Periodicity Flashcards
Define periodicty
Periodicity is the repeating pattern of physical or chemical properties going across the periods
Features of atomic radii
-Atomic radii decrease from left to right across a period (the right would be the smallest)
-Because the increased number of protons create more positive charge attraction for electrons which are in the same shell with similar shielding.
-Atomic radius increases down groups due to extra electron shells being added
What is the trend with the first ionisation energy
There is a general trend across to increase.
- This is due to the increasing number of protons as the electrons are being
added to the same shell
-The atomic radius increases as we go down the group. Outer electrons further from the nucleus. Attractive force is weaker.
-Energy required to remove an electron decreases.
Shielding increases as we go down the group. More shells between nucleus and outer shell. Attractive force is weaker. Energy required to
remove an electron decreases.
This data provides strong evidence for shells in atoms and proves Niels Bohr’s model of the atom is correct. BUT it didn’t explain data
shown going across a period.
Why is there a small drop between Mg and Al in ionisation energy
Mg has its outer electrons in the 3s sub shell, whereas Al is starting to fill the 3p subshell.
-Al’s electron is slightly easier to remove because the 3p electrons are higher in energy.
Why is there a small drop between phosphorous and sulfur for ionisation energy
- Sulfur’s outer electron is being paired up with an another electron in the same 3p orbital.
- When the second electron is added to an orbital there is a slight repulsion between the two negatively charged
electrons - Which makes the second electron easier to remove.
Mp and Bp
- For Na, Mg, Al- Metallic bonding : strong bonding – gets stronger the more electrons there are in the outer shell that are released to the sea of electrons.
- A smaller sized ion with a greater positive charge also makes the bonding stronger.
-Higher energy is needed to break bonds.
Why does Si have a high Bp and Mp
Si is Macromolecular: many strong covalent bonds between atoms
-High energy needed to break covalent bonds– very high mp +bp
Why is it easy to break the Cl2, S8, and P4 bonds
they are simple molecular : weak van der waals between molecules, so little energy is needed to break them – low mp+ bp
Why does S8 have a higher mp than P4
because it has more electrons
(S8 =128)(P4=60) so has stronger v der w between molecules
How are the elements ordered in the periodic table
By proton number -thesmaller number on the element-
What are the s block elements
Group 1 and group 2 elements + hydrogen and helium
What are p block elements
Group 3-7 elements
What are D block elements
Transition metals
Why does Ar have the lowest Mp
It is monoatomic and so has weaker vdw forces between atoms
Define ionisation energy
The minimum amount of energy required to remobve 1 mole of electrons from1 mole of atoms in the gaseous state
Give a half equation of Na getting its first ionisation energy removed
Na(g)–>Na+(g) +e-
What are the factors affecting ionisation energy
Shielding
The more electrons shells between the positive nucleus and negative electron that is being removed the less energy is required. There is a weaker attraction.
Nuclear Charge
The more protons in the nucleus the bigger the attraction between nucleus and outer electrons. This means more energy required to remove the electron.
Atomic Size
The bigger the atom the further away the outer electrons are from the nucleus. The attractive force between nucleus and outer electrons reduces - easier to remove electrons.
What is successive ionisation energy and give an example using Mg
The removal of more than 1 electron from the same atom
Mg+(g)–>Mg2+(g) +e-
What is the general trend for the energy required for successive ionisation energy
There is a generalincrease in energy as removing an electron from an increasingly more positive ion
What is one limitation to Niels Bohrs atomic model
It didnt explain subshells
Define electronegativitiy and features of electronegtaivity
The ability of an atom to attract a shaired pair of electrons in a covalent bond
Electronegativity increases across the period
Decreases down the group
Nuclear charge increases and atomic radius radius decreases therefore bondig pair closer to nucleus
Electron pair in covalent bond more attrcfated to atoms nucleus
