Chemistry
Electronic structure
Shells / energy levels (A group of atomic orbitals with the same quantum level)
Maximum number of electrons in each energy level (2n^2)
An electron is a negatively charged subatomic particle with nearly zero mass
These are split into subshells (a group of the same type of atomic orbitals within a shell)
These are split into orbitals ( a region around the nucleus which can hold up to 2 electrons with opposite spins)
2 shells
3 shells
1 shell
4 shells
2 electrons
8 electrons
18 electrons
32 electrons
The shell number is called the principal quantum number
Orbitals can hold one or two electrons with opposite spins to minimise repulsion
There are different orbital types: s- , p- , d- , f-
Each type of orbital has a different shape
D and F orbitals
S-orbitals
P-orbitals
Spherical shape
There are 3 separate p-orbitals at right angles (3 orbitals each fit 2 electrons meaning each 2p sub shell holds 6 electrons)
Each shell from n=4 contains 7 f-orbitals which fits 14 electrons
Each energy level contains one S-orbital
The greater the number of shells, the greater the radius of the sphere (one orbital fits 2 electrons)
Each energy level from n=2 contains 3 p-orbitals and the higher n is, the wider the radius is
Have the shape of a dumbell
Energy level 3 = 3S (2 electrons) + 3P (6 electrons) + 3d (10 electrons) = fits 18 electrons
Each shell from n=3 contains 5 d-orbitals which fits 10 electrons
Instead of shellls, we now use energy levels
Each energy level is divided into subshells. The first energy level has the first subshells which holds 2 electrons
1S^2, 2S^2, 2p^6, 3S^2, 3p^2, 4S^2, 3d^10, 4p^6
The 3rd and the 4th energy level overlap since the 3d orbital has a higher energy than 4s
The 3rd energy level fills before 4p energy level which is shown by the Aufbau principle
electron pairs have opposite spins
orbitals are simplified to boxes and each box can have 2 electrons with opposite spins (up or down)
this is the 'electron in a box' model
orbitals within a sub-shell have the same energy so an electron will occupy each orbital before pairing would start
the periodic table can be separated into blocks
p-block
d-block
s-block
highest energy electrons in s-subshell
highest energy electrons in p-shell
highest energy electrons in d-subshell
Example
K = 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^1
Na + = 1s^2, 2s^2, 2p^6
since it is an ion, it loses one electron from 3s^1
the 4s subshell fills before the 3d subshell. For any d-block ion, the s-subshell empties first
ionic bonding
the bonds between atoms always involve their outer electrons. Noble gases have full outer main levels of electrons and are very unreactive
when atoms bond together, they share or transfer electrons to achieve a more stable electron arrangement, OFTEN a full outer main level of electrons, likethe noble gases
an ion is a positively or negatively charged atom or group of atoms, where the number of electrons is different from the number of protons. When you lose or gain electrons, you form ions
electrostatic attraction- where opposite charges are attracted to each other (holds ions together)
each ion attracts to oppositely charged ions in all directions
the structure of an ionic compound (due to electrostatic attraction) will be an ionic lattice
properties of ionic compounds
the electrostatic attraction extends throughout the lattice and a giant lattice structure is formed. These constain billions of ions and the actual number is determined by the size of the crystal
tend to dissolve in polar substances such as water
conduct electricity in the liquid state or aqueous solution
high melting and boiling points (solid at room temperature)
electrostatic attraction results in strong bonding which extends throughout the compound, meaning ionic compounds have very high melting and boiling points. Energy must be supplied to break up the lattice by overcoming the force of electrostatic attraction. Melting points are higher for lattices with greater ionic charges
dissolve in polar solvents such as water. The water molecules break down the lattice and are then attracted to and surround each ion in solution. In compounds with larger charges, the ionic attraction may be too strong so they will not be very soluble
As a solid the ions are not free to move in their solid state. As a liquid, the ions are more free to move in their liquid state and the compound conducts electricity.
covalent bonding
a strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atom
a covalent bond is the overlap of atomic orbitals containing 1 electron to give a shared pair of electrons
this attraction is localised, only acting between the shared pair of electrons and the nuclei of the two atoms
example
boron forms 3 covalent bonds since it only has 3 outer electrons, meaning it will not form a full outer shell when covalently bonded (there will be 6 electrons)
this disproves octec rule of 8 outer electrons
sulfur - can form 2,4 or 6 covalent bonds (can have up to 12 electrons in their outer shell)
look at how many electrons are in the outer shell and this shows how many covalent bonds can form
a double bond is an electrostatic attraction between two shared electron pairs and the nuclei of the bonded atoms
a triple bond is an electrostatic attraction between three shared pairs of electrons and the nuclei of the bonded atoms
dative covalent bond
the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms where both electrons have come from one atom
they are also known as a co-ordinate bond (represented by ->). The shared pair of electrons was originally a lone pair of electrons of one of the bonded atoms
e.g. an ammonia molecule donates its lone pair of electrons to the nitrogen to a H+ ion
all 4 bonds are equivalent and you cannot distinguish between them
Averge Bond Enthalpy
a measure of covalent bond strength. The larger the value, the stronger the bond (given in exams)
most covalent compounds have small molecules. They are gases, liquids (or solids) with low melting points and low boiling points