Chapter 7 Periodicity Flashcards
How did Mendeleev arrange his table?
In increasing atomic mass
Groups with similar chemical properties
With spaces for elements he predicted were undiscovered
Switching some elements to better fit chemical trends, e.g tellurium and iodine
What did Mendeleev’s table not include?
The noble gases
Transition metals
How is our current periodic table arranged?
Increasing atomic number
In groups which have the same number of electrons in the same type of orbital in the outer shell
Periods which have the same number of shells
Define periodicity
Repeating trends in chemical, physical and atomic properties of elements across a period
Name 4 properties which exhibit periodicity
Electron configuration
Ionisation energy
Structure
Melting points
How does electron configuration exhibit periodicity?
From left to right, fills from the s to the p sub shell (with transition metals d sub shell also involved)
Each of the blocks have the same type of orbital holding the highest energy electron
What is the name of the groups?
Group 1= Alkali Metals
Group 2= Earth Alkali Metals
(Group 3= Boron groups/ triels)
(G4= Carbon family/ crystallogens)
(Group 5= Pnictogens)
(Group 6= Chalcogens)
Group 7= Halogens
Group 8= Noble Gases
Define the first ionisation energy
The amount of energy required to remove one electron form each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
What factors affect ionisation energy and how?
Atomic radius- increased atomic radius increases the distance between the nucleus and the outer shell, reducing the nuclear attraction
Nuclear charge- more protons, greater nuclear attraction
Electron shielding- greater amount of inner shells, greater repulsion from these upon outer shells, reduces nuclear attraction
What is the trend in successive ionisation energies?
Each successive ionisation energy will be greater than the last
Why do successive ionisation energies increase?
Fewer elections remain for the same positive nuclear charge, also same amount of electron shielding
The remaining electrons are pulled slightly closer to the nucleus
Greater nuclear attraction
Therefore, more energy needed to remove remaining electrons
How do successive ionisations give evidence for the shell theory? How can it be used to identify an element?
Sharp increase in ionisation energy at a point. Suggests electrons now being removed from a shell closer to the nucleus, has a lower energy, more strongly attracted to the nucleus, harder to remove
The ionisation number before the sharp increase shows the number of electrons in the outer shell, thus group
How does ionisation generally change across the period and why?
Increasing nuclear charge from protons being added
The amount of electron shielding is similar as no new shells are being added
Atomic radius decreases slightly as the outer shell is slightly greater attracted
Therefore, the nuclear attraction upon the outer shell increases, a greater amount of energy needed to remove an electron
How does ionisation energy change down a group and why?
A new shell is added at each period. There is a greater number of inner shells, increasing the amount of electron shielding.
The atomic radius also increases.
Therefore the nuclear attraction upon the outer shell decreases with less energy needed to remove an electron
(Although nuclear charge increases, increase not at a high enough proportion to be the dominant effect)
Why does the first ionisation energy fall from Be to B?
Highest energy electron in B is in the higher energy p orbital rather than the s orbital like in Be
Less energy needed to remove the electron, so a lower IE1
(make sure to specify higher energy and type of orbital for both)
Why does the first ionisation energy decrease from N to O?
Doubly occupied p orbital in O rather than singly occupied like in N
Added repulsion from pairing makes it easier to remove the highest energy electron, requiring less energy so a lower IE1
Where is the metal/non metal boundaries in the periodic table?
From Al to Po
Elements along this boundary are metals
Define metallic bonding
The strong electrostatic force of attraction between metal cations and the sea of delocalised elections
How does metallic bonding arise?
The outer shell of the metals delocalise, forming metal cations and a sea of delocalised electrons
In solid, in a giant metallic lattice, where the electrons can move throughout the structure, and cations in fixed positions
Why does the boiling point of the metals increase across the period?
Greater the charge of the cation and greater the number of delocalised electrons
Therefore, stronger metallic bonds, requiring more energy to overcome, so a greater boiling point
(Atom size is similar)
Do metals dissolve?
No, instead they react
Why can metals conduct electricity?
Delocalised electrons act as mobile charge carriers, carrying charge throughout the structure
What is the structure of group 4 elements normally? What properties do they possess?
Giant covalent
Insoluble, no electrical conductivity (- graphite/graphene)
Very high boiling points to overcome very many covalent bonds
What is a giant covalent structure?
A network of very many regular, repeating covalent bonds
What is the trend in melting point across periods and why?
Increases till group 4, then decreases (May then increase and decrease)
Stronger metallic bonding form 1-3
Very stronger giant covalent at 4
Much weaker London forces after 4
How does structure exhibit periodicity?
Metallic then giant covalent, then simple London forces
What is more useful for determine metallic bond strength- boiling or melting points? Why?
Boiling points
All the metallic bonds are broken when boiling rather than remaining with melting
What do you have to mention when comparing lattices?
The composition of the lattice- i.e ions, atoms or between molecules
