Module 3.1 Flashcards
How is the periodic table arranged
In groups and periods
In order of increasing atomic number
Metals are on the left and non -metals are on the right
What is a group (periodic table)
A vertical column in the periodic table
Each element in the same group has the same number of electrons in their outer shell
Elements in the same group have similar properties
What are periods (periodic tables)
A horizontal row in the periodic table
There are general trends in the periods
What is periodicity
The trend in properties that is repeated across each period
It can be used to predict the properties of elements
What do all elements in a group have
The same type of orbitals
The same number of electrons in the outer shell
How do you use the periodic table to tell the final subshell of an element
Depending on the ‘block’ of the periodic table
S block = s orbital
D block = d orbital
P block = p orbital
F block = f orbittal
How does atomic radius affect nuclear attraction
The larger the atomic radius, the weaker the nuclear attraction
How does nuclear charge effect nuclear attraction
The greater the charge, the greater the attraction
What is electron shielding
Electrons in inner shells repel outer shell electrons
The more inner shells the larger the shielding effect
How does electron shielding effect nuclear attraction
The larger the shielding effect, the smaller the nuclear attraction is on the outside ekectrons
How does nuclear attraction affect ionisation energy
(Forming positive ions)
The smaller the nuclear attraction, the smaller ionisation energy required
Electrons in the outer shell are removed first as they experience the lowest nuclear attraction
They require the lowest amount of ionisation energy to remove
What is ionisation energy
The energy needed to form POSITIVE ions
It is a measure of how easy it is to lose electrons
What is ionisation
The process where atoms gain or lose electrons to form an ion
Using ionisation energies how can you tell when there is a new shell
Electrons in the same shell require a similar amount of energy to remove
When the electrons between two shells are removed there is a large spike in ionisation energy
How does ionisation energy generally change across periods
Generally ionisation energy increases moving towards the non metals (as nuclear attraction is greater)
Why is the ionisation energy in group 3 lower than group 2
E.g Boron
Beryllium has a full stable orbital
Boron onky has 1 electron in the p orbital, which is easier to remove
(Easier to remove 1 unstable electron to obtain a stable orbital)
Why is the ionisation energy in group 6 lower than group 5
E.g Oxygen
Nitrogen has 3 electrons in the p orbital, it is half full and therefore stable
Oxygen has 2 electrons in the Px orbital
These electrons repel each other, so it is slightly easier to remove one of the Px electrons
How does atomic radius change across the period
It decreases, as there are more protons in the nucleus, so nuclear attraction is stronger, and the shells are pulled inwards
How does ionisation energy change moving down groups
Ionisation energy decreases as:
The number of shells increases (weaker force on outer electrons)
The shielding effect is greater
How does nuclear attraction change going down a group
There are more electron shells so it decreases
(Greater shielding effect)
What is the basic giant metallic structure
A lattice of cations surrounded by a sea of delocalised electrons
Why do metals have a high melting and boiling point
Electrons are free to move but the cations are stationary
The attraction between the cations and delocalised electrons is strong
A high temperature is needed to overcome the metallic bonds
Why can metals conduct electricity
The delocalised electrons can move freely within the lattice
Why are metals ductile and malleable
The sea of electrons allows the layers of cations to slide over each other
How do melting points change moving across the period (up to group 4)
The melting point increases steadily
This is because the elements have a giant structure
-the nuclear charge increases (electrons in the atom increase)
- There is a stronger attraction
How do melting points change after group 4 (going across the periods)
It decreases sharply between group 4 and 5 as the elements go from giant covalent structures, to simple molecular structures
From group 5+ the melting points remain low, as the elements have simple molecular structures
What are the physical properties of group 2 elements
High melting + boiling points
Light metals with low density
Form colourless compounds
How does reactivity change going down group 2
Reactivity increases
Going down the group nuclear attraction decreases, so the outer electrons require less ionisation energy to lose
How do group 2 elements react with water
They all lose electrons and form a hydroxide
Moving down the group each metal reacts more vigorusly with water
How do group 2 elements react with acids
They form salt and hydrogen gas
The reaction becomes more vigorous moving down the group
How does the solubility of group 2 hydroxides change going down the groupb
(Due to a lower lattice enthalpy)
Solubility increases going down the group (creating more alkaline solutions)
What group is the halogens
Group 7
How does reactivity change going down the halogens
Reactivity decreases going down the group
The number of electron shells increases so the shielding effect is higher - nuclear attraction is weaker
Going down the group it is harder for them to attract an additional electron
What change would be observed if chlorine displaces bromine
The solution would become the colour of bromine
It would turn orange / brown
What is the test for sulphate ions (So4 2-)
Add dilute HCL and barium chloride to the sulphate
A white precipitate (barium sulphate) should be produced
How is carbonate (CO3 2-) tested for
Add a dilute strong acid to the carbonate
Effervescence / colourless gas is produced
Pass the gas through limewater
The gas turns the limewater cloudy
How are halide ions tested for
Add nitric acid to the halid solution (to react with other elements)
Next add silver nitrate to the solution, to form a silver halide (which is insoluble)
Why is silver nitrate used to test for halide ions
As silver halides are insoluble, so the colour of the precipitate can be used to identify the halide
What precipitates do chlorine, bromine and iodine make after reacting with silver nitrate
Chlorine - white (silver chloride)
Bromine - cream (silver bromide)
Iodine - yellow (silver bromide)
If the precipitate colour when testing for halide ions is unclear, what should you do
Add nitric acid
Silver chloride + dilute nitric acid = colourless solution
Silver bromide + concentrated nitric acid = colourless solution
Silver Iodide + ANY nitric acid = no change
What colour / state is fluorine naturally
A pale yellow / green gas
What colour / state is chlorine naturally
A yellow / green gas
What colour / state is bromine naturally
A brown liquid
What colour / state is iodine naturally
A dark grey solid
How does the stability of group 2 carbonates change going down the group
Going down the group, the carbonates become more stable
This is because the charge density decreases (as atomic radi increases), so the carbonate ion is less distorted (its electrons are not attracted strongly)
