ACSI 2022 HL Notes Atomic Structure (Teacher)

IBDP Chemistry HL/ Atomic Structure Page 1 Anglo − Chinese School (Independent) Year 5 (2022) IBDP Chemistry HL (IBDP syllabus Topic 2) 2.1 The nuclear atom - Essential Idea: the mass of an atom is concentrated in its minute positively charged nucleus. 2.2 Electron configuration - Essential Idea: the electron configuration of an atom can be deduced from its atomic number. (IBDP syllabus Topic 12) 12. 1 Electrons in atoms - Essential Idea: the quantized nature of energy transitions is related to the energy states of electrons in atoms and molecules TOPIC 2 ATOMIC STRUCTURE TEACHER COPY – WITH SUGGESTED SOLUTIONS

IBDP Chemistry HL/Atomic Structure Page 2 2.1 The nuclear atom Essential idea: The mass of an atom is concentrated in its minute, positively charged nucleus.  Nature of Science: • Evidence and improvements in instrumentation—alpha particles were used in the development of the nuclear model of the atom that was first proposed by Rutherford. (1.8) • Paradigm shifts—the subatomic particle theory of matter represents a paradigm shift in science that occurred in the late 1800s. (2.3) Understandings: • Atoms contain a positively charged dense nucleus composed of protons and neutrons (nucleons). • Negatively charged electrons occupy the space outside the nucleus. • The mass spectrometer is used to determine the relative atomic mass of an element from its isotopic composition. Applications and skills: • Use of the nuclear symbol notation XΑ Ζ to deduce the number of protons, neutrons and electrons in atoms and ions. • Calculations involving non-integer relative atomic masses and abundance of isotopes from given data, including mass spectra. 2.2 Electronic configuration Essential idea: The electronic configuration of an atom can be deduced from its atomic number.  Nature of Science: • Developments in scientific research follow improvements in apparatus —the use of electricity and magnetism in Thomson’s cathode rays.(1.8) • Theories being superseded—quantum mechanics is among the most current models of the atom. (1.9) • Use theories to explain natural phenomena—line spectra explained by the Bohr model of the atom. (2.2) Understandings: • Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level. • The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies. • The main energy level or shell is given an integer number, n, and can hold a maximum number of electrons, 2n 2. • A more detailed model of the atom describes the division of the main energy level into s, p, d and f sub-levels of successively higher energies. • Sub-levels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron. • Each orbital has a defined energy state for a given electronic configuration and chemical environment and can hold two electrons of opposite spin.

IBDP Chemistry HL/Atomic Structure Page 3 Applications and skills: • Description of the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum. • Distinction between a continuous spectrum and a line spectrum. • Description of the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels. • Recognition of the shape of an s atomic orbital and the px, py and pz atomic orbitals. • Application of the Aufbau principle, Hund’s rule and the Pauli Exclusion Principle to write electron configurations for atoms and ions up to Z = 36. Guidance: • Details of the electromagnetic spectrum are given in the data booklet in section 3. • The names of the different series in the hydrogen line emission spectrum are not required. • Full electron configurations (eg 1s 22s22p63s23p4) and condensed electron configurations (eg [Ne] 3s23p4) should be covered. • Orbital diagrams should be used to represent the character and relative energy of orbitals. • The electron configurations of Cr and Cu as exceptions should be covered. 12.1 Electronic configuration Essential idea: The quantized nature of energy transitions is related to the energy states of electrons in atoms and molecules.  Nature of Science: • Experimental evidence to support theories – emission spectra provide evidence for the existence of energy levels. Understandings: • In an emission spectrum, the limit of the convergence at higher frequency corresponds to the first ionisation energy • Trends in first ionisation energy across periods account for the existence of main energy levels and sub-levels in atoms. • Successive ionisation energy data for an element give information that shows relations to electron configuration. Applications and skills: • Solving problems using E = hv • Calculation of the value of the first ionisation energy from spectral data which gives the wavelength or frequency of the convergence limit. • Deduction of the group of an element from its successive ionisation energy data. • Explanation of the trends and discontinuities in the first ionisation energy across a period. Guidance: • The value of Planck’s constant (h) and E = hv are given in the data booklet in sections 1 and 2. • Use of the Rydberg formula is not expected in calculations of ionisation energy. Important information for the topic 𝑐𝑐 = 𝑣𝑣λ 𝐸𝐸 = ℎ𝑣𝑣 Speed of light = 3.00 × 108 m s−1 Planck’s constant (ℎ) = 6.63 × 10−34 J s

IBDP Chemistry HL/Atomic Structure Page 4 2.1. The Nuclear Atom 2.1.1 The Subatomic Particles In An Atom An atom is the smallest component of an element. Atoms contain a positively charged dense nucleus composed of protons and neutrons (nucleons). Negatively charged electrons occupy the space outside the nucleus. Relative masses and charges of the three subatomic particles: Subatomic particle Proton (p) Neutron (n) Electron (e) Mass (kg) 1.672 ×10−27 1.674 ×10-27 9.109 ×10−31 Charge (× 10−19 C) 1.602 ×10−19 0 -1.602 ×10−19 Relative charge + 1 0 – 1 An atom can be represented by the nuclear symbol notation XΑ Ζ , where X is the symbol of the element. • A = total number of protons and neutrons (mass number or nucleon number) • Z = number of protons (atomic number or proton number) • (A − Z) = number of neutrons. • Atoms are electrically neutral as the number of protons is equal to the number of electrons. • Isoelectronic: Atoms / ions that have the same number of electrons. • Isotonic: Atoms / ions that have the same number of neutrons. NOS - Background to atomic theory and Dalton’s atomic theory. NOS – Thomson’s “plum-pudding” model of the atom. NOS – Rutherford’s gold foil experiment Electrons carry a negative charge of –1e where e = 1.6 × 10–19 C Protons carry a positive charge of +1e Neutrons carry no charge Scan me to view a video on the evolution of models of the atom

IBDP Chemistry HL/Atomic Structure Page 5 2.1.2 Isotopic Masses Isotopes are atoms of the same element with a different number of neutrons. They have the same atomic number but different mass numbers. • Isotopes have the same chemical properties but different physical properties. • Isotopes of elements have different nucleon num bers, relative isotopic masses and numbers of neutrons. • The isotope with more neutrons is referred to as the heavier isotope. • Isotopes of some elements are radioactive as the nuclei of these atoms are unstable and breakdown spontaneously. When they break dow