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Terms in this set (38)
Levels of electronic structure
- Shell/energy level
- Energy rises with shell number
- Shell number is also known as principal shell number ("n")
- The shell number corresponds to the number of subshells in that shell
- The formula for the number of electrons in that shell is 2n²
- The number of electrons that fill each shell (in ascending shell number) 2, 8, 8, 18
These contain orbitals. The shell number of a shell corresponds to the number of sub-shells it contains e.g. shell 4 contains 4 sub-shells.
- s-subshells contain one orbital
- p-subshells contain three orbitals
- d-subshells contain five orbitals
- f-subshells contain seven orbitals
As shell 1 only contains one subshell, it just contains a 1s orbital. Shell 2 has two subshells so it contains a 2s and 2d orbital.
What letter order to subshells go in and what is the exception?
Usually subshells are filled in the order s, p, d and f. However, the 4s subshell is actually filled before the 3d one and it is emptied before 3d when electrons are lost during ionisation as the 3d subshell has a higher energy. Despite this, when writing electron configuration, it is important that ALL the 3 subshells go before the 4 subshells.
- Region around the nucleus that can hold up to 2 electrons, provided the electrons have opposite spins
- Volume of space in which there is a high probability of finding an electron
- Normally drawn as a box containing two arrows, one pointing up and one pointing down (electrons with opposite spins)
- As orbitals hold two electrons, even though an s-subshell can only hold one orbital, it actually holds TWO electrons
Orbital filling rules
- In p, d, and f subshells, the electrons go into each orbital singly first and only pair up when they have to as there is a degree of repulsion between two electrons in the same orbital
- Elements that don't have two electrons in the same orbital of their outer shell are generally more stable due to the lack of repulsion
- Orbitals with the lowest energy are filled first (2s before 2d)
Periodic table 'blocks'
- Group 1 and 2 are s-block because their outer electrons go into s-orbitals
- Transitional metals are d-block because their outer electrons go into d-orbitals
- Group 3-8 are p-block because their outer electrons go into p-orbitals (except helium, which is s-block)
- The lanthanoids and actinoids are f-block because their outer electrons go into f-orbitals
Ways of writing electron configuration
- Full electron configuration, where the subshell filled is written and the number of electrons in each subshell is written next to it in superscript e.g. sodium is 1s²2s²2p⁶3s¹
- As sodium has 11 electrons the superscript numbers add up to 11 (2+2+6+1) and the subshell where they are found is written before it
- Shorthand notation of electron configuration is written as the noble gas found before the element and whatever additional electrons the element has e.g. sodium would be [Ne]3s¹ as neon is the last noble gas before sodium and sodium has one more electron than neon, found in the 3s subshell.
- Shorthand notation should only be used when the question asks for it
Electron configurations of ions
- To work these out just add (anions) or take away (cations) the right number of electrons from the configuration
- Don't forget that the 4s subshell is emptied BEFORE the 3d subshell
- AKA inert gases
- Group 8
If the outer shell of an atom has 8 electrons, it is 'happy'.
- It has a stable arrangement
- It isn't looking for anymore electrons (so it is largely unreactive)
The result of the electrostatic attraction between a positively charged ion and a negatively charged ion due to a transfer of electrons between atoms.
Properties of ionic compounds
- Usually made up of a metal and a non-metal
- High melting point and boiling point due to the high amount of energy needed to overcome the electrostatic interactions
- Usually solid at room temperature
- Can conduct electricity when molten or aqueous as the ions are free, mobile and able to carry charges whereas when solid they are in a fixed position
- Form a giant lattice
Giant ionic lattice
Repeating pattern of positive and negative ions
The number of ions of the opposite charge that an ion is surrounded by in a lattice.
What factors affect the strength of an ionic bond?
- As the charges on the ions increase, the strength of the bond increases e.g. MgO (Mg2+ and O2-) has a stronger ionic bond than NaCl (Na+ and Cl-)
- The smaller the ion, the higher the strength of the ionic bonds e.g. CaF2 has a higher boiling point than CaCl2 as the fluorine atom is smaller than the chlorine atom.
How does water dissolve ionic lattices?
- The cation (metal) is attracted to the delta minus charge on the oxygen atom in a water molecule
- The anion (non-metal) is attracted to the delta plus charges on the hydrogen atoms
- In solution the anion is surrounded by water molecules which orient themselves so the positive charges face the ion
- In solution the cation is surrounded by water molecules which orient themselves so the negative charges face the ion
- This pulls the two ions apart, weakening the bond between them
- Substances only dissolve if the HYDRATION (attraction of water around the ion) releases more energy than is needed to break the electrostatic forces
Why can't ionic compounds conduct electricity when solid?
- The ions are in a fixed position in a giant ionic lattice
- They are not mobile/free to move so they are unable to carry a charge
Why can ionic compounds conduct electricity when molten/aqueous?
-The ions are mobile/free to move
- They are able to carry a charge
Why are metals able to conduct electricity when solid and molten?
Metals are held together by metallic bonds, positive charges in a sea of delocalised electrons. These delocalised electrons are able to carry a charge both when the metal is solid and liquid.
The electrostatic attraction between two positive nuclei and a shared pair of electrons, produced by two orbitals, each containing one electron, overlapping.
Where are covalent bonds usually present?
- Between non-metals
- Non-metallic diatomic molecules (e.g. N2, H2, O2, halogens)
- Polyatomic ions
- Compounds of non-metallic elements (e.g. CO2, H2O, NH3)
How do covalent bonds work?
- Covalent bonds are localised
- There is an electrostatic attraction present between the shared pair of electrons and the positively charged nuclei of each of the bonding atoms.
- Bonded atoms often have outer shells with the same electron structure as the nearest noble gas.
Difference between ionic and covalent bonds
- Covalent bonds are localised between two atoms
- Ions attract ALL oppositely charges ions in ALL directions resulting in a large lattice
Number of bonds usually made by different atoms
Hydrogen - 1
Oxygen - 2
Nitrogen - 3
Carbon - 4
When an atom in a covalent compound has less than 8 electrons in its outer shell e.g. boron in BF3 only has 6
Expansion of the octet
When an atom in a covalent compound has more than 8 electrons in its outer shell e.g. sulfur has 12 in SF6
Double covalent bond
Two shared pairs of electrons e.g. O2
- Stronger than single covalent bonds
Triple covalent bond
Three shared pairs of electrons e.g. N2
- Stronger than single and double which makes N2 very stable and unreactive.
- A pair of electrons on an atom in a covalent compound that is not used for bonding/is not shared
- Represented by a pair of dots in a displayed formula
Dative covalent bond
When both bonding electrons in a covalent bond are donated from the same atom e.g. oxygen in carbon monoxide
- Represented by an arrow pointing at the atom that the electrons are being donated to in a displayed formula (also known as co-ordinate bond)
Melting points and boiling points of covalent compounds
- Very low so many are gases or liquids at room temperature
- Although covalent bonds are stronger than ionic bonds, the intermolecular forces between each molecule of the covalent compound are weak, so easy to overcome with very little heat
Average bond enthalpy
- A measure of covalent bond strength
- How much energy is needed to make/break a bond
What shape is an s-orbital?
What shape is a p-orbital?
Dumbbell shaped (like the number 8)
How can you use the periodic table to work out the electron configuration of an element?
- The row (period) the element is in tells you the outer shell number
- The block it's in tells you which subshell the outermost electrons are in
- How far the element is into a block tells you how many electrons are in that outer subshell
- For example, sulfur is in period 3 and it is the 4th element into the p-block. Therefore the last part of its electron configuration is 3p4. If you know this, you can just assume all the subshells before this one are full.
How are ionic bonds represented by a dot and cross diagram?
The symbol of the cation is written inside brackets with the charge of the ion in superscript outside the bracket. The outer shell arrangement of the anion is drawn inside brackets with dots and crosses showing where each electron originally came from, and the charge of the anion outside of the bracket.
How are covalent bonds represented by a dot and cross diagrams?
The chemical symbol of each atom is written and the shared electrons are represented by dots or crosses between them, and the rest of the outer shell of each atom is filled in.
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