Chapter 2: Chemical Periodicity Flashcards
explain the terms ‘effective nuclear charge’ and ‘shielding effect’
nuclear charge describes the electrostatic forces of attraction between protons and electrons.
[Electrons are simultaneously attracted to the nucleus and repelled by one another.]
shielding effect (aka electron-electron repulsion or screening effect) describes a context where an electron is partially shielded from the positive charge of the nucleus by the other/all electrons.
[core electrons (surrounding the nucleus) shield the valence electrons (outermost e-) from the full nuclear charge MOST EFFECTIVELY as it is directly between the valence e- and nucleus. hence: it is the CORE e- that provide substantial shielding]
hence, the RESULTANT/net nuclear charge acting on the VALENCE electron is termed the ENC, effective nuclear charge.
state how to calculate for ENC
ENC (Zeffective)
Z = nuclear charge (no. of protons)
o = shielding constant (no. of core e-s)
Zeffective = Z - o
explain how to relate ENC to the PT
electrons are added to the valence shell going across a period
while the proton number increases along a period, core electrons stay the same hence ENC increases across a period
define atomic radius/how its measured
the distance between the nucleus of an atom and its valence shell (in picometers, 10*^-12m)
atomic radius is determined by: ENC for valence e-, number of shells the e- occupy
state the atomic radii trends
atomic radii/size increases going down a group
atomic radii/size decreases across a period
explain atomic radii trends going across any period
across a period (L to R), nuclear charge increases due to increase in no. of protons but the shielding effect remains similar/insignificant (because the no. of core e- remain the same while valence e- increases)
therefore, ENC increases while no. of shells remain the same
as ENC increases, valence electrons feel increasingly attracted to the nucleus. they are drawn closer to the nucleus, hence atomic size is reduced going across any period
explain atomic radii trends going down any group
nuclear charge increases due to the increase in number of protons, as well as the shielding effect increases due to to increase in core electrons. hence ENC remains similar while the no. of shells increase.
the no. of shells increase so the outermost shell lies farther from the nucleus, hence the DISTANCE from the nucleus and the valence shell is greater (atomic size is increased going down any period)
define the term ionisation energy
the minimum amount of energy required to remove a valence electron from a gaseous atom or ion forming a gaseous cation of a higher oxidation number
to remove an e- from an atom, the attractive force between e- and nucleus must be overcome
the greater the attraction between e- and nucleus, the more difficult it is to remove the e- (hence the greater the IE is)
explain ionisation energy
the valence e- is the least firmly attached (least ENC, experiences the least amount of forces of attraction) so it is the first to go
[1st ionisation energy refers to the amount of energy needed to remove the 1st electron from gaseous atom.
ie. M (g) + energy → M+ (g) + e−]
[2nd ionisation energy refers to the amount of energy needed to remove the 2nd electron (NOT remove 2 e-!!!) from gaseous ion.
ie. M+ (g) + energy → M2+(g) + e−]
removing an e- is an endothermic process (where the gaseous atom requires energy)
explain why it takes much more energy to remove core e- than valence e-
core electrons are closer to the nucleus and experience greater NC due to fewer filled shells shielding them from the nucleus
state ionisation energy trends (for main group elements)
1st IE increases going across a period
1st IE decreases going down a group
explain 1st IE increasing across a period (general trend)
recall: no. of shells with e- remains constant while ENC increases, hence atomic size decreases
the larger the ENC (furthermore, smaller atomic size), the more tightly valence e- are held
hence the energy required to remove the first e- (1st IE) increases
HOWEVER THERE ARE EXCEPTIONS
explain 1st IE decreasing going down a group
recall: ENC remains similar while the no. of shells with e- increases going down a group hence atomic size increases
this greater separation between the electron and nucleus, means a weaker attraction between them
hence the energy required to remove the first e- (aka 1st IE) decreases OR is easier to remove an e-
state the exceptions to the increasing 1st IE trend across a period
Be & B, N & O
Be’s 1st IE is greater than B’s, and N’s 1st IE is greater than O’s
explain the exceptions to the increasing 1st IE trend across a period
Be: [He] 2s2, B: [He] 2s22p1
2p subshell is higher in energy level and further away from the nucleus compared to 2s subshell. hence B’s 2p e- is easier to remove than Be’s 2s e-
N: [He] 2s22p3, O: [He] 2s22p4
N’s 2p subshell contains all unpaired e-, while O’s 2p subshell has one pair of e-. within the paired e-, the electron-electron repulsion allows for the removal of the 1st e- to be done easier
