Lewis Structures Molecular Shape VSEPR Model (Valence Shell Electron Pair Repulsion Theory)
PART 1: Ionic Compounds Complete the table of Part 1 by writing: The Lewis dot structures for each metallic and non-metallic atom The Lewis dot structures for the common ions of the metal and nonmetal. The formula for the compound formed between them. The Lewis structure of the compound formed between them.
Here is an example of how the table should be completed using barium and chlorine: Element Lewis Dot Structure of the Element Lewis Dot Structure of its Ion Formula of the Compound Lewis Dot Structure of the Ionic Compound Ba Ba : Ba 2+ Cl.... _ BaCl 2.. _ : Cl. : Cl : Ba 2+ : Cl :...... 2
Ba: Located in group 2 of the periodic table. Therefore, the Lewis structure of the element has 2 dots. It will lose its two valence electrons to form the +2 ion. The Lewis structure of the ion of barium has no dots and a +2 charge. Cl: Located in group 17 of the periodic table. Therefore, the Lewis structure of the element has 7 dots. It will gain one electron to give it a complete octet in the valence shell to form the -1 ion. The Lewis structure of the ion of chlorine has 8 dots and a -1 charge.
Since barium forms a +2 ion and chlorine forms a -1 ion, they will combine in the ratio of 1 barium ion to 2 ions of chlorine: BaCl 2 The Lewis structure of the ionic compound is written by writing the Lewis structure of the cation beside the Lewis structure of the anion. If the formula of the compound requires more than one of any ion, that ion is enclosed in parentheses and the subscript is shown.
PART 2: Molecular Compounds Complete the table of Part 2 by: Summing up the valence electrons in molecular substance. Drawing the Lewis structure. Counting the total electron pairs on the central atom. Counting the total non-bonding electron pairs on the central atom. Writing the molecular geometry.
Here is how the table should be completed showing an example of ClO 3 - Molecule Total Number of Valence Electrons Lewis Dot Structure Total # of Electron Pairs on the Central Atom Total # of Nonbonding Electron Pairs on the Central Atom Molecular Geometry...... - ClO 3-26 : O Cl O :.... : O : 4 1 Trigonal pyramid..
A molecular model of each substance in the table of part 2 is shown in the following slides.
CH 4 H 2 O
H 3 O + HF
NH 3 NH 4 +
OH- F2
O 2 N 2
SO 4 2- CO 2
HCN NO 3 -
Molecular Shape Overall molecular shape depends upon bond angles. Bond angles are arranged so that electrons are spaced as far apart as possible.
VSEPR Model (Valence Shell Electron Pair Repulsion Theory) Electrons have like charges and repel one another. Therefore, they space themselves as far apart as possible. 1. Double and triple bonds are treated like single bonds. 2. VSEPR can be applied to any resonance structure.
Electron-Pair Geometry Orbital Geometry Geometrical arrangement of all electron pairs about the central atom. 2 Types of electron pairs: 1. Bonding or shared 2. Non-bonding or unshared or lone pairs.
Molecular Geometry Geometrical arrangement of atoms in space. Used to describe the shape of the molecule.
Steps for Determining Molecular Shape 1. Draw the Lewis Dot Structure. 2. Determine the orbital geometry by counting up the total number of electron pairs on the central atom, counting double and triple bonds as one pair. 3. Determine the molecular geometry by considering the number of lone pairs on the central atom.
Linear AB 2 Trigonal Planar AB 3 Tetrahedral AB 4
Trigonal Bipyramidal AB 5 Octahedral AB 6
Bond Angles Non-bonding electron pairs are not attracted to the nucleus of another atom. Therefore, they exert a greater repulsive force on adjacent electron pairs and compress the bond angle.
Orbital Geometry Molecular Geometry Linear Linear 180 o Trigonal Planar Tetrahedral Trigonal Bipyramidal Octahedral Trigonal Planar Bent Tetrahedral Trigonal Pyramidal Bent Trigonal Bipyramidal Seesaw T-Shaped Linear Octahedral Square Pyramid Square Planar Bond Angle(s) 120 o <120 o 109.5 o 107 o 105 o 90 o, 120 o, 180 o 90 o, <120 o, 180 o 90 o, 180 o 180 o 90 o 90 o 90 o
Polar Molecules Molecules in which the centers of positive and negative charge do not coincide. (a dipole exists) Dipole: Established when charges of equal magnitude, but opposite sign are separated by a distance.
Molecule Polarity Depends upon: 1. Bond Polarity within the molecule 2. Molecular Geometry
Symmetrical Geometries: Linear Trigonal Planar Tetrahedral Trigonal Bipyramid Octahedral Square Planar
Valence Bond Theory When a valence orbital of 1 atom overlaps with a valence orbital of another and electrons are concentrated (shared) between them to form a covalent bond.
Formation of H2
Hybrid Orbitals Formed when two or more atomic orbitals combine in order to form bonding orbitals. The number of hybrid orbitals equals the number of atomic orbitals combined.
sp Hybridization
BeCl 2 Be: 1s 2 2s 2.... :Cl Be Cl :.... 1s 2s 2p 1s 2sp hybrid 2p p orbital sp hybrid p orbital orbital
sp 2 Hybridization
BF 3 B: 1s 2 2s 2 2p 1 1s 2s 2p 1s 2sp 2 hybrid 2p.... : F :.. : F B F:....
sp 3 Hybridization
CH 4 C: 1s 2 2s 2 2p 2 1s 2s 2p 1s 2sp 3 hybrid H H C H H
NH 3 N: 1s 2 2s 2 2p 3 1s 2s 2p 1s 2sp 3 hybrid H : N H Lone pairs H
H 2 O O: 1s 2 2s 2 2p 4 1s 2s 2p 1s 2sp 3 hybrid H : O : Lone pairs H
PF 5 P: 1s 2 2s 2 2p 6 3s 2 3p 3 3s 3p 3d 3sp 3 d hybrid 3d F F F P F F
sp 3 d Hybridization
SF 6 S: 1s 2 2s 2 2p 6 3s 2 3p 4 3s 3p 3d 3sp 3 d 2 hybrid 3d F F F S F F F
sp 3 d 2 Hybridization
Sigma Bonds (σ) First bond formed between atoms. Head-to-head orbital overlap. Internuclear axis passes through the overlap region.
Pi Bonds (π) Additional bonds formed between atoms. Side-to-side orbital overlap. Overlap region is above and below the internuclear axis. Only forms when there are unhybridized p orbitals on the bonded atoms.
Ethene (ethylene) C 2 H 4 H H C: 1s 2 2s 2 2p 2 C C H H 1s 2sp 2 hybrids 2p σ bond orbitals π bond orbital
Ethyne (acetylene) C 2 H 2 H -C C -H C: 1s 2 2s 2 2p 2 1s 2sp hybrids 2p σ bond orbitals π bond orbitals