1 - Atomic Structure and the Periodic Table Flashcards
Basic things to memorise and harder exam questions, however the bulk of the concepts are not in flashcard form - only notes.
Define relative isotopic mass:
the mass of an atom of a particular isotope of an element relative to 1/12 the mass of a C-12 atom
Define relative atomic mass and give its symbol:
Ar - the weighted average mass of all the isotopes of a particular element, relative to 1/12 the mass of a C-12 atom
What is the difference between Mr and relative molecular mass?
-Mr is the sum of all the Ar’s of all the atoms in a compound
-relative molecular mass is only used for molecular substances, whilst Mr is only used for compounds with giant structures
How would you predict the mass spectra of a molecular substance?
-write out all the possible variants of the molecule using its isotopes
-find the probability of getting each of them using the isotopes’ abundances and a probability tree
How could you find the relative abundances of two isotopes?
-compare the intensity of the signals received
-in a mass spectrometer
Name 3 factors that affect ionisation energy:
-nuclear charge
-distance from nucleus
-shielding
(NDS)
Always try to mention each of these three factors in exams, and why certain factors have changed, increased, decreased, or how they outweigh other factors (better and more detailed than saying something about having a full outer shell)
How does the first IE change down a group and why?
-decreases, because:
-outer electron gets further from nucleus, less attractive forces on it
-increases shielding by electrons from inner shells
-nuclear charge increases too, but this is outweighed by the other 2 factors
What did the first IE graphs across a period show about electron configuration?
the existence of electron sub-shells
What did successive ionisation energies of a particular element show about electron configuration?
the existence of electrons residing in different energy levels
How many electrons can the first 4 shells hold?
2, 8, 18, 32
formula - 2n²
What is periodicity?
a repeating pattern of chemical and physical properties that occur across each period
Why is aluminium’s 1st IE lower than magnesium’s?
-electron removed from Al is in a higher energy level (3p rather than 3s), meaning it requires less energy to remove
-even tho Al has a higher nuclear charge than Mg, the effect of the electron being in a higher energy level outweighs the added charge
Why is silicon’s 1st IE higher than aluminium’s?
-Si has a higher nuclear charge than Al
-the electrons being removed in both elements are from the 3p subshell, meaning they are both equally shielded
-the electron in Si will be closer to the nucleus, and therefore requiring more energy to remove
Why is a potassium ion smaller than a potassium atom?
-one fewer subshell (no 4s in the ion)
-the same number protons will be holding 1 less electron, so they will be held more closely
-decreases ionic radius
Give some reasons for and against having hydrogen in Group 1:
-only has 1 outer electron in s-subshell
-hydrogen is not an alkali metal
-different chemical properties to the rest of Group 1
-can form a negative hydride ion (H⁻)
