3.1 Periodic table and energy 7.1-8.3 Flashcards
Atomic radius
The greater the distance between the nucleus and the outer electron the less the nuclear attraction
The force of attraction falls off sharply with increasing distance, so atomic radius has a large effect
Nuclear charge
The more protons there are in the nucleus of an atoms, the greater the attraction between the nucleus and the outer electrons
Electron shielding
Electrons are negatively charged and so inner shell electrons repel outer shell electrons
This repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons
Trend in first ionisation energy down a group
Atomic radius increases
More inner shells so shielding increased
Nuclear attraction on outer electrons decreases
First ionisation energy decreases
General trend in first ionisation energy across period two
Nuclear charge increases Same shell: similar shielding Nuclear attraction increases Atomic radius decreases First ionisation energy increases
Ionisation
Ionisation is the reaction where one electron is removed from an atom.
There will be one electron in the product of the equation.
So we will always have an e− term on the right-hand side of the equation.
Ionisation equation
The first ionisation energy equation of Na:
Na(g) → e− + Na +(g)
The second ionisation energy equation of Mg:
Mg+(g) → e− + Mg 2+(g)
First ionisation energy
The first ionisation energy is the energy required to remove an electron from every atom in a mole of atomic gas, to produce a mole of unipositive gaseous ions.
Second ionisation energy
The second ionisation energy is the energy required to remove an electron from every ion in a mole of unipositive gaseous ions, to produce a mole of dipositive gaseous ions.
Factors effecting ionisation energy
The main factors affecting ionisation energies are the nuclear charge, the distance from the nucleus and electron shielding.
Distance from nucleus- ionisation energy
Electrostatic attraction decreases sharply with distance.
This means that less energy is needed to remove electrons which are further away.
This means that as distance increases, ionisation energy decreases.
In practice, this means that the higher the principal quantum number of an electron, the lower its ionisation energy.
Nuclear charge- ionisation energy
The greater the number of protons in the nucleus, the greater the attraction of the electron to the nucleus.
A greater attraction of the electron means more energy is needed to remove the electron.
This means that ionisation energy is greater.
Shielding- ionisation energy
The greater the number of electrons between the nucleus and the outer electrons, the lower the effective nuclear charge.
This is because the positive charge felt by the electron is reduced by the electrons in between.
This means that the greater the number of electrons, the lower the ionisation energy.
Experimental evidence for shells
Ionisation energies can be used to show the presence of shells in atoms.
Atomic radius and ionisation energy in group two
Atomic radius increases as you go down Group 2.
This is because each extra electron shell is further away.
Ionisation energy decreases as you go down Group 2.
This is because the outer electrons are further away and experience less attraction to the nucleus.
Melting point in group 2
Melting points decrease as you go down Group 2.
This is because the ion cores have larger radii down the group.
The free electrons experience less attraction to the nuclei because of the larger radii.
The bonding is weaker, so the melting point is lower.
Magnesium has an anomalously low melting point.
This is because it has a different crystal structure to the rest of Group 2.
Group two reactions with water
Group 2 metals react with water to form metal hydroxides. For example:
Mg + 2H2O → Mg(OH)2 + H2
Reactivity increases as you go down Group 2. This is because the lower elements have lower ionisation energies.
Beryllium is an exception. If beryllium were to lose two electrons it would be tiny and have a very high charge density. This would make it unstable and so beryllium doesn’t react with water.
Solubility of salts in group two
The solubility of hydroxides increases as you go down Group 2.
Magnesium hydroxide is very insoluble. The phrase to use for this is ‘sparingly soluble’.
The solubility of sulfates decreases as you go down Group 2.
Barium sulfate is completely insoluble.
Trends in reactivity group 2
The reactivity of group 2 metals increases down the group.
The ionisation energies needed decrease down the group because increased shielding from extra shells of electrons and a greater distance from the nucleus makes it easier to remove an outer electron.
Both first and second ionisation energies decrease down the group. The second ionisation energy is greater than the first ionisation energy within an element.
Group 2 reaction with oxygen
Group 2 elements react with oxygen in the air to form metal oxides, For example:
2Mg + O2 → 2MgO
The elements react more readily with oxygen down the group because of the decrease in ionisation energies.
Group 2 reaction with dilute acids
Group 2 metals react with dilute acids to produce a salt and hydrogen. For example:
Ca + H2SO4 → CaSO4 + H2
Reactivity increases down the group because of the decrease in ionisation energies.
The salts produced become less soluble as you go down the group.
Group 2 oxides and water
Group 2 oxides steadily react with water.
The reaction produces a metal hydroxide salt. The salt dissociates into metal ions and hydroxide ions, giving the solution an alkaline pH.
Group 2 hydroxide solubility down the group
Group 2 hydroxides become more soluble down the group.
This means that the solutions become more strongly alkaline the further down Group 2 you go.
Group 2 trends in pH
Because Group 2 oxides get more soluble as you go down Group 2, the pH increases when reacting Group 2 oxides with water as you go down the group.
E.g. Calcium oxide (CaO) will have a higher pH than magnesium oxide (MgO).
