Bonding Concepts

Most of the universe is made up of compounds
- Everything is trying to get to the lowest energy state possible
- Compounds are more stable than atoms on their own (atoms achieve a low energy state)
The most stable configuration is when energy is released (groud state, lowest energy configuration)
- Forming a bond releases energy (exothermic)
- Bonds store potential energy
- Some atoms form multiple bonds (become more stable, release more energy for each bond formed)
  
  $$ A + B → AB + ΔH $$
In compounds, the nucleus of one atom is attracted to the electron of the other atom (negative force - attraction, negative energy - exothermic)
- A is holding onto B’s valence electrons, B is holding onto A’s valence electrons
- Electron density overlaps
As radius (distance) between atoms increases, the energy of interaction approaches 0 (very little interaction between the nucleus of one atom and electron of the other atom, use Coulomb’s law)

As radius (distance) between atoms decreases, the energy of the interaction increases (becomes more exothermic)
When atoms come too close together, they start to repel (can’t release any more energy, nuclei are close together and repel more than they are attracted to valence electrons)
- Nucleus (very small) only repel when they are on top of each other (not being shielded by electrons
Vibrates between repelling and attracting and eventually settles at lowest potential energy state
Reactants (two separate atoms) and products (compounds and energy releases) have equal amounts of energy (law of conservation of energy)
The graph shows that as the nucleus and electron cloud of the other atom overlap more, the potential energy decreases (exothermic), then the frame of reference changes to the repulsion of nuclei (becomes endothermic)
The lowest point on the graph is the most stable (lowest energy conformation possible, balance between attraction and repulsion), number on the y-axis is the enthalpy of the bond
- The opposite (absolute value) is the amount of energy it takes to break the bond (Hess’s Law), the amount of energy to break or form are opposite values
Bond length (distance between the nuclei) depends on the size of the atoms being bonded

Link to simulation for repulsion vs. attraction as two atoms move toward each other

Lewis Dot Structures

<aside> ⚛️ Typical bonding arrangement of a given atom (does not always have to be used)

</aside>

Determine by imagining a box the element’s symbol
Place each valence electron on one side of the box
Do not start creating pairs until each side has at least one
Place a maximum of 2 electrons per side
The number of unpaired/lone electrons shows bonding sites (where a bond can form)

NBA for Period 2	Li	Be	B	C	N	O	F	Ne
	$\Huge{{\cdot \above{0pt} {Li}} \above{0pt} {}}$	$\Huge{{\cdot \above{0pt} {~Be~\cdot}} \above{0pt} {}}$	$\Huge{{\cdot \above{0pt} {\cdot~B~\cdot}} \above{0pt} {}}$	$\Huge{{\cdot \above{0pt} {\cdot~C~\cdot}} \above{0pt} {\cdot}}$	$\Huge{{\cdot \above{0pt} {\cdot~N\colon}}\above{0pt} {\cdot}}$	$\Huge{{\cdot \above{0pt} {\colon O\colon}}\above{0pt} {\cdot}}$	$\Huge{{\cdot\cdot \above{0pt} {\colon F\colon}}\above{0pt} {\cdot}}$	$\Huge{{\cdot\cdot \above{0pt} {\colon {Ne} \colon}}\above{0pt} {\cdot\cdot}}$
Number of Valence Electrons	1	2	3	4	5	6	7	8
Number of Bonding Sites	1	2	3	4	3	2	1	0

Noble gases have the maximum number of valence electrons, no lone electrons (no bonding sites)
Bonding atoms have more unpaired electrons than the atom does at ground state (paired electrons are excited to an empty orbital, against Aufbau principle)
- Slightly endothermic, but the bonds forming (exothermic process) accounts for this
- Core electrons cannot be excited to become unpaired (too endothermic)