cap 2 Flashcards
Define
First ionisation energy
The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of ions.
down a group: first ionisation energy trend
and why
Decreases. Because core charge stays constant and number of shells increases as you move down a group. So, the valence electrons are less attracted to the nucleus as they are further from it. So the energy required to overcome the attraction between the nucleus and valence electron is less.
left to right across a period: first ionisation energy trend
and why
Increases. Because core charge increases and the number of occupied shells (shielding) remains constant as you move across a period. So the valence electrons becomne more strongly attracted to the nucleus, and more energy is required to remove an electron.
Define
Atomic radius
The distance from the nucleus to the valence-shell electrons. (Half the distance between two identically bonded atoms).
Down a group: atomic radius trend
And why
Increases because core charge stays constant and the number of shells increases as you move down a group.
Across a period: atomic radius trend
And why
Decreases as there are more electrons, but the number of fully occupied electron shells stays the same, so the shielding effect does not change. This means that there is a greater electrostatic attraction between the electrons and protons due to the larger number of protons, causing the electrons to be pulled closer to the nucleus, decreasing the atomic radius.
Define
Valency
The number of electrons gained or lost to form a full valence (outermost) electron shell. A measure of the atom’s bonding capacity,
Down a group: valency trend
Stays the same.
Across a period: valency trend
?!??!?!?!?!?!!?!?!?!?!?!
Define
Electronegativity
The ability of an atom to attract electrons in a covalent bond towards itself. (I.e., the ability of an atom to form covalent bonds). The strength of this attraction (same with all ES attraction) depends on the magnitude of the charges involved and the distance between them.
Down a period: electronegativity trend
And why
Decreases. Because core charge stays constant and the number of shells increases down a group. Therefore, valence electrons are less strongly attracted to the nucleus as they are further from it.
Across a group: electronegativity trend
Increases. The number of occupied shells int he atoms remains constant but the core charge increases across a period. Therefore, the valence electrons become more strongly attracted to the nucleus.
Define
Core charge/effective nuclear charge
A measure of the attractive force felt by the valence electrons towards the nucleus. The resultant attractive force experienced by valence electrons once the impact of the shielding effect provided by electrons in inner shells is taken into account.
Subtract the total number of inner-shell electrons from the number of protons in the nucleus.
Down a group: core charge trend
No change because the valence electrons are more weakly held moving down a group as, although the core charge stays constant, the valence electrons are further from the nucleus (there are more shells in the atom) and thus experience a weaker electrostatic attraction.
Across a period: core charge trend
Increases as the valence electrons are more strongly held moving across a period as, although the distance between valence electrons and the nucleus remains essentially constant (same electron shell), the valence electrons are subject to a higher core charge and thus experience a stronger electrostatic attraction.
