Chapter 7 Periodicity Flashcards
What is periodicity?
The repeating pattern of properties across a period
What are the vertical columns on the periodic table known as?
Groups
What are the horizontal rows on the periodic table known as?
Periods
How are elements arranged in order of within the periodic table?
Atomic number
What do the groups on the periodic table tell us?
The group number shows the number of electrons in the outer shell, therefore this means that elements in the same group have similar chemical properties.
What do the periods on a periodic table tell us?
Tells us the number of the highest energy electron shell in an elements atom
What is the trend across a period in the periodic table?
Each period starts with an electron in a new highest energy shell.
For example, across period 2, the 2s sub shell fills with two electrons, followed by the 2p sub-shell with six electrons
What is first ionisation energy?
The energy that is required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous +1 ions.
Eg Na (g) —> Na+ (g) + e-
What are the factors that affect ionisation energies?
Atomic radius
Nuclear charge
Electron shielding
How does atomic radius impact ionisation energy?
The greater the distance between the nucleus and outer electrons, the larger the atomic radius. Therefore, the electron being removed is further away, so there is less attraction between the nucleus and the electron being removed.
How does nuclear charge impact ionisation energy?
The more protons there are in the nucleus, the greater the nuclear charge. Therefore, the stronger the attraction between the nucleus and the electron being removed is.
How does electron shielding impact ionisation energy?
The more inner shells of electrons there are between the nucleus and the outer electrons, the greater repulsion there is between the inner electrons and the electron being removed. This reduces the attraction between the nucleus and electron being removed so first ionisation energy decreases.
How many ionisation energies does an element have?
An element has as many ionisation energies as there are electrons.
Explain the ionisation energy of Helium.
Helium has two ionisation energies, however the second ionisation energy is greater than the first ionisation energy. This is because there are two protons attracting the two electrons in the 1s sub-shell. After the first electron is lost, the single electron is pulled closer to the nucleus due as there is a stronger nuclear charge. Therefore, more ionisation energy is needed to remove the second electron.
What is the second ionisation energy?
The energy that is required to remove one mole of electrons from one mole of +1 gaseous ions to produce one mole of gaseous +2 ions.
What do successive ionisation energies allow us to predict?
The number of electrons in the outer shell based on the number of electrons before the large jump signalling the move to the next shell
The group of the element on the periodic table
The identity of the element
What is the trend in first ionisation energies down a group?
First ionisation energies decrease down a group.
This is because the atomic radius increases and there are more inner shells so shielding also increases. Therefore, the nuclear attraction to outer electrons decreases.
What is the trend in first ionisation energies across a period?
First ionisation energies increase across a period. This is because the nuclear charge increases as an extra proton is added each time you move across the period, there is the same number of shells so similar electron shielding occurs , due to this the nuclear attraction between the nucleus and outer electrons is increased so the atomic radius decreases.
Why is there a drop in ionisation energy from group 2 to group 3?
This marks the start of filling the p sub-shell in the Group 3 element which is of higher energy than that of the s sub-shell in the group 2 element. Therefore, the p sub shell electron of the group 3 element is easier to remove than one of the s sub shell electrons in the group 2 element.
The first ionisation energy of group 3 is less than the first ionisation energy of group 2.
Why is there a drop in ionisation energy from group 5 to group 6?
Moving from group 5 into group 6, the electrons begin to pair up in a p sub shell, this increases repulsion between the electrons therefore making it easier to remove an electron from a group 6 element than a group 5 element.
The first ionisation energy of a group 6 element is less than the first ionisation energy of a group 5 element
What is metallic bonding?
The electrostatic forces of attraction between the positive metal cations which are fixed in position and the negative sea of delocalised electrons that are mobile.
This forms a giant metallic lattice
What are the properties of metals?
High electrical conductivity- electrons are mobile and free to carry the charge when solid or liquid
High melting and boiling points- due to the electrostatic attraction between the cation and delocalised electrons which require a lot of energy to overcome
Insoluble- as the metallic bonds are too strong to be overcome by the interactions of a solvent
They don’t dissolve in polar solvents as the interactions between the solvent and metal would lead to a reaction
What non metals form giant covalent lattices?
Boron
Silicon
Carbon
What are the properties of Giant covalent lattices?
High melting and boiling point- due to the strong covalent bonds which require a lot of energy to overcome
Insoluble in almost all solvents- the covalent bonds are too strong to be overcome by the interaction of solvents
Poor conductivity (with the exception of Graphite and Graphene) - as there are no free mobile charge carriers.
Explain the drop in ionisation energy between Nitrogen and Oxygen
In nitrogen and oxygen the highest energy electrons are in a 2p sub-shell. In nitrogen, all three p orbitals are singularly occupied however in oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen than a nitrogen atom.
Therefore, the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen
Explain the drop in first ionisation energy between Beryllium and Boron
This marks the start of filling the 2p sub shell in Boron, which is of higher energy that that of the 2s sub shell in Beryllium. Therefore, itv is easier to remove an electron from the 2p orbital of Boron than it is the 2s orbital of Beryllium.
This means Boron has a lower first ionisation energy than Beryllium
What is the structure of diamond?
Giant covalent lattice
Four carbon atoms are covalently bonded to four other carbon atoms forming a tetrahedral shape, bond angle 109.5.
What elements have a simple molecular structure across period 2 and 3?
Nitrogen Phosphorus
Oxygen Sulfur
Fluorine Chlorine
Neon Argon
What elements have a giant metallic structure across period 2 and 3?
Lithium
Sodium
Beryllium
Magnesium
Aluminium
How do you know when a substance is simple molecular or giant covalent?
If it is a gas, liquid or a low melting point solid then it is a simple molecular substance. However, if it is a high melting point solid then it is giant covalent.
Do Giants covalent structures form intermolecular forces?
No, they only have covalent bonds.
Explain why the melting point of Sulfur is higher than that of Chlorine.
Sulfur exists as S8 molecules and Chlorine as Cl2 molecules. This means that the London Forces are stronger between S8 molecules as there are more electrons than Cl2 molecules.
Why is the Group 4 element Germanium a good conductor of electricity?
Has a giant metallic lattice rather than a giant covalent lattice like Carbon and Silicon, therefore is has negative delocalised electrons which are free to move and carry the charge throughout the structure.
Why does Arsenic have a higher melting point than nitrogen and phosphorus?
Has a giant covalent lattice whereas in period 2 and 3, group 5 had a simple molecular lattice.
Explain, in terms of its structure, whether you would also expect diamond to be a good conductor of electricity.
No because it has a giant covalent lattice structure so all outer electrons are used up in covalent bonds.
Explain whether you would expect gallium to be soluble or insoluble in water.
Insoluble because the metallic bonds are too strong to be overcome by the interactions of a solvent.
Explain in terms of bonding why Magnesium has a higher melting point than sodium
Magnesium has more delocalised electrons within its metallic lattice attracting the positive 2+ cations so the electrostatic attraction requires a lot more energy to overcome in the Magnesium atom than the Sodium atom.
Describe the structure and bonding of aluminium. Include the names of the particles involved in the bonding within your answer.
Aluminium has a giant metallic lattice structure. This is due to its positive 3+ cations which are fixed in position attracting the negative delocalised electrons which are mobile and free to move around the structure. This forms the metallic structure due to the electrostatic attraction between the positive cations and negative delocalised electrons.
Explain why the first ionisation energy of nitrogen is higher than the first ionisation energy of carbon.
Because across a period there is a general increase in first ionisation energy, the nuclear charge increases as nitrogen has more protons in its nucleus. The number of shells remains the same so inner electron shielding stays similar, however the increased nuclear attraction reduces the atomic radius, pulling the outer electrons closer slightly. Therefore, there is a stronger attraction between the nucleus and the electron being removed of Nitrogen than Carbon.
The variation in first ionisation energies across a period of the Periodic Table provided evidence for what structures within an atom?
Sub shells
From the following elements: Lithium, Beryllium and Fluorine.
Predict which one will have the largest second ionisation energy. Explain your answer.
Lithium
The atomic radius decreases as the only electron is the outer shell is lost
-therefore, there is only 2 electrons remaining in the final shell
-this decreases the amount of inner electron shielding so it has little effect
-there is a greater nuclear attraction between the nucleus and the electron being removed so second ionisation energy increases
