Chapter 7 Flashcards
How is the periodic table arranged?
By increasing atomic number, in groups and periods
What does the period number of an element tell you?
The number of the period gives the number of the highest energy electron shell in an element’s atom
What does the period number of an element tell you?
The number of the period gives the number of the highest energy electron shell in an element’s atom
What is periodicity?
The repeating trends in properties of elements across a period
How is the periodic table divided?
Into the:
s block
p block
d block
What does the position of an element, say about its electronic structure?
If an element is found in the s block, the highest energy sub shell would be the s sub shell.
How to deduce electron config. from an element’s position: (example)
Chlorine is found in the p block, therefore the highest energy sub shell is a p sub shell.
Chlorine is in the 3rd period, therefore the highest energy sub shell will be a 3p sub shell
Chlorine is in the 5th column of the p block, therefore the highest energy sub shell will contain 5 electrons.
Thus chlorine = 1s22s22p63s23p5
What is the periodic trend in electron configuration across period 2?
Across period 2, the 2s sub shell fills with two electrons, followed by the 2p sub shell with 6 electrons
What is the periodic trend in electron configuration across period 3?
Across period 3, the 3s sub shell fills with two electrons, followed by the 3p sub shell with 6 electrons
What is the first ionisation energy?
The first ionisation energy is 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
Na(g) –> Na+(g) + e-
What are the factors affecting ionisation energy?
Atomic radius
Nuclear charge
Electron shielding
What is ionisation energy?
It measures how easily an atom loses electrons to form positive ions
What does a large difference in ionisation energy suggest when looking at a graph of successive ionisation energies?
It suggests a change from one shell to another.
What predictions can be made about this period 3 element from a table of successive ionisation energies? (example)
The ionisation energies steadily increase but then there is a large increase between the third and fourth ionisation energies.
This shows that the fourth electron is being removed from an inner shell.
Thus, there are 3 electrons in the outer shell and the element must be in Group 3.
Since it is in period 3, the element must be aluminium.
Explain the trend in first ionisation energy down a group?
As you go down a group, atomic radius and electron shielding increases and nuclear attraction decreases, so therefore the first ionisation energy decreases.
Explain the trend in first ionisation energy across a period?
As you go across a period, nuclear charge increases, atomic radius decreases and electron shielding remains the same so therefore the first ionisation energy increases.
Why does the first ionisation energy decrease between beryllium and boron?
The fall in first ionisation energy from beryllium to boron marks the start of filling the 2p sub-shell.
The 2p sub-shell in boron has a higher energy than the 2s sub-shell in beryllium.
Therefore, in boron the 2p electron is easier to remove than one of the 2s electrons in beryllium and so the first ionisation energy of boron is less than of beryllium.
Why does the first ionisation energy decrease between nitrogen and oxygen?
The fall in first ionisation energy from nitrogen to oxygen marks the start of electron pairing in the p orbitals of the 2p sub shell.
In nitrogen and oxygen, the highest energy electrons are in the 2p sub shell
In oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electrons from an oxygen atom than a nitrogen atom.
Therefore the first ionisation energy of oxygen is less than that of nitrogen.
What is metallic bonding?
The strong electrostatic attraction between cations and delocalised electrons
What is a giant metallic lattice?
When billions of metals atoms are held together by metallic bonding
What are the properties of metals?
Strong metallic bonds
High electrical conductivity
High melting and boiling points
Why do metals have a high electrical conductivity?
They are able to conduct electricity because when a voltage is applied across the metal, the delocalised electrons can carry the charge through the structure
Why do metals have a high melting and boiling point?
The melting point depends on the strength of the metallic bonds holding together the atoms in the giant metallic lattice
High temperatures are necessary to provide a large enough energy to overcome the strong electrostatic attraction between the cations and electrons, resulting in high melting and boiling points
What is a giant covalent lattice?
Billions of atoms which are held together by a network of covalent bonds
What are the typical properties of giant covalent structures?
High melting and boiling points
Insoluble in almost all solvents
Non-conductors of electricity (apart from graphite and graphene)
Why do giant covalent lattices have high melting and boiling points?
Because they contain strong covalent bonds which require high temperatures in order to overcome
Why are giant covalent lattices insoluble?
The covalent bonds present are far too strong to be broken by interaction with solvents
Why are most giant covalent lattices non conductors of electricity?
In carbon (diamond) and silicon, all four outer shell electrons are involved in covalent bonding, so none are available for conducting electricity
Why are graphite and graphene conductors of electricity?
Only three electrons of the four outer-shell electrons are used in covalent bonding. The remaining electron is released into a pool of delocalised electrons, which carries charge through the structure
What is the periodic trend in melting point across period 2 and 3?
Across both periods:
The melting point increases from group 1-4
There is a sharp decrease in melting point between group 4 and group 5
The melting points are comparatively low from group 5 to group 8
What does the sharp decrease in melting point between group 4 and 5 suggest?
Suggests a change from giant structures to simple molecular structures
Explain the periodic trend in melting point across period 2 and 3.
The variation in melting point across both periods is a result of a different structure - from giant structures to simple molecular structures.
On melting, the 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.
