Periodicity (7) Flashcards
What happens to the atomic number in the periodic table?
It increases as you go left to right
Groups
Vertical columns of the periodic table, tells you how many electrons on your outermost shell
Periods
Horizontal columns that tell you how many shells the element has
Periodicity
Trends in the periodic table such as electron config, ionisation energy, structure and melting point
S Block
Group 1+2
P block
Groups 3-8
D Block
Transition metals
F Block
Bottom two rows
Ionisation Energy
The energy required for an atom to lose an electron in order to become a positive ion
First Ionisation Energy
The energy required to release 1 electron from one mole of gaseous atoms of an element to form 1 mole of gaseous cations.
How do electrons sit on shells?
The electrons sit on shells as they are attracted to the nucleus, but the higher the shell number, the further away the electron will be from the nucleus so the weaker the bond will be.
Factors Affecting Ionisation Energy
Size of atomic radius, Nuclear charge, Electron shielding
Atomic Radius
The bigger the radius of the atom, the bigger the distance between the outermost electron and the nucleus which will weaken the bond so the electron can be lost easier.
Nuclear Charge
The more protons that exist in the nucleus, the greater the attraction to the electron as protons are positive and attract the negative electrons
Electron Shielding
The more shells that are between the nucleus and outermost electron, the more electron shielding there is and the more the innermost shells will repel the outermost electron so the bond between the nucleus and electron will be weaker.
How many ionisation energies does an atom have?
It will have as many ionisation energies as there are electrons on its outer shell e.g Helium will have 2.
Second Ionisation Energy
The energy needed to remove an electron from 1 mole of gaseous +1 ions of an element to form 1 mole of gaseous 2+ ions
Is the second ionisation energy greater than the first?
Yes it is because in helium for example, after the first electron is removed, the remaining electron on the outer shell is pulled closer towards the nucleus which means more energy is required to remove this electron.
How can you make predictions about an element?
Look at the table of ionisation energies and see how the energy is increasing. If it’s a steady slow increase, it means the electrons being removed are from the same shell. If there is a sharp jump, this indicates it’s moved to a lower shell.
How does the ionisation energy trend in a period?
There is an increase of ionisation energies in a period, and a peak will be a noble gas as they are really hard to take electrons from. A sharp decrease in ionisation energy means you’ve started the next period.
How does the ionisation energy trend in a group?
As you go down a group, the ionisation energy will decrease.
Why is the trend like that in a period?
It will increase because the nuclear charge is increasing as there are more protons. The atomic radii also decrease which makes the nucleus have a stronger bond with the electron. Electron shielding is the same as they have the same number of shells
Why is the trend like that in a group?
It decreases within a group because the atomic radii increase, the nuclear charge and the electron shielding all increase because the atoms are getting bigger.
Why does the energy level within the periods sometimes fall?
There are some drops when moving on to fill the next orbital e.g from 2s2 to 2p6
Comparing beryllium and boron
Boron follows beryllium and is the start of the p 2p orbital so it has one electron in that orbital which is easier to remove than the two close electrons in beryllium
Comparing oxygen and nitrogen
Oxygen follows nitrogen and is when you start filling in the opposite spin electrons for the 2p orbital so naturally, because the electrons repel each other, it is easier to remove the electron from the oxygen rather than the nitrogen that has a single electron in each orbital.
Metallic Bonds
Strong electrostatic attraction between positive metal ions and negative delocalised electron floating within its structure
Structure of metal
Metal atoms will donate their outer shell electrons in order to become stable metal cations and form a sea of delocalised electrons that can float around the structure. They are held together in a giant metallic lattice
Properties of metal
High melting point because of strong bonds in it.
Good conductor because of the delocalised electrons
Malleable and ductile because the positive ions can be dispositioned without the lattice collapsing
It is not soluble but could react with polar molecules
Giant covalent structures
Diamond is a giant covalent lattice and uses all 4 of its covalent bonds to make a tetrahedral structure with 109.5 bond angle.
Graphene and graphite are also giant covalent structures but the carbons bond to form planar hexagonal layers.
Properties of giant covalent structures
High melting and boiling points because of the strong covalent bonds
They are insoluble
Diamond doesn’t conduct electricity but graphite and graphene can as they only use three of the four covalent bonds
Trend of melting points
The melting points increase from group 1 to 3 where all the metals are and then there is a sharp decrease where the non-metals are
