1. Atomic structure ( ORBITALS + IONISATION) Flashcards
Define an orbital
a region of space in an atom containing up to two electrons with opposite spins
Give the 4 orbitals in an atom
s p d f
What is the shape of the s orbital compared to the p orbital
s orbital is spherical and p orbital is dumbell shaped
What are the names of the 3 P orbitals
Px, Py and Pz
How many electrons can the p orbital hold
6
Give the maximum number of electrons in each shell
shell 1 - 2
shell 2 - 8
shell 3 - 18
shell 4 - 32
Give the order in which electrons fill subshells
1s–> 2s –> 2p –> 3s –>3p –> 4s –> 3d –> 4p –> 5s –> 4d –> 5p
Give the type of orbital found at each energy level
1 - s
2 - sppp
3- spppddddd
4- spppdddddfffffff
Why is the 4s orbital filled before the 3d orbital
Because 4s is slightly lower in energy
Describe 2 ways of writing electron structure
Spin diagrams - electrons in boxes with arrows
Short hand - uses the noble gas that comes before it.
eg: oxygen is 1s2 2s2 2p4 but can be written as [He] 2s2 sp4. (must use square brackets)
–> This is because helium has 2 electrons in its outer shell so it can substitute the 1s2
When forming ions, is the 4s orbital removed before the 3d orbital
Yes ( apart from copper and chromium )
What is the exception to the rule of filling 4s orbital before 3d orbital
Chromium and copper
Its more stable this way
How many electrons can the d orbital hold
10
(so it could be written as eg: 3d^10)
How many electrons can the f orbital hold
14
In what order do electrons fill subshells
First singularly then pairing up
What is meant by s block elements
The elements in group 1 and 2 have highest energy electron in an s orbital
What is meant by d block elements
Transition metals
What is meant by p block elements
Elements from group 3 to group 7
Define isoelectronic
When either atoms, isotopes or ions ( referred to as species) have the same number of electrons
Define first ionisaton energy
First ionisation energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state
[ M(g) → M+(g) + e- ]
Define second ionisation energy
is the energy required to remove one mole of electrons from one mole of 1+ ions in the gaseous state.
[ M+(g) → M2+(g) + e- ]
Give 3 factors that affect ionisation energy
size (distance of the outer shell electrons from the nucleus) , shielding, nuclear charge
What happens to first ionisation energies going down a group
It decreases
Explain why first ionisation energies decreases down a group
Atomic radius (size) increases - increased number of shells so distance from outer shell electrons to nucleus increases
Shell number increases - more shielding between outer shell electron and nucleus due to the inner electrons, so force of attraction decreases
Nuclear charge increases - more protons in nucleus ( but this increase is outweighed by the other factors)
What happens to ionisation energy going across a period
They generally increase
Explain why ionisation energy increases across a period
Size of the atom decreases - due to the greater force of electrostatic attraction between the positive nucleus and outer shell electron
Nuclear charge increases - due to higher number of protons in the nucleus
Shielding stays the same - number of shells stays the same as electrons are removed from the same shell so experiences simillar sheilding
Explain why sucessive ionisation energies increase in size
The electrons are now closer to the nucleus and there are more protons than electrons. The electron must also overcome the increasing attraction of the positive ion.
What do the big jumps on a ionisation energy graph show + what do they provide evidence for
shows when an electron is being removed from a shell
Provides evidence for shells
Describe how to work out the group of an element and its structure from an ionisation energy graph
To work out the group, look for how many electrons are removed before the first big jump. Look from left to right
To work out structure, look from right to left and count how many points there are before a big jump
Define periodicity
The regularly repeating patern of atomic, physical and chemichal properties with increasing atomic number
How do chemical and physical properties of isotopes differ
the chemical properties are the same (because the number of protons doesn’t change and electron configuration remains the same) but the physical properties (such as mass/density) are different because of different amounts of neutrons
Why is there a drop in ionisation energy between group 2 and 3
An element in group 3 has its outer shell electron in a p orbital, not an s orbital.
P orbital has a slightly higher energy than s orbital
P orbital is therefore found further away from the nucleus.
These 2 factors override the increased nuclear charge so ionisation energy decreases
Why is there a drop in ionisation energy between group 5 and 6
Group 5 elements have their outer shell electrons in singularly half filled p orbitals which are more stable than those with paired filled subshells.
Elements in group 6 have their outer shell electron paired with another electron in the p orbital. The repulsion between the 2 electrons makes it easier to remove so ionisation energy decreases
Note: Ionisation energy for group 6 elements is less than group 5 but greater than group 3
When going across the period does melting and boiling point increase or decrease for metals
Increases as metallic bonds get stronger due to decreased atomic radius - there is a stronger electrostatic attraction between positively charged ions and a sea of
delocalised electrons
Describe the melting and boiling point for giant covalent lattice structures ( Carbon, Silicon)
Very high due to many strong covalent bonds between atoms so high energy needed to break covalent bonds
What is the trend of melting/boiling points across a period?
1) for metals, metallic bonding , melting/boiling points increase across period as there are more electrons that are delocalised, smaller sized ions and ions with a greater charge
2) then a peak (e.g. silicon) as silicon forms a giant covalent structure (macromolecular) many strong covalent bonds
3) the next non-metals are simple molecular with weak London forces however S forms S8 so is slightly higher than P (P4)
4) then decreases to halogens (diatomic) and then lowest at noble gases (monatomic)
Give a reason why log is used to plot ionisation energies on a graph
Numbers produced are smaller and easier to plot on a graph
What does successive ionisation energies provide evidence for
Existence of quantum shells and the group to which the element belongs to
What does first ionisation energy of successive elements provide evidence for
Electron sub - shells
