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Results from the electrical attraction between cations and anions
Atoms completely give up electrons to other atoms
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
The bonding electrons are shared equally by the bonded atoms
Results in a balanced distribution of electrical charge
Example: hydrogen-hydrogen bond
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
Neutral group of atoms that are held together by covalent bonds
Chemical compound whose simplest units are molecules
Indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts
Shows the types and numbers of atoms combined in a single molecule of a molecular compound
Molecule that contains only two atoms
The average distance between two bonded atoms
The energy required to break a chemical bond and form neutral isolated atoms
Measure in kilojoules per mole (kJ/mol)
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:
Forms bonds in which it is only surrounded by two electrons
Forms bonds in which it is surrounded by six electrons
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.
Indicates the kind, number, arrangement, and bonds but not the unshared pairs of atoms in a molecule
Covalent bond in which one pair of electrons is shared between two atoms
Covalent bond in which two pairs of electrons are shared between two atoms
Covalent bond in which three pairs of electrons are shared between two atoms
Either double or triple 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.
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 – a charged group of covalently bonded atoms.
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 – the three-dimensional arrangement of molecule’s atoms in space.
Molecular Polarity – the uneven distribution of molecular charge.
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 – 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.