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Chapter 5 - Electrons and bonding - Coggle Diagram
Chapter 5 - Electrons and bonding
Shells
Shells are regarded as energy levels
Energy level increases as shell number increases
Principal quantum number (n) - Shell occupied by the electrons
Maximum number of electrons in a shell is 2n^2
Maximum number of electrons
Shell 1 = 2
Shell 2 = 8
Shell 3 = 18
Shell 4 = 32
Atomic orbitals
Shells are made up of atomic orbitals, clouds of negative charge
Region in space where there is a high probability of finding an electron
S orbitals
Electron cloud is within the shape of a sphere
Each shell from n = 1 contains one s-orbital
Greater shell number = greater radius of s-orbital
Electron cloud - negative-charge cloud with the shape of the orbital
p orbitals
Electron cloud is within the shape of a dumb-bell
Three seperate p-orbitals at right angles to one another
Each shell from n=2 contains three p-orbitals
Greater shell number = further p-orbital is from nucleus
Can hold up to two electrons
D orbitals
Each shell from n = 3 contains five d-orbitals
F orbitals
Each shell from n = 4 contains seven f-orbtials
Filling orbitals
Increasing energy
Within each shell, the new type of sub-shell added has a higher energy
Exception
3d sub shell is at a higher energy level than the 4s sub shell
4s sub-shell fills before the 3d subshell
Order: 1s 2s 2p 3s 3p 4s 3d 4p 4d 4f
Aufbau principle - Lowest energy level is filled first
Opposite spins
An electron is shown as an arrow indicating its spin, either up or down
Two electrons in an orbital must have opposite spins, this is to counteract repulsion between negative charged of electrons
Paulis exclusion principle - Electrons cannot spin the same way
Orbitals with same energy
Within a sub shell, orbitals have the same energy
One electrons occupies each orbital before the pairing starts
This prevents any repulsion between paired electrons until there is no further orbital available at the same energy level
Hund's rule - electrons prefer orbitals on their own, they only pair up when no empty orbitals are available
Electron configuration
Shows how sub-shells are occupied by electrons
Electron configurations can be expressed in terms of the previous noble gas in the periodic table plus outer electron sub shells
Useful for emphasising similarities in the electron configuration of the outer shell
Shorthand - Electron configurations can be expressed more simply in terms of the previous noble gas in the periodic table plus outer electrons subshells
Ions
s and p block
Highest energy sub-shells lose/ gain electrons
d block
4s subshell is at a lowr energy level than the 3d subshell, so is filled first
Once filled, the 4s shell has a higher enegry level than the 3d
4s sub-shell fills before the 3d sub-shell
4s sub shell empties before the 3d sub-shell
Ionic bonding
Electrostatic attraction between positive and negative ions
Dot-and-cross diagrams
Electrons in the original atoms are shown as either dot or crosses
Structure of ionic compounds
Each ions attracts oppositely charged ions in all directions
The result is a giant ionic lattice
Properties
Melting and boiling points
Melting points are higher for lattices containing ions with greater ionic charges as there is a stronger attraction between the ions
High melting/ boiling points are large quantity of energy is needed to overcome the strong electrostatic forces of attraction between oppositely charged ions
Solubility
Many ionic compounds dissolve in polar solvents, which break down the lattice
In a compound made with large charged, ionic attraction may be too strong for water to break down the lattice, so it may not b very soluble
Solubility required the lattic to be broken down and the water molecules must attract and surround the ions
Conductivity
Solid state - doesn't conduct as ions are in fixed positions in the giant ionic lattic, no mobile charge carriers
Liquid/ dissolved in water - solid ionic lattice breaks down, ions are now free to move as mobile charge carriers
Covalent bonding
Electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
Occurs between non metals and polyatmic ions
Orbital overlap
Overlap of atomic orbitals, each containg one electron
Shared pair of electrons is attratced to the nuclei of the bonding atoms
Localised
Covalent bonds only acts between the shared pair of electrons and the nuceli of the two bonded atoms, this is a molecule
Single covalent bond
Dot and cross
Allow origin of each electron to be shown
Triangle is used to show extra electrons
Displayed formula
Relative positions of atoms due to bonds between them
Lone pairs, electrons that aren't shared, they can be added to displayed formula
n of covalent bonds
C = 4, N = 3, O = 2, H = 1
For elements in period 2, their shell can hold just eight electrons
But for elements in period 3, they can have up to 18 electrons in their outer shell, so more electrons are available for bonding
Expansion of the octet
Multiple covalent bonds
Two atoms share more than one pair of electrons
Double covalent bonds
Electrostatic attraction is between two shared pairs of electrons and the nuclei of the bonded atoms
Triple bonds
Electrostatic attraction is between three shared pairs of electrons and the nuclei of the bonded atoms
Dative bonds
Shared pair of electrons has been supplied by one of the bonding atoms only
Shared electron pair was originally a lone pair of electrons on one of the bonded atoms
Shown with an arrowhead
Average bond enthalpy
Measurement of covalent strength
Larger the average bond enthalpy, the stronger the covalent bond