1.4 Learning Outcomes down a group  explain the general trends in values of atomic radii (covalent radii only) down a group across a period (main group elements only) define and explain first ionisation energy explain the general trends in first ionisation energy values: across a period (main group elements) and explain the exceptions to the general trends across a period define and explain second and successive ionisation energies  

1.4 b Learning Outcomes describe how second and successive ionisation energies provide evidence for energy levels recognise the relationship and trends in successive ionisation energies of an individual element explain how chemical properties of elements depend on their electronic structure explain how atomic radius, screening effect and nuclear charge account for general trends in properties of elements in groups I and VII  

Trends in the Periodic Table of Elements Atomic radius Ionisation energy Electronegativity Two reasons required for each trend - down a group - across a period

Atomic Radius Hard to determine using one atom Why? can’t define the size of the atom exactly because orbitals are areas of high probability of finding an electron So position of electron hard to measure Nucleus easier to measure or pinpoint

Atomic Radius [covalent radius] Half the distance / between the nuclei of two atoms / of the same element / joined by a single covalent bond. Distance between two hydrogen nuclei H-H = 0.074 nm [1 nm = 10-9 m ] Atomic radius = 0.037 nm

Across a Period [left to right] Atomic Radius gets smaller Reasons Increasing nuclear charge so electrons pulled in more No increase in screening effect from electrons already there

Trends in the Periodic Table Terms The nuclear charge [i.e. the charge in the nucleus as a result of the number of protons in the nucleus] – check cartoon The screening effect of inner electrons - The screening effect is observed in atoms with an atomic number of greater than 2. In these atoms, electrons in inner energy levels shield outer electrons from the attraction of protons in the nucleus

Electronic Screening Effect Electrons on inner sublevels shield the attraction of protons in the nucleus for outer electrons

Down a Group Radius gets bigger Reasons New main energy level [shell] filling Increased nuclear charge cancelled out to some extent by increased screening effect of an extra layer of e-s in sublevels nearer the nucleus e- is also further away to begin with.

Atomic (Neutral) versus Ionic Size Metals Nonmetals Group 1 Al 143 50 e- Group 13 Group 17 e- e- 152 186 227 Li Na K 60 Li+ F- 136 F Cl Br 64 99 114 e- e- 95 Na+ Cl- 181 Al3+ e- e- 133 K+ Br- 195 Cations are smaller than parent atoms Anions are larger than parent atoms

First Ionisation Energy The energy required to completely remove the ‘most loosely bound’ electron from a [mole of] ‘neutral gaseous atom’ [s], in its [their] ‘ground state’ Na(g) = Na+(g) + e- +494 kJ mol-1 States must be given (e.g. (g))

Across a Period Ionisation energy Increases (i) Increasing nuclear charge Number of protons increases so attraction increases Electrons in same main energy level (ii) Decreasing atomic radius Atomic radius decreases so electrons are closer to positive nucleus so held tighter.

Down a Group Decreases [gets less] (i) Increasing atomic radius means electron further from attractive force of the nucleus (ii) Screening effect of inner electrons Positive charge of nucleus is increasing but inner e- shells shield the outer electron from this increased charge.

Exceptions to general trend Look at the general trend. Argon has the highest value. Why? e- from ‘full’ 3p sublevel First ionisation Energies Are any elements higher than general trend? Yes Mg and P Why? Mg: e- from ‘full’ 3s sublevel P: e- from ‘half-filled’ 3p sub-level Must state 3 in all cases

Second Ionisation Energy The removal of the second electron from a monopositive positive ion (formed when the first electron is removed.) Na+(g) = Na2+(g) + e- (States must be given)

Further evidence for the existence of Energy Levels Take a Mg atom and remove the electrons one by one using energy – measure the amount used in each case. Plot a graph of the successive ionisation energies. Use log10 because values so big.

Being held by average of 1 proton 2 layers of shielding Remove first electron Being held by average of 1 proton 2 layers of shielding e- Energy Requird 1 2 3 4 5 6 7 8 9 10 11 12 738kJ 12+

Mg+ Remove first electron Being held by average of 1 proton 2 layers of shielding e- Energy Requird 1 2 3 4 5 6 7 8 9 10 11 12 738kJ 12+ Notice that the radius of the ion has decreased Mg+

Mg2+ Remove second electron Energy Requird 1 2 3 4 5 6 7 8 9 10 11 12 Remove second electron Being held by average of 1.1 protons (Increased effective nuclear charge) Still 2 layers of shielding Atom slightly smaller Harder to remove [+712] 738kJ 1450 kJ 12+ Notice that the radius of the ion is much smaller Mg2+

Remove third electron Being held by average of 1.2 protons (Increased effective nuclear charge) Now only 1 layer of shielding Radius now much smaller Removing e- from a full shell Much harder to remove [+6282] e- Energy Requird 1 2 3 4 5 6 7 8 9 10 11 12 738kJ 1450 kJ 7732 kJ 12+

Remove fourth electron Being held by average of 1.33 protons (Increased effective nuclear charge) Now only 1 layer of shielding Atom a bit smaller Harder to remove [+ 2808] e- Energy Requird 1 738 kJ 2 1 450 kJ 3 7 732 kJ 4 5 6 7 8 9 10 11 12 10 540 kJ 12+

Huge jump e- removed from full n = 1 main level Energy [log] Huge jump e- removed from full n = 1 main level 135000 6000 Big jump e- removed from full n = 2 main level 700 Electrons removed

Successive Ionization Energies Potassium - 1s2,2s2,2p6,3s2, 3p6, 4s1 1s2 Very Important: Proof of Existence of E levels 2s1 2p6 3s1 3p6 4s1

Trends in Electronegativity Values Electro negativity: Is the relative attraction of an atom for shared pairs of electrons in a covalent bond. The most commonly used scale is the Pauling Scale which runs from zero [ 4 is high electronegativity] to four

Going across a period, Electronegativity increases

WHY ??? Increasing nuclear charge Decreasing atomic radius

Going down a group, Electronegativity decreases

WHY ??? Increasing atomic radius Increasing Screening effect

Electronegativity Values

Summary of Trends

Exam Q’s 2002

Exam Q’s 2002

Exam Q’s 2004

Exam Q’s 2004

Exam Q’s 2009

Exam Q’s 2009

Exam Q’s 2009

Exam Q’s 2013

Exam Q’s 2013

Exam Q’s 2013

Exam Q’s 2014 Bohr E level: The definite energy of an electron in an atom / shell / orbit Orbital: region where the probability of finding an electron is 95% 1s2, 2s2, 2p6, 3s2, 3p2

Exam Q’s 2014 Write the electron configuration (s, p) of an atom of silicon showing the distribution of electrons in atomic orbitals in the ground state. (6) Hence, state how many (i) main energy levels, (ii) atomic orbitals, are occupied in the silicon atom in its ground state. 1s2, 2s2, 2p6, 3s2, 3p2 (i) main energy levels = 3 (ii) atomic orbitals = 8

Exam Q’s 2014 (b) Define first ionisation energy. (6) Explain why the first ionisation energy value of silicon is: (i) greater than that of aluminium, (ii) less than that of carbon. (9) The minimum energy required to remove the most loosely bound electron // from a gaseous atom in its ground state i)Greater nuclear charge // smaller atomic radius ii)Greater atomic radius so most loosely bound electron is further from nucleus // greater screening as more shells

Exam Q’s 2014 Line emission spectra Sharp increase from 4th to 5th shows that this is the first electron to be removed from the second shell Sharp increase from 12th to 13th shows that this is the first electron to be removed from the first shell Line emission spectra