Periodicty Flashcards
What happens to atomic radius across a period?
It decreases.
This is due to the increased number of protons in the nucleus while the outer electrons remain in the same shell.
Why does atomic radius decrease across a period?
Because there are more protons in the nucleus and the same amount of shielding.
This results in a stronger attraction between the nucleus and outer shell electrons.
What effect does the stronger attraction have on outer shell electrons?
Outer shell electrons are pulled closer to the nucleus.
List the elements in order from sodium to chlorine.
Na, Mg, Al, Si, S, Cl.
What happens to atomic radius down a group?
It increases.
This is due to more shells/energy levels.
Why does atomic radius increase down a group?
There is a greater distance between the nucleus and outer electron.
This results from increased shells.
What effect does increased shielding have on atomic radius?
It leads to a weaker attraction between the nucleus and outer shell electrons.
This causes outer shell electrons to be further away.
Which elements are examples of increasing atomic radius down a group?
Sodium (Na), Rubidium (Rb), Cesium (Cs).
These elements illustrate the trend in atomic radius.
What is electronegativity?
The power of an atom to attract the 2 electrons in a covalent bond.
How does electronegativity change across a period?
Electronegativity increases across a period.
What causes the increase in electronegativity across a period?
More protons in the nucleus and smaller atomic radius lead to stronger attraction between the nucleus and the 2 electrons in a covalent bond.
List elements in order of increasing electronegativity: Na, Mg, Al, P, S, Cl.
Na < Mg < Al < P < S < Cl
What is electronegativity?
Power of an atom to attract the 2 electrons in a covalent bond.
How does electronegativity change down a group?
Electronegativity decreases.
What are the reasons for the decrease in electronegativity down a group?
More shells / more shielding and larger atomic radius.
What is the trend of electronegativity in Group 1 elements?
Electronegativity values are lower as you move down the group: Li (4), Na (2), K (1), Rb (1), Cs (1).
What is the trend of electronegativity in Group 7 elements?
Electronegativity values are lower as you move down the group: F (4), C (3), Br (2), I (1), At (1).
What is the general trend of 1st ionisation energy across a period?
Increases across a period due to more protons and smaller atomic size.
Atoms get smaller, leading to a stronger attraction from the nucleus to the outer shell electron.
What happens to the number of protons across a period?
The number of protons increases as you move across a period.
This increase contributes to the overall increase in ionisation energy.
What remains the same across a period regarding electron shielding?
The electron shielding remains the same across a period.
Atoms have the same shell configuration.
What effect does atomic size have on ionisation energy across a period?
Atomic size decreases across a period.
Smaller atoms have a stronger attraction from the nucleus to the outer shell electron.
What is 1st ionisation energy?
The energy required to remove the outermost electron from a neutral atom in its gaseous state.
What is the trend in 1st ionisation energy for Na, Mg, Al, Si, P, S, Cl, Ar?
The 1st ionisation energy increases from Na to Ar.
From which orbital is the electron lost in Group 2 elements?
From the s orbital.
From which orbital is the electron lost in Group 3 elements?
From the p orbital.
Which orbital is higher in energy, s or p?
The p orbital is higher in energy than the s orbital.
What is the relationship between the energy levels of 4s and 3d orbitals?
4s is of lower energy than 3d.
What is the relationship between the energy levels of 2s and 2p orbitals?
2s is of lower energy than 2p.
What does distance from the nucleus affect?
It affects the energy levels of the orbitals.
What happens to a Group 5 element during ionization?
A Group 5 element loses an electron from the orbital with 1 electron (p3).
What happens to a Group 6 element during ionization?
A Group 6 element loses an electron from the orbital with 2 electrons (p).
Why is it easier for a Group 6 element to lose an electron compared to a Group 5 element?
Extra electron-electron repulsions make it easier to lose an electron from p4 than p3.
Explain why metallic bonding has a high metling and boiling points
Strong attraction between metal ions and delocalised electrons
Why is Aluminium > Mg> Na
Higher charge
More delocalised electrons
Smaller ions
(Stronger metallic bonding)
Why is Sulphur have a high melting and boiling points
It is a simple molecular structure and goes around in 8.
It has 128 electrons so stronger vdw forces so harder to overcome
Why are giant covalent structures hard to break?
Have to break many strong covalent bonds
Why are simple moleculars easy to break.
- weak van der waals forces between molecules
- bigger molecules = more elextrons = more vdw = higher mp and bp
Why are monotomic elements very easy to break?
Very weak van der waals forces between atoms
