3. Periodicity Flashcards
according to which property are elements arranged in the PTE
their atomic number/number of protons/Z
metals
- electrons
- charge of ions
- conductivity
- form
- melting point/boiling point
- lose electrons
- positive charge (cations)
- good heat and electricity conductors (in solid state and when liquid/molten)
- malleable and ductile
- high melting and boiling point (hard to break metallic bonds - strong intermolecular bonds)
non-metals
- electrons
- charge of ions
- conductivity
- form
- melting point/boiling point
- gain electrons
- negative charge (anions)
- non-conductors/isolators of both heat and electricity
- brittle solids
- low melting point - many in gaseous state at room temperature (sublimation means that the boiling point is very low)
octet rule
elements usually react in such a way as to obtain the electron configuration of the noble gas closest to them in the PTE
isoelectric
ions of different elements that have the same number of electrons
trends in the PTE
1|atomic radius
2|electron affinity
3|size of ions
4|ionization energy
5|electronegativity
atomic radius
size of an atom
increases down and to the left
electron affinity
measure of the energy change when an electron is added to a neutral atom to form a negative ion (affinity to gaining electrons)
absolute value increasing up and to the right
size of ions
cations<atoms<anions - number of protons stays the same but number of electrons increases - weaker attraction force, further away from the nucleus - number of shells may change as well
ionization energy
the energy needed to remove one mole of electrons from one mole of an isolated, gaseous atom/ion
increasing up and to the right
electronegativity
chemical property that describes the ability of an atom (or a functional group) to attract electrons toward itself in a covalent bond (bond not equally shared)
increases up and to the right
the smaller the atomic radius the greater the electronegativity
FONCl (4 most electronegative elements)
alkali metals
- I. group (H excluded)
- soft elements
- rigorous reaction with water and air
- reactivity increases down the group (because atomic radius gets greater and so electrons are more easily lost because the attraction force is smaller)
- strong reducing agents - small nuclear charge so they can easily lose e- (want to get to noble gas e- configuration)
- melting point: increases up the group (inverse proportional with atomic radius - easier to break them apart because their nuclei are further apart and the attraction force is weaker - metallic bond)
- when recating with halogens, produce salts (vigorous reaction)
halogens
- VII. group
- diatomic molecules
- reactive elements - only have to gain one electron to achieve n-g- configuration
- reactivity increases up the group (atomic radius gets larger so it’s harder to gain electrons because the attraction force of the nucleus is lesser)
- strong oxidizing agents (proportional to reactivity) - want to gain electrons
F2>Cl2>Br2>I2 - melting point: increases down the group (proportional with molar mass) - intermolecular forces in covalently bonded compounds: Van der Waals bond strength is proportional with molar mass
- when reacting with alkali, produce salts (vigorous reaction)
- presence of halogens can be detected by placing silver nitrate solution
why do C and Si peak on the melting point graph and why is C melting point greater than Si m.p.
because they have giant covalent structures that are difficult to break (require a lot of energy)
C>Si because it has one less shell and therefore a smaller radius, greater attraction force between nucleus and shells and more E is needed to break the bonds
metalloids conductivity
semi-conductors
metallic oxides
- ionic compounds
- high melting and boiling points
- conduct electricity when molten/aqueous but not solid
- when reacted with water form basic solutions
non-metallic oxides
- ionic compounds
- variable melting/boiling points (SiO2)
- don’t conduct electricity at all
- when reacted with water form acidic solutions
why do non-metals have lower melting/boiling points than metals
they have much weaker intermolecular forces - easier to separate
why is 2+ the most common valency in transition metals?
because they first lose electrons from 4s subshell
catalytic behaviour (processes)
1| Haber process
2|contact process
3|hydrogenation reactions
4|hydrogen peroxide breakdown
Haber process
- iron (Fe)
N2 + 3H2 <-> 2NH3
contact process
- vanadium (V)
2SO2 + O2 <-> (V2O5) 2SO3 - all gaseous state
hydrogenation reactions
- nickel (Ni)
C2H4 + H2 -> (Ni) C2H6
hydrogen peroxide breakdown
- manganese (IV) oxide (MnO2)
2H2O2 -> (MnO2) 2H2O + O2
copper/chromium/manganese colored compounds
blue
yellow/orange/red
purplish
why don’t Sc3+ and Zn2+ show usual transition metal behavior
Sc3+ has completely empty d subshell
Zn2+ has a completely fulfilled d subshell
- electron transfer from split orbitals not possible - transparent solutions
- they can form colored complexes but they prefer not to (Sc2+ and Zn+)
ligands
neutral molecules or anions that contain a non-bonding electron pair that form coordinate covalent bonds with metal ions - to form complex ions
(have to provide all electrons for the bonding because metal ions have none to spare)
coordination number
number of lone pairs bonded to the metal ion
complex ion shapes according to coordination number
2: linear
4: tetrahedral, square planar
6: octahedral
types of ligands
1|monodentate
2|bidentate
3|polydentate
monodentate ligands
- H2O (water)
- F- (fluorine)
- Cl- (chloride)
- NH3 (ammonia)
- OH- (hydroxide ion)
- SCN- (thiocyanate ion)
- CN- (cyanide)
bidentate ligands
- H2NCH2CH2NH2 (1,2 diaminoethane)
- (C2O4)2- (ethanedioate ion)
polydentate ligands
- EDTA (ethylenediaminetetraacetic acid)
factors affecting the color of transition metal complexes
1|nature of transition metal
2|oxidation state of the metal
3|ligand (identity)
4|the stereochemistry of the complex (geometric shape)
which complex ions don’t have color
- non-transition metal
- Sc3+ and Zn2+ because their d subshells are either empty or full
coordinate covalent bonding
mutually shared electron pair with only one source - the ligand - also called dative bonding
how are d subshell orbitals split
- ligands split them
- 3 orbitals on a lower energy level and 2 on a higher (electrons jump from those 3 to 2)
- the energy difference between split orbitals is determined by the position of the ligand in the spectrochemical series
how is color created by splitting of d orbitals
- electrons that jump from lower to higher E level absorb energy from surrounding light to be promoted, stay there (never emit) and reflect the complementary color
- e- jump because they want to be single (to get away from their orbital pair)
spectrochemical series
I-<Br-<S^2-<Cl-<F-<OH-<H2O<SCN-<NH3<CN-<CO
the stronger the ligand, the…
…bigger the E difference between orbitals/the bigger the E absorbed, the smaller the wavelength absorbed, the bigger the wavelength reflected, the smaller the E reflected
- absorb high E value and reflects low E value (complementary colors)
