Module 3: Chapter 7 - Periodicity Flashcards
Describe Mendeleevs periodic table
Mendeleevs periodic table consisted of 63 elements arranged in order of atomic mass. He lined up the elements in groups with similar properties. If the group properties did not fit, Mendeleev swapped around elements and left gaps assuming that these elements were yet to be discovered. As new elements were discovered which filled these gaps, their properties matched those that Mendeleev predicted causing his table to become accepted by the wider scientific community.
How many elements have been discovered (as of 2023)?
118 Elements (A level periodic table only includes 114)
How many periods are there in the periodic table?
7 horizontal periods
How many groups are there in the periodic table?
18 vertical groups
How are elements in is the periodic table arranged?
Reading from left to right, elements are arranged in order of increasing atomic number, each successive element has 1 extra proton. These elements are then organised into vertical columns called groups and horizontal rows called periods.
How are elements arranged into groups in the periodic table?
The elements are aranged into vertical columns called groups. Each element in a group has atoms with the same number of valence electrons and similar chemical properties
What is the valence shell?
The outer shell
How are elements arranged into periods in the periodic table?
The elements are arranged in horizontal rows called periods. The number of the period gives the number of the highest energy occupied electron shell in an elements atom
What is periodicity?
The repeating trend in properties of the elements across each period
What are examples of periodicity of properties?
- Electron configuration
- Ionisation energy
- Structure
- Melting points
What is the periodic trend in electron configuration across a period?
Each period starts with an electron in a new highest energy shell:
* Across period 2, the 2s sub-shell is filled with 2 electrons, followed by the 2p subshell with 6 electrons
* Across period 3, the same pattern of filling is repeated for the 3s and 3p subshells
* Across period 4, the 4s shell is filled, followed by the 3d subshell and then the 4p subshell (although the 3d subshell is involved, the highest occupied energy shell is n=4)
What is the periodic trend in electron configuration down a group?
Elements in each group have the same number of valence electrons, and also contain the valence electrons in the same subshell (s,p,d,f). This similarity in electron configuration gives elements in the same group their similar chemistry
How are elements in the periodic table divided into blocks?
The elements in the periodic table can be split into blocks corresponding to their highest energy sub-shell. This gives 4 distinct blocks (s,p,d,f)
How can the sub-shell blocks of the periodic table be used to determine electronic configuration?
The electron configurationmof an element would be all the previous subshells completely filled, and then the highest energy subshell filled with how far the element is into the subshell row. (i.e Phosphorus is 1s2 2s2 2p6 3s2 3p3 as it is 3 into the 3p row of the p block)
What are the new numbering groups compared to the old numbering groups?
- The old numbers were 1-7 then 0 (excluding transition metals), this was based on the s and p blocks
- The new numbers are 1-18 (including the transition metals), this numbers each column in the s, d, and p sequentially
What is the name of group 1?
Alkali metals
What is the name of group 2?
Alkaline earth metals
What is the name of groups 3-12?
Transition metals
What is the name of group 15?
Pnictogens
What is the name of group 16?
Chalcogens
What is the name of group 17?
Halogens
What is the name of group 18?
Noble gases
What are metalloids?
An element whose properties are an intermediate between those of a metal and non-metals
What is ionisation energy?
A measure of how easily an atom loses electrons to form positive ions in kJ Mol⁻¹
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 factors affect ionisation energy?
- Atomic Radius
- Nuclear Charge
- Electron Shielding
How does atomic radius effect ionisation energy?
The greater the distance between the nucleus and the outer electrons the less the nuclear attraction. The force of attraction falls off sharply with increasing distance, so atomic radius has a large effect. As atomic radius increases the ionisation energy decreases
How does nuclear charge effect ionisation energy?
The more protons there are in the nucleus of an atom, the greater attraction between the nucleus and the outer electrons. As Nuclear charge increases, ionisation energy increases
How does electron shielding effect ionisation energy?
Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons. This repulsion, called the shielding effect, reduces attraction between the nucleus and the outer electrons. As electron shielding increases, ionisation energy decreases
How many ionisation energies does an element have?
as many ionisation energies as there are electrons
Why do successive ionisation energies increase in energy?
After the first electron is lost, the remaining electrons are pulled closer to the nucleus. The nuclear attraction on the remaining electrons increases and more ionisation energy will be needed to remove this second electron.
What is the second ionisation energy?
The energy required to remove one electron from each ion in one mole of gasous 1+ ions of an element to form one mole of gaseous 2+ ions
How can you determine the number of valence electrons from successive ionisation energies?
A large increase in ionisation energies suggest that the electron has been removed from a different shell, closer to the nucleus, with less shielding
Are Ionisation energies endothermic or exothermic?
Endothermic
What predictions do successive ionisation energies allow to be made?
- The number of electrons in the outer shell
- The group of the element in the periodic table
- The identity of an element
Identify this element from period 3:
Aluminium
What is the general pattern in first ionisation energies across each period?
General increase
What is the general pattern in first ionisation energies between the end of one period and the start of the next?
Sharp decrease
What is the general pattern in first ionisation energies down a group?
First ionisation energies decrease down a group
Explain the trend in ionisation energies down a group
As you travel down a group, the atomic radius increases and there are more inner shells of electrons, causing the shielding effect to increase. This decreases the effective nuclear charge on outer electrons, causing the first ionisation energy to decrease
Explain the trend in first ionisation energies across a period
As you travel across a period, the nuclear charge increases, the valence electrons are in the same shell so the number of inner shells stays the same meaning that there is similar electron shielding between each element, and the atomic radius decreases. This causes the effective nuclear charge on the outer electrons to increase, causing the first ionisation energy to increase
Why is there a slight decrease in first ionisation energy between the second and third element of a period, such as between Be and B?
The fall in first ionisation energy between Berrylium and Boron marks the start of filling the 2p subshell. The 2p subshell in Boron has a higher energy than the 2s sub-shell in Berrylium. Therefore, in Boron, the 2p electron is easier to remove than one of the 2s electrons in Berrylium. Therefore the first ionisation energy decreases between Berrylium and Boron
Why is there a slight decrease in first ionisation energy between the fifth and sixth element of a period, such as between N and O?
The fall in ionisation energy between nitrogen and oxygen marks the start of electron pairing in the p-orbitals of the 2p subshell. In nitrogen and oxygen the highest energy electrons are in a 2p subshell. In oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom. Therefore the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen
What are metalloids?
Metalloids/semi-metals are elements near to the metal/non-metal divide. They can show in-between properties of metals and non-metals
What state are metals at room temperature?
All metals except mercury are solids at room temperature
What is the one property constant throughout all metals?
They are electrical conductors
Describe the formation and structure of a giant metallic lattice
In a solid metal structure, each atom donates its negative outer-shell electrons to form a sea of delocalised valence electrons which spreads out throughout the entire structure. The positive metal cations left behind consist of the nucleus and the inner shell electrons. The cations are fixed in position, maintaining the structure and shape of the metal. The delocalised electrons are mobile and are able to move throughout the structure, only the electrons move.
Why are metals so useful?
They have a rigid structure which is able to conduct. The cations are fixed in position, maintaining the structure and shape of the metal. The delocalised electrons are mobile and are able to move throughout the structure, only the electrons move.
What is metallic bonding?
The strong electrostatic attraction between positive metal ions and delocalised electrons
What is the ratio of metal cations to delocalised electrons in a group 1 metal?
1:1
What is the ratio of metal cations to delocalised electrons in a group 2 metal?
1:2
What is the ratio of metal cations to delocalised electrons in a group 3 metal?
1:3
What are the properties of most metals?
- Strong metallic bonds
- High electrical conductivity
- High melting and boiling points
Explain how metals can conduct
Metals conduct electricty in both solid and liquid states (unlike ionic substances which cannot conduct as a solid). When a voltage is applied across a metal, the delocalised electrons move through the structure carrying charge. This is a flow of charge and therefore a current.
Which metal has the highest melting point?
Tungsten (W)
What does the melting point of metals depend on and why do metals have high melting points?
The strength of the metallic bonds. For most metals, high temperatures are required to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and electrons. This strong attraction results in most metals having high melting and boiling points
What is the solubility of metals?
Metals are completely insoluble. They are completely insoluble in non-polar solvents and any interaction between a polar solvent and the charges in a metallic lattice would lead to a reaction.
Which non-metals form giant covalent lattice structures
- Boron
- Carbon
- Silicon
What is the structure of a diamond?
- Tetrahedral arrangement of carbon atoms
- All bond angles are 109.5° by EPRT
What are the melting and boiling points of a giant covalent lattice?
Giant covalent lattices have high melting and boiling points. This is because covalent bonds are strong. High temperatures are necesary to provide the large quantity of energy needed to break the strong covalent bonds
What is the solubility of giant covalent lattices?
Giant covalent lattices are insoluble in almost all solvents as the covalent bonds holding together the atoms in the lattice are far too strong to be broken by interacction with solvents
What is the electrical conductivity of Giant covalent lattices?
Giant covalent lattices are non-conductors of electricty. The only exceptions are graphene and graphite - allotropes of carbon - as they have delocalised electrons which are able to conduct.
What is the structure of graphite?
Graphite is composed of parallel layers of graphene. The layers are bonded by weak London forces.
What is the structure of graphene?
Graphene is composed of hexagonally arranged carbon atoms linked by strong covalent bonds. Graphene has a trigonal planar geometry meaning that each carbon atom only uses 3 of the 4 outer shell electrons in covalent bonding. The 4th electron becomes delocalised and allows graphene to conduct.
What are the unique properties of graphene?
- Graphene has the same electrical conductivity as copper
- It is the thinnest material ever made
- It is the strongest material ever made
Describe the periodic trend in melting points across periods 2 and 3?
- The melting points increase from group 1 to group 14 (4)
- There is a sharp decrease in melting points between group 14 (4) and group 15 (5)
- The melting points are comparatively lower from group 15 (5) to group 18 (0)
Explain the periodic trend in melting points across periods 2 and 3?
The sharp decrease in melting point after carbon and silicon respectively marks the change from giant structures to simple molecular structures. On melting giant structures have strong forces to overcome so have high melting points. Simple molecular structures have weak forces to overcome, so have much lower melting points.
What are the equations for the first 2 ionisation energies of sulfur?
S(g) -> S⁺(g) + e⁻
S⁺(g) -> S²⁺(g) + e⁻
Why is there a sharp drop in first ionisation energy between the end of one period and the start of another?
The sharp drop in first ionisation energy between the end of one period and the start of another is caused by the addition of a new inner shell. This causes both the atomic radius to increase and electron shielding to increase. The decreases the effective nuclear charge on the valence electrons, causing first ionisation energy to decrease.
