Topic 3&13 Periodicty (3.1, 3.2, 13.1, 13.2) Flashcards
3.1 Elements in the same group tend to have similar chemical/ physical properties, true or false?
True (This is because their atoms have the same number of electrons in the highest occupied energy level/ the same number of electrons in the outermost shell)
3.1 Period tells you the number of valence electrons, true or false?
False (Period=energy level outer electrons occupy, Group=number of outer electrons)
3.2 Define atomic radius.
Distance between nucleus and outer electrons of an atom/ Half the distance between the nuclei of two bonded atoms of the same element
3.2 What is the trend for atomic radius and why?
Decreases across the period because there are more protons, attracting/ pulling the energy levels closer. Increase down the group because number of occupied electron shells increases.
3.2 Define ionic radius.
Distance between nucleus and outer electrons of an ions
3.2 What is the trend for ionic radius and why?
Cations are smaller than the parent atoms as the number of electron shells decreased by one.
Anions have more electrons so they are bigger than the parent atoms due to electron repulsion in the full outer shell.
All cations in a period have the same number of electrons, as is the case with anions. But protons still increase so the ionic radius decreases.
3.2 Define ionisation energy.
Energy required to remove an electron from a neutral gaseous atom. Note: pay attention if question asks for FIRST ionisation energy
3.2 What is the trend for ionisation energy and why?
Generally increase across a period because higher concentration of protons increases the effective nuclear charge, and electrons are removed from the same energy level.
Decrease down the group because increased distance between nucleus and outer electrons reduces the electrostatic attraction.
3.2 First and second ionisation energy equation for element X
X(g) → X+(g) + e-
X-(g) → X2+(g) + e-
Energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+
3.2 Define electron affinity.
The energy released when an electron is added to a neutral gaseous atom to form a negative ion.
3.2 What is the trend for electron affinity and why?
Increases across the period because the nuclear charge increases while the shielding and distance remains roughly the same because electrons are added to the same energy level. Aside from high nuclear charge, non-metals also have incomplete outer energy levels so attract electrons the most.
Decreases down the group because it takes less energy to gain an electron to become and anion.
3.2 Define electronegativity.
Relative measure of the ability an atom has to attract a shared pair of electrons.
3.2 What is the trend for electronegativity and why?
Increases across the period because of increase in nuclear charge, resulting in more attraction between nucleus and bonding electrons. Also elements with many valence electrons prefers to gain electron to complete their valence shells.
Decreases across the period as the distance increases between nucleus and bonding electrons, increasing the screening so there is less attraction.
3.2 What is the trend for melting points in the alkali metals (group 1) and why?
It decreases as you go down the group because the attractive forces between delocalised outer electrons and positive ions decrease when the atomic radius increases.
3.2 What is the trend for melting points in halogens (group 17) and why?
It increases when you go down the group because the van der Waals’ forces increase with the number of electrons in the molecule.
