B2 redox and periodicity Flashcards
Oxidation number of element
0
Oxidation number of group 1
1+
Oxidation number of group 2
2+
Oxidation number of fluorine
1-
Oxidation number of oxygen in H2O2
1-
Oxidation number of hydrogen in H2O2
1+
What substance gets the negative oxidation number in a compound
The most electronegative
Sum of oxidation numbers in a compound
0
Sum of oxidation numbers in an ion
Charge of ion
What happens to oxidation number when a substance is reduced
Decreases
What happens to oxidation number when a substance is oxidised
Increases
When writing systematic names, when does the Roman numeral oxidation number come
In brackets after the element
e.g Titanium (IV) Oxide
In metal and acid reactions, what is reduced and what is oxidised
Metal oxidised
Hydrogen reduced
Meaning of oxidation
Loss of electrons
Increase in oxidation number
Meaning of reduction
Gain of electrons
Decrease in oxidation number
Define Disproportionation
Reduction and oxidation of the same element in the same reaction
Define Oxidising Agent
Takes electrons from the species being oxidised
Define Reducing Agent
Adds electrons to the species being reduced
How to write elemental half-equations
Calculate overall charge on each side
Add electrons to more positive side
Ensure fully balanced
How is the periodic table arranged
Increasing atomic number
How to write half-equations for acidic conditions
Balance the element changing oxidation state
For every oxygen, add an H2O to the other side
For every hydrogen, add an H+ to the other side
Use electrons to ensure both sides have the same charge
What do elements in periods in the periodic table have
Repeating trends in physical and chemical properties
What do elements in groups of the periodic table have
Similar chemical properties
Define Periodicity
The repeating trends in physical and chemical properties across different periods
How to explain why an element is in the s/d/p block
It’s highest energy electron is in the s/d/p subshell
Define First Ionisation Energy
The energy required to remove 1mol of electrons from 1mol of gaseous atoms
How to explain differences in ionisation energies
Electron shells
Atomic radius
Shielding
Nuclear charge
Shielding + atomic radius outweigh nuclear charge
Nuclear attraction
What must always be stated when describing changes in ionisation energies
Shielding - even if has no effect
When given an ionisation energy graph for an element, how can you tell what group that element is in
Greatest jump in ionisation energies
Jump 4 - 5
In group 4
General equation for first ionisation energy
General equation for second ionisation energy
Draw a graph for ionisation energies across period 2
What element has highest first ionisation energy
Helium
On first ionisation energy graphs, there is a general increase - what are the exceptions
New shell - decrease
p subshell - decrease
group 6 - decrease (new electron in singular paired orbital so more repulsion)
How are elements arranged in the periodic table?
In order of increasing atomic number.
What is the first ionisation energy?
The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
What are the 3 factors that affect ionisation energy?
- Atomic radius
- Nuclear charge
- Electron shielding
How does atomic radius affect ionisation energy?
The greater the distance between the nucleus and the outer electrons, the less the nuclear attraction. The force of attraction decreases with increasing distance.
How does nuclear charge affect ionisation energy?
The more protons in the nucleus, the greater the attraction between the nucleus and the outer electrons.
How does electron shielding affect ionisation energy?
Negatively charged outer-shell electrons repel inner shell electrons. The shielding effect increases and reduces the attraction between the nucleus and the outer electrons.
What determines how many ionisation energies an electron has?
How many electrons it has.
E.g., helium has 2 electrons, and therefore 2 ionisation energies.
What is the second ionisation energy?
The energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.
What can the values of successive ionisation energies mean?
They can tell us how many electrons are in the outer shell and what group the element is in.
e.g - the largest increase is between the 3rd and 4th ionisation energies:
- shows 4th electron is being removed from an inner shell
- therefore, there are 3 electrons in the outer shell
- element must be in group 3
What is the trend in ionisation energy down a group?
Decrease down the group.
Why does ionisation energy decrease down a group?
- Atomic radius increases.
- More inner shells => shielding increases.
- Nuclear attraction to outer electrons decreases.
- The effect of increasing nuclear charge is outweighed by the increased radius and increased shielding.
What is the trend in ionisation energy across a period?
General increase.
Why does ionisation energy generally increase across a period?
- Nuclear charge increases.
- Same shell: similar shielding.
- Nuclear attraction increases.
- Atomic radius decreases (outweighed).
What are the exceptions to the general increase in first ionisation energy across a period?
- Decrease between beryllium and boron.
- Decrease between nitrogen and oxygen.
- Decrease between the last element of one period and the first element of the next.
Why does first ionisation energy decrease between beryllium and boron?
The electron in boron is coming from a higher energy subshell (2p vs. 2s). Boron needs less extra energy to remove the electron.
Why does first ionisation energy decrease between nitrogen and oxygen?
In O2, the electron removed is a paired electron in the 2p orbital. Electrons repel each other in the 2p orbital, so less energy is required to remove the electron.
Why does first ionisation energy decrease between the last element of one period and the first element of the next?
Increased distance between the nucleus and outer shell electrons and increased shielding by inner electrons outweigh the increased number of protons.
Describe metallic bonding.
The electrostatic attraction between cations and delocalised electrons. The cations consist of the nucleus and the inner electron shells of the metal atoms.
Describe metallic structure.
The cations are in fixed positions. The delocalised electrons are mobile and are able to move throughout the structure, forming a giant metallic lattice.
Describe metals’ electrical conductivity.
Metals conduct electricity in solid and molten states because the delocalised electrons can move through the structure, carrying charge.
Describe the melting and boiling points of metals.
Most metals have high melting and boiling points. The melting point depends on the strength of the metallic bonds, and high temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic attraction.
Describe the solubility of metals.
Metals do not dissolve in water.
What substances form giant covalent lattices?
- Graphite
- Diamond
- Silicon dioxide
Describe the melting and boiling points of giant covalent lattices.
They have high melting and boiling points because covalent bonds are strong. High temperatures are necessary to provide the large quantity of energy needed to break the strong covalent bonds.
Describe the solubility of giant covalent lattices.
They are insoluble in almost all solvents because the covalent bonds holding together atoms in the lattice are too strong to be broken by interactions with solvents.
Describe the electrical conductivity of giant covalent lattices.
Diamond and silicon dioxide cannot conduct electricity (no delocalised electrons), while graphite can conduct electricity.
Describe graphite.
Graphite is composed of parallel layers of hexagonally arranged carbon atoms. The layers are bonded by weak London forces, and the bonding only uses 3/4 of carbon’s 4 outer shell electrons. The spare electron is delocalised between the layers, allowing electricity to be conducted.
Use your answer to is to explain why electron configuration is an example of a period trend (electron configuration and block)
across period 2, the 2s subshell fills first, followed by the 2p
same pattern or trend of filling the sub shell related in other periods
why does successive ionisation energy increase with ionisation number
radius decreases
attraction between the electrons and the nucleus increases
Explain how the successive ionisation energies provide evidence for the electron shells in sodium atoms
large increase between …. and… and …
large increase shows new shell
Explain how the first ionisation energies of mg (larger) and al give further details of electron structure
Mg has (outer) electron in (3)s sub-shell
Al has (outer) electron in (3)p sub-shell
(3)p sub-shell has higher energy than (3)s sub-shell
explain why potassium is placed immediately after argon in periodic table
potassium has 1 more proton
how could the student obtain a sample of magnesium phosphate after reacting mg and phosphoric acid?
filter to obtain ppt
dry the soil and evaporate to remove water
No3 overall charge for oxidation numbers
-1
explain why the first ionisation energy of mg is greater than the first ionisation energy of Sr
mg removes electrons from shell closes to nucleus/ smaller atomic radius
Greater nucalr attraction
Explain why second ionisation erenrgy of Sr is greater than fist ionisation of Sr
2nd electron removed from cation/ positvelty charged oops
protons electron ratio in 1+ ion is greater than atom
greater nuclear attraction
explain why iodine is less reactive than bromine
larger atomic radius
greater shielding
weaker nuclear attraction
disporopration
oxidation and reduction occurring of same element
