Atomic and Molecular Structure
Study Guide
Atomic and Molecular Structure
Recognize Lewis structures for common elements and their characteristics. For example,
For argon:
Argon has 8 valence electrons.
Argon has a pair of electrons on each side of the atom.
Argon is a noble gas.
The total number of electrons in a Lewis structure is equal to the sum of the numbers of valence electrons on each individual atom. Electrons are placed initially as lone pairs until each outer atom has eight electrons in bonding pairs and lone pairs. For example, oxygen has 6 valence electrons. Nitrogen has a lone valence electron pair in the Lewis structure for ammonia.Radon (a noble gas) has no valence electrons. When calculating total valence electrons, extra electron pairs are added to which atom. When bonded to a central atom, hydrogen will have no lone electron pairs.
If the molecule is ionic, the Lewis structure will show the electrical charge. For example, hydroxide (OH-)is a polyatomic ion; hydroxide has a negative charge and the Lewis structure will have a negativecharge. For ammonium:
Ammonium is a polyatomic ion with a charge of 1+. It has a positive charge.
In the Lewis structure for ammonium, nitrogen is the central atom.
Ammonium has 4 hydrogens with single bonds to the nitrogen
Hydroxide (OH-) is a polyatomic ion with a negative charge.
Draw Lewis structures following the rules in the lesson on Lewis structures. Keep in mind the following exceptions:
- Hydrogen atoms will never have more than two electrons in its valence shell. Its
Lewis structure will only be bonded once and will not have lone valence electron pairs.
- Beryllium can have four electrons in its valence shell.
- Boron and aluminum can have six electrons in their valence shells.
Determine the number of valence electrons for the elements.
Keep in mind the special characteristics of the Lewis structure for boron:
Boron is a Group III element.
Boron has 3 valence electrons.
Boron has 3 single dots on 3 sides of the atom.
Boron is usually the central atom.
Study packing rules (or the atomic packing factor), which apply to ionic compounds:
Packing rules are rules by which molecules are shaped by the
repulsion/interaction among the atoms in an ionic compound. Atoms form centers, faces and corners of crystalline structures, with smaller atoms taking up the space between the larger ones.
VSEPR applies to covalent (molecular) compounds. It serves as a model for understanding how atoms and electrons are assembled into molecules.
Orbitals – s, p, d, and f.
Bonds:
Single bonds consist of one single bond. They are generally weaker than double bonds.
A double bond consists of one sigma and one pi bond. They are the same length as single bonds.
A triple bond consists of 1 sigma and 2 pi bonds
Sigma bonds are characteristic of s orbitals.
Sigma bonds occur as a result of overlapping s-orbitals. Orbital overlap forms sigma and pi bonds.
Pi bonds are most frequently found in p-orbitals.
Pi bonds are also most frequently found in double bonds. Sigma bonds occur as a result of overlapping s-orbitals.
Bond angles – the angle between three bonded atoms
Torsion angle – the angle between the plane formed by the first three atoms and the plane formed by the last three atoms
Resonance structures- structures that occur when it is possible to write two or more valid electron dot formations that have the same number of electron pairs for a molecule or ion. They are simply a way to envision the bonding in certain molecules.
If you refer to the two dot structures for ozone (O3), for example,
you will see that each structure can be converted to the other by shifting electron pairs without changing the positions of the oxygen atoms.
If the structures you draw follow the octet rule (8 valence electrons around every atom), then there will be no resonance.
If the Lewis structure doesn't follow the octet rule for every atom in the molecule, then the resonance structures will be equal to the number of bonds.
Notice thatCO3 has 3 bonds and an unpaired electron. There will be three resonance structures. The same guideline can be used for SO3. It also has 3 resonance structures. NH3 has no resonance because it satisfies the octet rule - the nitrogen has four pair of electrons and each hydrogen has one pair.
Nitrogen has five valence electrons. N3- will have fifteen valence electrons plus one extra. The extra electron will migrate around the ion - so the ion will have three resonant structures. At any given time there will be two double bonds and one single bond.
Here is a good method:
1.Count the valence electrons for each atom in the compound.
2. Draw a skeletal structure.
3. Check the non-central atoms to see if they obey the octet rule.
4. Check the central atom for the octet.
Not all molecules have resonance. If the octet rule is obeyed, there will be no resonance. Study the examples in the lessons and do the practice questions. For example, acetylene(C2H2) will have no resonance because the octet rule is obeyed. Methane (CH4) is another one with no resonance.
Resonance hybridization refers to the intermediate properties of all resonance structures.
Review the different types of molecular geometry. For example, the polyatomic ion nitrate (NO3-) has a trigonal planar molecular geometry.
Repulsion – a force or interaction among atoms in a molecule in which the atoms are repulsed by one another.The following repulsion force strength is accurate: Bonding Pair-Bonding Pair < Lone Pair-Bonding Pair < Lone Pair-Lone Pair.
Study bond angles, the angle between three bonded angles. For example, a bent angle will have an angle smaller than 120o. A T-shaped molecule will have an angle of 90o. The correct angle between the atoms in a seesaw molecule is <90o, <120o. A linear molecule will have an angle of 180o. An octahedral will have 90o. Water has a bent molecular geometry.
Molecular geometry and symmetry can help us determine a substance’s color, magnetic properties, and polarity.