapes:

    • Trigonal Bipyramidal: No lone pairs present.

    • Angles: 120° and 90°.

    • Specific Shapes:

    • Seesaw:

      • Geometry with 1 lone pair, where the lone pair is placed in an equatorial position.

      • Angles: 120° apart among atoms.

    • T-shaped:

      • Geometry with 2 lone pairs.

    • Linear:

      • Geometry with 3 lone pairs.

      • Angles: All lone pairs positioned at 120° apart.

Six Electron Domain Molecules

  • Shapes:

    • Octahedral:

    • No lone pairs present.

    • Bond Angles: 90° between all atoms.

    • Specific Shapes:

    • Square Pyramidal:

      • Geometry with 1 lone pair.

    • Square Planar:

      • Geometry with 2 lone pairs located above and below the square plane.

POLARITY OF MOLECULES

  • Polar and Non-polar Molecules:

    • Molecules with polar bonds can be non-polar if bond arrangement is symmetrical, as partial charges may cancel out.

    • Conversely, if polar bonds are arranged asymmetrically, the resultant dipole moments do not cancel, rendering the molecule polar.

TYPES OF INTERMOLECULAR FORCES

  • Hydrogen Bonds:

    • Example: Found between H₂O molecules in water, indicating strong intermolecular attraction.

  • Permanent Dipole-Dipole Forces:

    • Example: Present in HCl molecules in hydrogen chloride, indicating a permanent dipole due to the difference in electronegativity between H and Cl.

  • London Dispersion Forces:

    • Example: Found between I₂ molecules in iodine crystals, representing the weakest form of intermolecular attraction due to transient dipoles.