Topic 3: Periodicity Flashcards
atomic radius
the distance from the centre of the nucleus to the outermost shell of electrons
bonding atomic radius
- aka covalent radius
- half of the distance between the centre of the nuclei of 2 covalently-bonded atoms
why is atomic radius (RnB) > covalent radius (Rb)?
- the covalent bond is formed by the overlapping of atomic orbitals
- the overlapping region becomes common ground
periodicity
the tendency of similar properties of the element to recur at intervals, if they are arranged in order of increasing atomic number
periodic trend of atomic radius
- across a period, atomic radius decreases due to increased nuclear charge, which pulls electrons closer to the nucleus
- down a group, atomic radius increases due to increased no. of shells and increased shielding effect
ionic radius
the effective distance from the nucleus of the ion up to which it has an influence in the ionic bond
characteristics of ionic radius
- radius of cation radius of parent atom
electron affinity
- the amount of energy released when a mole of electrons is added to a mole of gaseous atoms
- represents the ability of an atom to hold an additional electron
- the higher the tendency, the higher the e.a.
factors affecting electron affinity
- nuclear charge: higher nuclear charge = higher e.a.
- size of atom: the larger the atom, the lower the e.a.
- e. config.: stabler configs = lower e.a.
periodic trends of electron affinity
- along a period, e.a. increases, because nuclear charge increases and atomic size decreases
- down a group, e.a. decreases, because although size and nuclear charge both increase, effect of increased size is much more pronounced
electronegativity
a measure of the ability of atoms to attract a shared pair of electrons toward itself in a covalent bond
polar covalent bond
bond in which electrons are shared unequally, resting in a partial polar charge
periodic trends of electronegativity
- along a period, electronegativity increases due to increased nuclear charge, resulting in increased attraction between nucleus and electrons
- down a group, electronegativity decreases due to increase in atomic size, resulting in reduced attraction
types of oxides
- acidic oxide
- basic oxide
- neutral oxide
- amphoteric oxide
Group I trends
Down the group:
- reactivity increases
- m.pt/b.pt increases
- density decreases
Group VII trends
down the group:
- m.pt/b.pt increases
- density increases
- reactivity decreases
transition element
- element whose atom has an incomplete d sub-shell
- or can give rise to a cation with an incomplete d sub-shell
properties of transition metals
PHYSICAL PROPERTIES
- high m.pt/b.pt
- high tensile strength
- high density
- malleable and ductile
- high thermal & electrical conductivity
CHEMICAL PROPERTIES
- form compounds with more than one oxidation no
- incomplete d-subshell
- form complex ions
- form colored compounds
- act as catalysts when either elements or compounds
why does chromium show a large jump in IE for the 7th e-?
Chromium’s config: [Ar] 3d5
- because there’s not much diff between 3d and 4s orbitals
- but taking from 3p6 is harder (also because the orbital is full)
first ionisation energy
- the amount of energy required
- to remove 1 mol of e-s from 1 mol of atoms in gaseous state
what does group number tell you about electron placement?
number of outer shell/valence electrons
what does period number tell you about electron placement?
number of occupied electron shells
Why do halogen m.pts/b.pts increase down the group?
- coz their Mr also increases down the group
- thus van der waals/intermolecular forces also increase
