Topic 3.2 - Periodic Trends Flashcards
Define ‘periodicity’
recurring trends observed in the properties of elements as they are arranged in the periodic table
What is nuclear charge
atomic number - positive charge (protons) present in the nucleus of an atom
What is effective charge
- net positive charge pulling electrons towards the nucleus
- the stronger the pull on the valence electrons towards the nucleus, the higher the effective nuclear charge
What is electron shielding
when the electrons in the inner energy levels of an atom shield the outermost electrons from the full effect of the positive charge of the nucleus
What is atomic radius
the distance from the nucleus to the outermost electrons (highest energy level occupied)
What is atomic radius
the distance from the nucleus to the outermost electrons (highest energy level occupied)
What is the trend in ‘atomic radius’
-increase down a group
- decrease across a period
Why does atomic radius decrease across a period
nuclear charge increases which means that there is an increase in the number of protons and electrons. Therefore, there is a stronger attraction as electrons are pulled closer to the nucleus, reducing atomic radius
Why does atomic radius increase down a group
- nuclear charge increases = more electron energy levels occupied + more electrons
- however, there is a decrease in the effective nuclear charge experienced by the outermost electrons due to electron shielding
- negative charged electrons repel and move further away from the nucleus = increase in atomic radius
What is ionic radius
size of an ion
What is the trend in ‘ionic radius’
- increases across a period
- increases down a group
Why does ionic radius increase across a period
ions to the left of the pt (cations) lose electrons and become smaller than the original atom
ions to the right of the pt (anions) gain electrons and become bigger than the original atom
Why does ionic radius increase down a group
the number of occupied electron energy levels increases, therefore the ion becomes bigger
What is ‘ionization energy’
The first ionization energy of an element is the energy required to remove one mole of electrons from one mole of gaseous atoms in their ground state
What is the trend in ionization energy
- increase across a period
- decrease down a group
Why does ionization energy increase across a period
more energy required to to strip electrons from the outer shell moving across a period as effective nuclear charge increases and there is an increase in attraction. Elements on the right-hand side also almost have a full shell.
Why does ionization energy decrease down a group
electrons occupy more energy levels and are further away from the nucleus. There is also electron shielding which reduces the attraction between electrons and the nucleus. Therefore less energy is required to strip and electron
What is ‘electron affinity’
the energy change that occurs when an atom gains an electron to form a negative ion
*amount of energy released or absorbed when an atom gains an electron
What is the trend in electron affinity
- increases across a period
- decreases down a group
*same as ionization energy
Why does electron affinity increase across a period
elements on the right-hand side of the pt. have a higher effective nuclear charge and attract electrons the most as they are negative
Why does electron affinity decrease down a group
effective nuclear charge decreases as there is electron shielding and therefore don’t attract electrons
What is ‘electronegativity’
ability of an atom to attract electrons in a covalent bond
What is the trend in electronegativity
- increase across a period
- decrease down a group
Why does electronegativity increase across a period
increase in effective nuclear charge from left to right which means more pulling power and a stronger attraction between the nucleus and the bond electrons
Why does electronegativity decrease down a group
the bonding electrons are furthest from the nucleus which leads to electron shielding and reduced attraction
What is the difference between the trends in atomic radius and ionic radius
ionic radius = increases across a group (loses electrons to form ions)
atomic radius = decreases across a group (nuclear charge increases)
Why is there a drop in ionization energy between elements in group 15-16
the electron removed from group 16 is taken from a doubly occupied p orbital which is easier to remove as it is repelled by its partner
Describe Group 18: noble gases
- very unreactive as they are stable and have a stable octet
- monatomic = only exist as single atoms
- highest ionization energy (hardest to strip electron)
Why are other elements reactive
- unstable and incomplete electron energy levels and will lose and gain electrons to achieve the electron configuration of their nearest noble gas
Describe Group 1: alkali metals
- soft shiny metals
- far too reactive to be found in nature due to low ionization energy which means electrons are lost easily
describe the reaction of alkali metals with water
highly reactive with water = produce alkaline solutions = hydroxide ion formed
describe the reactivity trends of alkali metals
increases down a group because ionization energy decreases as it is easier to lose electrons due to a decrease in effective nuclear charge
*form cations easier
Describe Group 17: halogens
- diatomic molecules
- very reactive non-metals due to their readiness to accept electrons due to the increased effective nuclear charge which exerts a stronger fulling force on electrons
describe the reactivity trends of halogens
- decreases down a group because atomic radius increases and the attraction for outer electrons decreases due to a decrease in effective nuclear charge due to electron shielding
describe the reaction of halogens and alkali metals
metal + halogen = ionic halide
- halogen atom gains one electron from the group 1 metal resulting in both having the stable octet of a noble gas
- forms an ionic compound as there is an ionic bond between the two oppositely charged ions, due to an electrostatic attraction
e.g. NaCl, table salt
What happens when two halogens react
the more reactive halogen displaces the ions of the less reactive halogen from its compounds
halide + silver =
produces a precipitate
metal + oxygen =
metal oxide
metal + water =
metal hydroxide + hydrogen gas
e.g. 2K(s) + 2H20(l) -> 2KOH(aq) + H2(g)
oxides of metals have what structure
giant ionic structure + high electrical conductivity
oxides of non-metals have what structure
molecular covalent + no electrical conductivity
oxides of metalloids have what structure
giant covalent structure + low conductivity
which oxides have the most iconic character
down a group because they have low electronegativities as they have a decreased effective nuclear charge due to electron shielding and electrons are further away as they occupy more electron energy levels. Therefore the difference between the electronegativity of these elements and oxygen is low
metal + acid =
salt + hydrogen
metal oxide + acid
salt + water (neutralisation)
Why are oxides of metals basic
they are ionic and basic because metal oxide reacts with acids and form salt + water. They neutralise the acids so they are basic
neutralisation formula + examples
acid + base = salt + water
NaOH + HCl -> NaCl + H20
why are oxides of non-metals acidic
reacts with bases to form salt and water
why is aluminium oxide amphoteric
- considered to be an ionic oxide with some covalent character
- shows amphoteric properties reacting with both acids and bases
metal oxide (base) + water =
basic solution (alkaline)
what charactersises a basic and acid solution
base = OH (hydroxide)
acid = H (hydrogen)
non-metal oxide (acid) + water
acidic solution
- sulfur trioxide + water = sulfuric acid
- phosphorus pentoxide + water = phosphoric acid
what are characteristics of an amphoteric oxide
it can react with both acids and bases to give a salt
what are the trends in melting point
ionic/covalent giant structures = high mpt
molecular covalent structure = low mpt
