Module 2.2 Electrons, Bonding and Structure Flashcards
The principle quantum number, n, indicates
The shell that the electrons occupy
The larger the value of n,
The further the shell is from the nucleus and the higher the energy level
Number of electrons in the first shell
2
Number of electrons in the second shell
8
Number of electrons in the third shell
18
Number of electrons in the fourth shell
32
An orbital can
Hold a max. of 2 electrons with opposite spins
An orbital is a region of space where electrons are most likely to be found
4 different types of orbital
s
p
d
f
S orbitals
Spherical shape
Can hold 2 electrons
P orbitals
3D dumbbell shape
Can hold 6 electrons
Electron energy levels
1s
2s
2p
3s
3p
4s
3d
4p
4d
4f
Filling shells and sub-shells rules
The lowest energy level is filled first
Each energy level must be full before the next higher energy level starts to fill
Each orbital is filled singly before pairing starts
The 4s orbital is at a slightly lower energy level than the 3d orbital. Therefore, 4s fills before 3d
Electron configuration of ions
Electrons in the highest energy levels are lost first
Ionic bonding
Metal + non-metal
Electrons are transferred from the metal atom to the non-metal atom
Oppositely charged ions formed that attract each other
Covalent bonding
Two non-metals
Electrons are shared between the atoms and are attracted to the nuclei of both bonded atoms
Metallic bonding
Occurs in metals
Electrons are shared between all atoms
Ionic bonds
Electrons are transferred from the metal atom to the non-metal atom
Oppositely charged ions form, which are bonded together by electrostatic attraction
The metal ion is positive (cation)
The non-metal ion is negative (anion)
Giant ionic lattices
Each ion is surrounded by oppositely charged ions
These ions attract each other from all directions, forming a 3D giant ionic lattice
All ionic compounds exist as giant ionic lattices in the solid state but when aqueous, the ions are free to move so they can conduct when aqueous or molten
Ionic compound properties: high melting and boiling points
Ionic compounds are solids at room temperature- a large amount of energy is needed to break the strong electrostatic bonds that hold the oppositely charged ions together in the lattice. Therefore, ionic compounds have high melting and boiling points.
The greater the charge, the stronger the electrostatic forces between the ions so more energy is required to break up the ionic lattice during melting
Ionic compounds: Electrical conductivity
In a solid ionic lattice:
The ions can’t move - can’t conduct
When aqueous or molten:
The solid lattice breaks down - the ions are free to move
Can conduct
Why do ionic lattices dissolve in polar substances?
The slight charges within the polar substances are able to attract the charged ions in the giant ionic lattice. This means that the lattice is disrupted and the ions are pulled out of it.
Covalent bonds
The negatively charged shared pair of electrons is attracted to the positive charges of both nuclei
This attraction overcomes the repulsion between the two positively charged nuclei
Electrons are shared
Carbon will always make
4 covalent bonds
Nitrogen will always make
3 covalent bonds
Oxygen will always make
2 covalent bonds
Hydrogen will always make
1 covalent bond
In a dative covalent bond,
one atom provides both bonding electrons from a lone pair of electrons
For elements in groups 15-17 something weird happens from period 3. What is it?
After covalent bonding, one of the bonding atoms may finish up with more than eight outer shell electrons - “expansion of the octet”
Simple covalent structures
The atoms within each molecule are held together by strong covalent bonds
The different molecules are held together by weak intermolecular forces (London forces)
Simple covalent properties: Melting and boiling points
Low melting and boiling points:
Weak intermolecular forces
A relatively small amount of energy is needed to break them
Simple covalent properties: Electrical conductivity
Don’t conduct!
No charged particles free to move
Simple covalent properties: Solubility
Generally soluble in non-polar solvents:
Weak London forces are able to form between covalent molecules and the solvent
This helps the lattice break down and the substance dissolves
Giant covalent properties: Melting and boiling points
High melting and boiling points:
High temperatures needed to break the strong covalent bonds within the lattice
Giant covalent properties: Electrical conductivity
Can’t conduct!
No free charged particles (graphite is an exception)
Giant covalent properties: Solubility
Insoluble in both polar and non-polar solvents:
The covalent bonds in the lattice are too strong to be broken by either polar or non-polar solvents
Linear
2 bonded pairs around central atom
180*
Trigonal planar
3 bonded pairs around central atom
120*
Tetrahedral
4 bonded pairs around central atom
109.5*
Trigonal bipyramid
5 bonded pairs around the central atom
90* and 120*
Octahedral
6 bonded pairs around the central atom
90*
Each lone pair reduces the bond angle by
2.5*
A lone pair repels more than a bonded pair
What is the Pauling scale?
A scale used to measure the electronegative of an atom
What is electronegativity?
The attraction of a bonded atom for the pair of electrons in a covalent bond
Bonds between atoms with different electronegativity values result in
Polar bonds
Relative strength of bond types going from least to most
London dispersion forces
Permanent dipole-dipole
Hydrogen bonds
Ionic and covalent bonds
The size of London dispersion forces increases with increasing
numbers of electrons. The greater the number of electrons, the larger the induced dipoles and the greater the attractive forces between molecules.
A hydrogen bond is a strong permanent dipole-permanent dipole attraction between
An electron deficient hydrogen atom on one molecule
and
A lone pair of electrons in a highly electronegative atom (N, O, F, Cl) on a different molecule
Why is ice less dense than water?
Ice has an open lattice structure
Particles are further apart
Why does water have a higher melting point and boiling point than expected?
Hydrogen bonds are strong so take more energy to break
Why does water have a very high surface tension/is more viscous than expected?
Because strong hydrogen bonds hold the surface together
