• Share
  • Email
  • Embed
  • Like
  • Save
  • Private Content
Chemical Bonding
 

Chemical Bonding

on

  • 3,941 views

 

Statistics

Views

Total Views
3,941
Views on SlideShare
3,937
Embed Views
4

Actions

Likes
9
Downloads
0
Comments
1

2 Embeds 4

http://geneseechemistry.weebly.com 3
http://www.weebly.com 1

Accessibility

Categories

Upload Details

Uploaded via as Microsoft PowerPoint

Usage Rights

© All Rights Reserved

Report content

Flagged as inappropriate Flag as inappropriate
Flag as inappropriate

Select your reason for flagging this presentation as inappropriate.

Cancel

11 of 1 previous next

  • Full Name Full Name Comment goes here.
    Are you sure you want to
    Your message goes here
    Processing…
  • the best one
    Are you sure you want to
    Your message goes here
    Processing…
Post Comment
Edit your comment

    Chemical Bonding Chemical Bonding Presentation Transcript

    • Chapter 6 CHEMICAL BONDING
      • Mutual electrical attraction between the nuclei and valence electrons of different atoms that finds that atom together
      • Why are most atoms chemically bonded together?
        • Atoms are less stable existing by themselves than when they are combined
        • By bonding with each other, atoms decrease in potential energy, which means they create more stable arrangements of matter
      CHEMICAL BOND
    • CHEMICAL BONDING VIDEO
      • Ionic bonding
        • Results from the electrical attraction between cations and anions
        • Atoms completely give up electrons to other atoms
      • Covalent Boning
        • Results from the sharing of electron pairs between two atoms
        • The shared electrons are “owned” equally by the two bonded atoms
      TYPES OF BONDS
    • COVALENT BONDS VIDEO CLIP
    • IONIC BOND VIDEO CLIP
      • Non-polar
        • The bonding electrons are shared equally by the bonded atoms
        • Results in a balanced distribution of electrical charge
        • Example: hydrogen-hydrogen bond
      • Polar
        • There is an uneven distribution of charge
        • The bonded atoms have an unequal attraction for the shared electrons.
      TYPES OF COVALENT BONDS
    • POLAR VS NON-POLAR COVALENT BONDS
      • Molecule
        • Neutral group of atoms that are held together by covalent bonds
      • Molecular Compound
        • Chemical compound whose simplest units are molecules
      • Chemical Formula
        • Indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts
      • Molecular Formula
        • Shows the types and numbers of atoms combined in a single molecule of a molecular compound
      • Diatomic Molecule
        • Molecule that contains only two atoms
      MOLECULAR COMPOUNDS
      • Bond length
        • The average distance between two bonded atoms
      • Bond Energy
        • The energy required to break a chemical bond and form neutral isolated atoms
        • Measure in kilojoules per mole (kJ/mol)
        • Example:
          • 436 kJ/mol of energy is needed to break hydrogen-hydrogen bonds in one mole of hydrogen molecules
      CHARACTERISTICS OF COVALENT BONDS
      • Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level
      • Exceptions to the rule:
        • Hydrogen
          • Forms bonds in which it is only surrounded by two electrons
        • Boron
          • Forms bonds in which it is surrounded by six electrons
      OCTET RULE
      • Electron configuration notation in which only the valence electrons of an atom of a particular element are shown
      • Valence electrons are indicated by dots placed around the element’s symbol
      ELECTRON DOT NOTATION
      • Formulas in which atomic symbols represent nuclei and inner-shell electrons
      • Dot pairs or dashes between two atomic symbols represent electron pairs in covalent bonds
      • Dots that are adjacent to only one atomic symbol represent unshared electrons
      • Unshared pair (lone pair) – a pair of electrons
      • that is not involved in bonding and that
      • belongs exclusively to one atom.
      LEWIS STRUCTURES
      • Indicates the kind, number, arrangement, and bonds but not the unshared pairs of atoms in a molecule
      STRUCTURAL FORMULAS
      • Single
        • Covalent bond in which one pair of electrons is shared between two atoms
      • Double
        • Covalent bond in which two pairs of electrons are shared between two atoms
      • Triple
        • Covalent bond in which three pairs of electrons are shared between two atoms
      • Multiple bonds
        • Either double or triple bonds
      BONDS
    • Multiple Covalent Bonds
      • Double Bond – a covalent bond produced by the sharing of two pairs of electrons between two atoms.
      Triple Bond – a covalent bond produced by the sharing of three pairs of electrons between two atoms. N N O O N N O O
    • IONIC BONDING AND IONIC COMPOUNDS
      • Ionic Compound – a compound composed of positive and negative ions that are combined so that the number of positive and negative charges are equal.
      • Formula Unit – the simplest collection of atoms from which an ionic compound’s formula can be established.
      • Ex. NaCl
    • FORMATION OF IONIC COMPOUNDS Na + + Cl - NaCl Ca 2+ + F - CaF 2 Ca 2+ + N 3- Ca 3 N 2
    • CHARACTERISTICS OF IONIC BONDING
      • Crystal lattice – an orderly arrangement of ions.
      • Lattice energy – the energy released when one mole of an ionic crystalline compound is formed for gaseous ions.
      • Ionic bonds are stronger than molecular bonds.
    • POLYATOMIC IONS
      • Polyatomic ions – a charged group of covalently bonded atoms.
    • METALLIC BONDING
      • Within a metal, the vacant orbitals in the atoms’ outer energy level overlap, allowing outer electrons of atoms to roam freely throughout the entire metal.
      • Delocalized electrons – electrons that do not belong to one atom, but can freely move about the metal’s network of empty atomic orbitals.
      • Metallic Bonding – the chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.
      • The shiny appearance of metals are due to the absorption of a high range of light frequencies, resulting in exciting/de-exciting electrons.
      • Malleability – the ability of a substance to be hammered or beaten into thin sheets.
      • Ductility – the ability of a substance to be drawn, pulled, or extruded through small opening to produce a wire.
      • Heat of Vaporization – the amount of heat required to vaporize a metal, which is the measure of the strength of the bonds that hold metal together.
    • MOLECULAR GEOMETRY
      • Molecular Geometry – the three-dimensional arrangement of molecule’s atoms in space.
      • Molecular Polarity – the uneven distribution of molecular charge.
    • VSEPR THEORY
      • Valence Shell Electron Pair Repulsion (VSEPR Theory) – the repulsion between the sets of valence-level electrons surrounding an atom cause these sets to be oriented as far apart as possible.
      • Example: CH 4 (bonding pairs only, no lone pairs)
      Key: Must consider both bonding and lone pairs in minimizing electron repulsion. Lewis Structure VSEPR Structure
    • • Example: NH 3 (both bonding and lone pairs). Lewis Structure VSEPR Structure Molecular Shape
    • VSEPR APPLICATIONS The previous examples illustrate the strategy for applying VSEPR to predict molecular structure: 1. Construct the Lewis Dot Structure 2. Arrange the bonding and lone electron pairs in space such that repulsions are minimized .
    • Case: Linear Structure ( AX 2 ): angle between bonds is 180° Example: BeF 2 180°
    • Case: Trigonal Planar Structure ( AX 3 ): The angle between bonds is 120° Example: BF 3 120°
    • Case: Pyramidal ( AX 3 E ): Bond angles are <120° structure is nonplanar due to repulsion of lone-pair. Example: NH 3 107° VSEPR Structure Molecular shape Lewis
    • Case: Tetrahedral ( AX 4 ): the angle between bonds is ~109.5° Example: CH 4 109.5°
    • Note: for ‘Tetrahedral’, the actual angle may vary slightly from 109.5°, due to size differences between bonding and lone pair electron densities bonding pair: more elongated, less repulsive lone pair: puffier, more repulsive
    • Example of distorted tetrahedron: water ( AX 2 E 2 ): the angle is reduced to 104.5° by repulsion of the lone pairs “ bent” VSEPR structure Molecular shape
    • Case: Trigonal Bipyramidal ( AX 5 ): non-equivalent bond positions: three in-plane (equatorial, 120°), and two at 90° to plane (axial) Example, PCl 5 90° 120°
    • Octahedral ( AX 6 ): all angles are 90° Example SF 6 90° Lewis VSEPR
    • INTERMOLECULAR FORCES
      • Intermolecular Forces – the force of attraction between molecules.
      • Dipole – equal and opposite charges that are separated by a short distance.
      • Dipole-Dipole Forces – forces of attraction between polar molecules.
      H - Cl
      • Hydrogen Bonding – the intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electromagnetic atom in the nearby molecule.
      • London Dispersion Forces – the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles.
        • Act between all atoms and molecules
        • The only intermolecular forces acting among noble-gas atoms, nonpolar molecules, and slightly polar molecules
      • Only intermolecular among noble gasses and non-polar molecules.