1 Atomic Structure

MATTER H2 CHEMISTRY / 9729 GCE ADVANCED LEVEL © INTELEAGUE 2018 9729_CH_C1_2018 1 CHAPTER 1 Atomic Structure Learning Outcomes Candidates should be able to: (a) identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses (b) deduce the behaviour of beams of protons, neutrons and electrons in an electric field (c) describe the distribution of mass and charges within an atom (d) deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (and charge) (e) (i) describe the con tribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number (ii) distinguish between isotopes on the basis of different numbers of neutrons present (f) describe the number and relative energies of the s, p and d orbitals f or the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals (g) describe the shapes of s, p and d orbitals [knowledge of wave functions is not required] (h) state the electronic configuration of atoms and ions given the proton number (and charge) (i) explain the factors influencing the ionisation energies of elements (see the Data Booklet) (j) deduce the electronic configurations of elements from successive ionisation energy data (k) interpret successive ionisation energy data of an element in terms of the position of that element within the Periodic Table Trends and variations in atomic and physical properties For elements in the third period (sodium to chlorine), and in Group 2 (magnesium to barium) and Group 17 (chlorine to iodine) candidates should be able to: (l) recognise variation in the electronic configurations across a Period and down a Group (m) describe and explain qualitatively the trends and variations in atomic radius, ionic radius, first ionisation energy and electronegativity: (i) across a Period in terms of shielding and nuclear charge (ii) down a Group in terms of increasing number of electronic shells and nuclear charge (n) define the terms relative atomic, isotopic, molecular and formula mass, based on the 12C scale (o) calculate the relative at omic mass of an element given the relative abundances of its isotopes (p) explain what is meant by a transition element, in terms of d block elements forming one or more stable ions with partially filled d subshells (q) state the electronic configuration of a first row transition element and its ions (r) explain why atomic radii and first ionisation energies of the transition elements are relatively invariant

2 PROPERTIES OF THE 3 SUBATOMIC PARTICLES Property Proton Electron Neutron Relative charge +1 –1 0 Relative mass 1 1 1836 1 Atom: The smallest component of an element having the chemical properties of the element Isotopes: Atoms of an element that have the same proton number but different nucleon numbers Proton number: The number of protons in the nucleus of each atom of an element Nucleon number: The total number of neutrons and protons in the nucleus of an atom of an element Deflection of charged particles 𝜃 ∝ 𝑞 𝑚 Isotopic → Same number of protons Isotonic → Same number of neutrons Isoelectronic → Same number of electrons Relative atomic mass, A r: Average mass of 1 atom of an element compared with 1 12 the mass of 1 atom of 12C Relative isotopic mass: Average mass of 1 atom of an isotope of an element compared with 1 12 the mass of 1 atom of 12C Atomic orbital: Region of space with a 90% probability of finding an electron RELATIVE ABUNDANCE OF ISOTOPES The Ar of Boron, which consists of the isotopes 10B and 11B is 10.8. What is the percentage of 11B atoms in the isotopic mixture? Let a be the relative abundance of 11B atoms. Relative abundance of 10B atoms is (1 – a). a(11) + (1 – a)(10) = 10.8 a = 0.8 so % of 11B atoms in the mixture = 80.0%

3 Summary of the relationship between principal quantum shells, subshells and orbitals Principal quantum no. Subshell No. of orbitals Max no. of e– in each principal quantum shell 1 1s 1 2 2 2s 2p 1 3 8 3 3s 3p 3d 1 3 5 18 4 4s 4p 4d 4f 1 3 5 7 32 s Orbitals 1s 2s p Orbitals px py pz x y z x y z x y z x y z x y z

4 d Orbitals dx2-y2 dz2 dxy dxz dyz 1. Aufbau Principle Electrons in their ground states occupy orbitals in order of energy levels 2. Hund’s Rule of Multiplicity When filling subshells that contain >1 orbital with the same energy, each orbital must be singly occupied before electrons are paired. 3. Pauli Exclusion Principle An orbital can’t hold >2 electrons and the 2 electr ons sharing the same orbital must have opposite spins Exceptions Chromium: 1s2 2s2 2p6 3s2 3p6 3d5 4s1 Copper: 1s2 2s2 2p6 3s2 3p6 3d10 4s1 x z y z x y x y z x z y y x z

5 Cations: Remove electrons from the outermost subshell Anions: Add electrons to the lowest energy orbital Exceptions Scandium forms Sc3+ which has the 3d0 configuration Zinc forms Zn2+ which has the 3d10 configuration TRENDS IN ATOMIC RADIUS Increases down the group • Number of quantum shells increases as more electrons are added to a new principal quantum shell • Outermost electrons are further away from nucleus • Although nuclear charge increases due to the increase in number of protons, the number of quantum shells is the more important factor Decreases across the period • Nuclear charge increases due to increasing number of protons • Shielding effect remains relatively constant as electrons are added to the same outermost shell • Effective nuclear charge increases, resulting in stronger electrostatic forces of attraction between nucleus and outermost electron so it is pulled closer to the nucleus Relatively invariant across transition elements • Nuclear charge increases due to increasing number of protons • Shielding effect increases as electrons are added to the inner 3d subshell • Shielding effect nullifies, to a considerable extent, the influence of each additional proton in the nucleus • Effective nuclear charge remains almost constant

6 TRENDS IN IONIC RADIUS Decrease in ionic radius across the cations and anions • Na+, Mg2+ and Al3+ are isoelectronic as they have the same number of electrons (10) • P3–, S2– and Cl– are isoelectronic as they have the same number of electrons (18) • Across the 2 isoelectronic series, nuclear charge increases due to increasing number of protons • Shielding effect remains the same due to the same number of electrons • Increase in effective nuclear charge and stronger attraction between outermost electron and nucleus • Sharp increase in ionic radius from cationic to anionic series is due to the anions having 1 more quantum shell than cations despite their higher nuclear charge 1st ionisation energy: Energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of unipositively charged gaseous ions X (g) → X+ (g) + e– Ideal gas has negligible IMF so no extra energy is needed to overcome IMF, only for 1st I.E. 2nd ionisation en ergy: Energy needed to remove 1 mole of electrons from 1 mole of unipositively charged gaseous ions to form 1 mole of gaseous ions with double positive charge X+ (g) → X2+ (g) + e– Na+ Mg2+ Al3+ Si4+ P3- S2- Cl- Ionic Radius/nm

7 Successive Ionisation Energies of Mg • 1st big increase from 2 nd to 3 rd ionisation energies implies that the 3 rd electron is in an inner quantum shell so there are 2 electrons in the valence 3s subshell → Mg is a Group 2 element • Gradual increase from 3rd to 10th ionisation energies implies that the next 8 electrons are removed from the 2nd quantum shell • Slightly greater increase from 8 th to 9th ionisation energies as more energy is needed to remove the 2s electrons which are closer to the nucleus than the 2p electrons