211 - Periodicity and Tred

trends in radial distribution functions (3)

1. as n increases, for any value of l, the maximum RDF increases

2. as l increases, the maximum RDF gets closer to the nucleus

3. for 3s, 3p and 3d orbitals, the total electron density within 400 ppm of the nucleus follows the order s>p>d

electron shielding definition

phenomenon in which the inner/core electrons create repulsive forces that partially block the positive charge of the nucleus felt by the valence electrons

Z_eff definition

the effective nucleus charge is the net charge of the nucleus felt by the valence electrons

trends in Z_eff (2)

1. increases across the periods as nuclear charge increases, but electrons are added to the same energy level so shielding remains similar but valence electrons experience greater nuclear attraction

2. increases significantly from the end of one period to the beginning of the next because electron are added to higher energy orbitals so inner electrons increase in shielding ability

slater’s rules over simplifications (2)

1. treats Zeff felt by s and p orbitals of the same n value as the same

2. treats n-1 orbitals as pure 100% shielders so there is no increase in Zeff down in the rows (except Al → Ga due to d block contraction effect)

Hund’s Rule definition

the ground state of an atom has maximum spin multiplicity

Exchange energy definition

the quantum stabilisation energy that arises due to the exchange of electrons with parallel spins in degenerate orbitals

Elements that are exceptions to the Aufbau principle (6)

1. Cu

2. Co

3. Mo

4. Pd

5. Ag

6. Au

first ionisation energy definition

the minimum energy required to remove the least tightly held electron from each atom in a mole of gaseous atoms in their ground state

first ionisation energy trends (4)

1. across a period, ionisation energy increases as nuclear charge increases and electrons are added to the same energy level which increases the electrostatic attraction between the nucleus and the valence electrons

2. down a group, ionisation energy decreases as electrons are added to higher energy orbitals resulting in greater distance from the nucleus and increased electron shielding from inner electrons

3. From the end of one period to the beginning of the next, ionisation energy decreases significantly as additional electrons are added to higher energy orbitals and core electrons increase in shielding ability

4. First period elements have very high ionisation energy due to an unusually high overall percentage of nuclear charge felt by the valence electrons

exceptions to the trends in first ionisation energy (2)

1. Ga has a higher than expected IE compared to Al → d-block contraction effect

2. Tl (Thallium) has a higher IE than In due to f-block contraction effect

d/f contraction effect definition

phenomenon when the d (Ga) and f (Tl) orbitals are introduced that results in a less significant decrease in atomic size across a period because the d/f electrons are added to an inner shell and are not effective at shielding so the valence electrons still experience high Zeff

selected trends in first ionisation energy (2)

1. IE(B) < IE(Be): boron has a lower ionisation energy than beryllium because its valence electron is in the higher energy 2p subshell which is less tightly held due to shielding from he filled 2s subshell 

2. IE(O) < IE(N): oxygen has a lower ionisation energy than nitrogen because pairing of electrons in the 2p orbital (O) increases electron-electron repulsion and reduces exchange energy stabilisation, making it easier to remove an electron from oxygen 

electron affinity definition

the change in energy when one mole of gaseous atoms in their ground state each gains one electron to form one mole of gaseous negative ions

trends in electron affinity (4)

1. across a period, electron affinity decreases/becomes more negative because nuclear charge increases so additional electrons experience more nuclear attraction so more energy is released when an electron is added

2. down a group, electron affinity increases/becomes more positive because electrons are added to higher energy orbitals and experience less nuclear attraction so less energy is released when an electron is added

3. 2^nd period electron affinities are lower than 3^rd period electron affinities due to their small atomic radii which produces high electron-electron repulsions when electrons are added

4. 1^st group electron affinities are higher than 13^th group electron affinities because 1^st group elements gain electrons to complete a stable s-orbitals while 13^th group elements must place additional electrons in the higher energy p-orbital, which increases repulsion

 kinks in the electron affinity trends (2)

1. N’s electron affinity is less negative than C and O (instead of trend that electron affinity becomes more negative across a period)

   • C gains an electron with minimal repulsion as the electron is placed in the 2p orbital

   • N has a half-filled 2p orbital so an additional electron is forced to pair up, which increases electron-electron repulsions 

   • O readily gains an electron to complete the p-orbital  

2. P’s electron affinity is less negative than Si and S (instead of trend that electron affinity becomes more negative across a period)

   • Si gains an electron with minimal repulsion as the electron is placed in the 3p orbital 

   • P has a half-filled 3p orbital so an additional electron is forced to pair up, which increases electron-electron repulsions 

   • S readily gains an electron to complete the p-orbital

Pauling electronegativity definition

the ability of an atom in a covalent bond to attract towards itself electron density

Mulliken electronegativity definition

the average of an atom’s ionisation energy and electron affinity, representing its tendency to accept or donate electrons

Allred-Rochow electronegativity definition

the electrostatic force exerted by the nucleus on valence electrons based on effect nuclear charge and covalent radius

Allen electronegativity definition

the average energy of the valence electrons in a ground state free atom

trends in electronegativity (3)

1. increases across a period due to decreased atomic radius which leads to an increase in electrostatic attraction between the nucleus and valence electrons

2. decreases down a group due to an increase in atomic radius which leads to a decrease in electrostatic attraction between the nucleus and valence electrons

3. particularly high for 1^st and 2^nd period elements as they are smaller with fewer shells so shielding is minimal

Kinks in the trends in electronegativity (2)

1. Ga has a higher electronegativity than Al due to the d block contraction effect which results in an increase in Zeff due to weak shielding of the d electrons so valence electrons experience a greater nuclear charge

2. Tl has a higher electronegativity than In due to the f block contraction effect results in an increase in Zeff due to the weak shielding of the f electrons so valence electrons experience greater nuclear charge

polarisability definition

how easily the electron cloud of an atom/ion can be distorted by an external electric field

polarising power definition

the ability of an atom/ion to distort the electron cloud of another species

trends in polarising power (2)

1. increases across a period, as nuclear charge increases, ionic radius decreases, charge density increases and cations become more compact and better at distorting anion electron clouds

2. decreases down a group as ionic radius increases, shielding increases, charge density decreases and cations become larger and less effective at distorting anion electron clouds

Van der Waals Forces definition

weak intermolecular forces that arise due to temporary or permanent dipoles

types of Van der Waals Forces (3)

1. London dispersion forces

2. dipole-dipole interactions

3. dipole-induced dipole interactions

London dispersion forces definition

temporary attractive forces due to fluctuating electron clouds creating instantaneous dipoles

dipole-dipole interactions

attraction between permanent dipoles in polar molecules

dipole-induced dipole interactions

permanent dipole induces a temporary dipole in a non-polar molecule

Hydrogen bonding

a special case of dipole-dipole interactions in which a hydrogen atom covalently bonded to a highly electronegative atom (N, O or F) exhibits an attraction to another electronegative atom

types of covalent bonds (3)

1. sigma

2. pi

3. delta

sigma bonds

electron density concentrated between 2 nuclei

pi bonds

electron density concentrated above and below the plane of the molecule

delta bonds

electron density concentrated in two symmetrical regions above and below the bonding axis

trends in sigma and pi bond strength (1)

decrease down a group due to increase in atomic size which produces more diffuse orbitals with less effective overlap

special case in sigma and pi bond strength trend

Pi and sigma bonds are comparable and strong in C-C due to many different hybridisations but N-N sigma bond strength decreases due to N’s small size and the lone pair of electrons which create repulsions that decrease bond strength 

catenation definition

process in which an element preferentially forms allotropes with predominance of sigma bonds

oxidation state definition

the charge which would result if the electrons in each bond to an atom were assigned to the more electronegative atom

valence state definition

number of valence electrons used in bonding

trend in s-block oxidation states

Oxidation states correspond to elements group number 

type of compounds formed by s block elements (G1, 2)

lattices

type of compounds formed by p block elements

covalent compounds

promotion energy definition

energy required to excite an electron from its ground state to a higher energy state

inert pair effect

the tendency of the 2 outermost electrons in the atomic s-orbital to remain unshared in compounds of post-transition metals

types of solids (4)

1. molecular covalent

2. covalent network

3. metallic

4. ionic

molecular covalent solids

solids made up of discrete molecules held together by intermolecular forces

network covalent solids

solids where atoms are bonded together in a continuous 3D network entirely by strong covalent bonds

metallic solids

solids consisting of a positive metal ion in a sea of delocalized electrons that move freely throughout the structure

ionic solids

solids composed of positive and negative ions held together by strong electrostatic forces

crystalline solids

solids with a regular repeating atomic arrangement

amorphous solids

solids that lack a repeating pattern of long range order, resulting in a disordered structure

coordination number

Number of atoms bonded to the atom of interest