Bonding

You don’t need to know the text with a strikethrough for the test (it’s good to know though).

Note: Some equations sometimes don’t show up on the Google Docs app, try to use the website whenever possible.

YouTube and Helpful links will be next to important concepts/topics, don’t hesitate to watch some videos to clear up misconceptions or confusion.

Scroll down for practice problems - answers here

General Bonding (Chapter 8)

This unit is largely conceptual, so make sure you understand the why behind the concepts.

It’s important to know what exactly we mean when we say **chemical bonds **in this chapter. In simple terms; a chemical bond is a force of attraction between 2 atoms or a group of atoms

Chemical Bonds
3 Extreme Types
Ionic Bonds - Non-metal and metal - Stolen Electron (+) ion is attracted to (-) ion, (-) ion takes the atom. Lattice properties - conduct in water - extremely high melting points	Covalent Bonds - 2 Non-metals - Share Electron Nucleus of an atom is attracted to the valence electron(s) of another.	Metallic Bonds - Metals (+) metal cation is attracted to delocalized (free) valence electrons.

2 reasons why Chemical Events Occur
Movement toward minimum energy (enthalpy)	Movement toward maximum disorder (entropy)
The Entropy of the Universe is always increasing
This explains why chemical bonds form: the system is at lower energy after they form.

Lattice Energy - exothermic The energy released when forming 1 mol of an ionic solid from its gaseous ions.

Bond Energy - endothermic Energy required to break a mole of bonds in the gaseous phase.

Coulomb's Law Lattice energy = Where Q is the charge of the ion, d is the distance between them, and k is Coulomb’s constant.

Molecular Structure Important: understanding this is essential to understanding the properties of matter 2 main factors: Shape of Molecules - 3D geometry Distribution of Charge - electronegativity

Electronegativity - Video (Includes Periodic Trends) A measure of an atom’s ability to attract the valence electrons of another atom Note: we generally use Pauling’s scale for electronegativity, though there are others.
Increases as you go across rows This is explained by the increasing Zeff (Shielding electrons are constant, but protons increase, so the valence electrons are pulled in tighter, causing a smaller atomic radius). More nuclear charge and smaller atomic radius means the atom is able to pull on another atom’s valence electron more effectively.	Decreases as you go down columns This is explained by the constant Zeff (radius is increasing so less pull outside the atom). The larger radius means the atom pulls on another atom’s valence electron less effectively.
Polarity If a molecule’s shape isn’t symmetrical, chances are it’s polar. Add the vectors of dipoles to see if there’s a dipole moment.

Lewis Structures - Video
How to draw them, 3 steps
Step 1. Draw the structure of the compound itself, draw single bonds to start. Count the number of valence electrons for all the elements in the compound (eg, CH4 would be 4 (C) +4*1 (H): 8 electrons)	Step 2. Subtract the number of bonded electrons from the valence electrons--that is, subtract 2 from the valence electrons for every 1 bond(s).	Step 3. If you need more electrons than valence electrons available to fill the shells of the atoms, form a double bond for every 2 missing electrons. If you have extra electron pairs, place them on the central atom, this molecule isn’t following the octet rule, and have an expanded octet. (expanded octets use empty D orbitals) If you have an odd number of electrons, one atom gets 7 electrons. (Groups II and III usually are odd)
The stronger the bond, the closer the bond. Polar bonds are weaker than nonpolar bonds.
Resonance Structures - Video (Don’t need to know the Resonance Hybrid) If you have double/triple bonds on not all atoms; resonance structures are just the lewis dot structure with the bonds in different places, since they aren’t actually held between 2 atoms. Move electrons, not atoms Expected when you’re able to draw more than one structure following octet rules

The VSEPR Model - Video
Electron groups are a non-bonded pair of electrons or a bonded atom (single, double, and triple bonds all count as 1 electron group)
Electron groups try to stay as far from each other as possible, they repel each other
Note: lone pairs usually take more room than bonded pairs, but we don't need to know more

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Properties of Covalent Bonds

Pure Covalent Bonds - Video Equal sharing of electrons

Polar Covalent Bonds - Video Unequal sharing of electrons Polar if a molecule has polar bonds and geometry where the polarity doesn’t cancel each other (in terms of adding vectors)

Formal Charge - Video Used to choose most appropriate structure when there’s multiple non-equal lewis structures (Formal charges closest to 0 with negative charges on most electronegative atoms are best description of bonding of atom/ion)

Formal charge is = The sum of formal charges in molecule/ion must equal overall charge

Covalent Bonding (Chapter 9)

I recommend looking over your notes packet for this chapter, there’s a lot of written examples and more context.

The Atomic Orbital Approach (Valence Bond) - Video
Why? Lewis Dot Structures are good, but not always.
Why are they good?
Predict Geometries	Predict Polarities of Molecules
Why are they bad?
They give no information on the energies of electrons	They give no information about orbitals used in bonding
The valence bond approach is helpful for this
A covalent bond is formed from a pair of electrons with opposite spins in overlapped atomic orbitals. - Sigma Bond (𝞼) To explain bonding for atoms, it’s assumed that electrons are promoted and hybrid orbitals are formed. Orbitals that don’t need to be used don’t hybridize, thus they’re called unhybridized orbitals

Hybrid Orbitals - Video Have new properties, different from orbitals used to form them
Shape	Energy
The geometries are the same as predicted from VSEPR

Hybridization Chart (not mine, copied straight from the notes packet)
# electron groups	Hybridization	Geometry	Angle
2 AX2		Linear	180°
3 AX3 AX2E		Equilateral Triangle	120°
4 AX4 AX3E AX2E2		Tetrahedral	109.5°
5 AX5 AX4E AX3E2		Trigonal Bipyramidal	120°, 90°, 180°
6 AX6		Octahedron	90°, 180°

2 Types of Covalent Bonds
Sigma Bonds (𝞼) From from hybridized orbitals - overlap head to head	Pi Bonds (ℼ) Form from unhybridized orbitals - overlap side to side
# of Electron groups is the number of hybrid orbitals needed
Unhybridized p-orbitals will either be empty or be ½ filled (can form a pi bond by overlapping another p-orbital side to side)

Practice Problems

Click here for answers

1. Determine which of the following bonds is more polar using only Periodic Trends:
   1. Mg-O or Be-O
   2. Mg-S or K-Cl
   3. Ca-S or Sr-S
   4. Sr-At or Ba-I
2. Draw the Lewis Structures (and any Resonance structures) for the following molecules: 5. 6. 7.
3. List the Electron Geometry, Molecular Geometry, and Angles for each of the following: 8. 9. 10. 11. 12.

Answers to Practice Problems (w/ Explanations)

1. Determine which of the following bonds is more polar using only Periodic Trends:

   1. Mg-O is more polar; we know that electronegativity increases as we go across a period (row) and decreases down a group (column). Since O is in both bonds, we can ignore it’s electronegativity and look at the others. Be is higher than Mg in the same group, so it must have a higher electronegativity, this means the electronegativities of O and Be are more similar than that of **Mg **and O, so the electrons spend less time with a specific atom in Be-O than they do in **Mg-O. **When asked for which is more polar, essentially you’re looking for the compound in which the electron spends more time with a certain atom.
   2. K-Cl, they’re on different periods + further apart in terms of columns.
   3. Sr-S, they’re further apart.
   4. Hard to tell; not really discernable with periodic trends without knowing actual #’s

2. Draw the Lewis Structures (and any Resonance structures) for the following molecules (understand this by clicking here): 5.

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7. 

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3. List the Electron Geometry, Molecular Geometry, and Angles for each of the following (understand this by VSEPR Model): 8. Linear, Linear, 180° 9. Tetrahedral, Bent, 109.5° 10. Octahedral, Octahedral, 90° & 180° 11. Octahedral, Square Pyramidal, 90° & 180° 12. Trigonal Planar, Trigonal Planar, 120°