Lesson Notes By Weeks and Term v5 - Grade 10
Atomic structure and periodic table – Week 7 focus
Download the Lessonotes Mobile South Africa app for faster lesson access on Android and iPhone.
Atomic structure and periodic table – Week 7 focus
Download the Lessonotes Mobile South Africa app for faster lesson access on Android and iPhone.
Subject: Physical Sciences
Class: Grade 10
Term: 1st Term
Week: 7
Theme: General lesson support
This page supports the lesson note with a companion video and a short classroom-ready summary.
For class groups and homework, share this lesson page so learners also get the summary, objectives, and full lesson context.
This week, we delve into the fascinating world of the atom and how these tiny building blocks are organised on the periodic table. Understanding atomic structure and the periodic table is fundamental to comprehending chemistry and the physical properties of matter. From the fertilizers that help grow our food to the materials used in our cell phones, understanding how atoms interact is crucial. In South Africa, with our rich mineral resources, a strong understanding of atomic structure is essential for extracting, processing, and utilizing these resources responsibly and sustainably.
2.1 Atomic Structure: Atoms are the basic building blocks of all matter. They consist of a central nucleus surrounded by electrons.
Nucleus: The nucleus contains: Protons: Positively charged particles. The number of protons determines the element's identity (atomic number).
Neutrons: Neutral (no charge) particles. Protons and neutrons contribute significantly to the atom's mass.
Electrons: Negatively charged particles that orbit the nucleus in specific energy levels or shells. Electrons are arranged in energy levels, which are designated by numbers (n=1, 2, 3, etc.), with n=1 being closest to the nucleus. Within each energy level, there are sublevels (s, p, d, f). The first energy level (n=1) has only the 's' sublevel, the second (n=2) has 's' and 'p', the third (n=3) has 's', 'p', and 'd', and so on. The number of electrons each sublevel can hold is: s=2, p=6, d=10, f=14. 2.2 Atomic Number (Z) and Mass Number (A): Atomic Number (Z): The number of protons in the nucleus of an atom. It uniquely identifies an element. For example, all atoms with 6 protons are carbon atoms.
Mass Number (A): The total number of protons and neutrons in the nucleus of an atom. A = number of protons + number of neutrons. 2.3 Isotopes: Isotopes are atoms of the same element (same atomic number) that have different numbers of neutrons, and therefore different mass numbers. For example, Carbon-12 (¹²C) and Carbon-14 (¹⁴C) are isotopes of carbon. Both have 6 protons, but ¹²C has 6 neutrons, and ¹⁴C has 8 neutrons. 2.4 Relative Atomic Mass (Ar): Because elements exist as a mixture of isotopes, we use the relative atomic mass, which is a weighted average of the masses of all naturally occurring isotopes of that element.
Calculating Relative Atomic Mass: Ar = [(% abundance of isotope 1 × mass of isotope 1) + (% abundance of isotope 2 × mass of isotope 2) + ...] / 100
Example: Chlorine has two isotopes: Chlorine-35 (⁷⁵Cl) with an abundance of 75.77% and Chlorine-37 (⁷⁷Cl) with an abundance of 24.23%. Ar(Cl) = [(75.77 × 35) + (24.23 × 37)] / 100 Ar(Cl) = (2651.95 + 896.51) / 100 Ar(Cl) = 3548.46 / 100 Ar(Cl) = 35.48 u (atomic mass units) 2.5 Electron Arrangement (Electronic Configuration): Electrons occupy specific energy levels (shells) around the nucleus. The arrangement of electrons determines an element's chemical properties. The order of filling electrons into the orbitals is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s… Valence Electrons: The electrons in the outermost energy level. These electrons are involved in chemical bonding.
Electronic Configuration Notation: E.g., Sodium (Na, Z=11): 1s² 2s² 2p⁶ 3s¹ Abbreviated Notation (Noble Gas Core): Uses the previous noble gas to shorten the configuration. E.g., Sodium (Na, Z=11): [Ne] 3s¹ 2.6 The Periodic Table: The periodic table organizes elements based on their atomic number and recurring chemical properties.
Periods: Horizontal rows. Elements in the same period have the same number of electron shells.
Groups: Vertical columns. Elements in the same group have the same number of valence electrons and similar chemical properties. For example, Group 1 (alkali metals) all have one valence electron and are highly reactive. Group 17 (halogens) have seven valence electrons and are also highly reactive. Group 18 (noble gases) have a full outer shell (8 valence electrons, except for Helium which has 2) and are generally unreactive. 2.7 Ions: Atoms can gain or lose electrons to form ions.
Cations: Positively charged ions formed when an atom loses electrons. Metals typically form cations. For example, Sodium (Na) loses one electron to form Na⁺.
Anions: Negatively charged ions formed when an atom gains electrons. Non-metals typically form anions. For example, Chlorine (Cl) gains one electron to form Cl⁻. The charge of a common ion can be predicted by the number of valence electrons an element has. Elements in group 1 tend to lose 1 electron forming ions with a +1 charge. Group 2 elements tend to lose 2 electrons forming ions with a +2 charge. Group 16 elements tend to gain 2 electrons forming ions with a -2 charge. Group 17 elements tend to gain 1 electron forming ions with a -1 charge. Guided Practice (With Solutions)
Question 1: An atom has 17 protons and 18 neutrons. a) What is its atomic number? b) What is its mass number? c) What element is it? d) How many electrons does it have if it's neutral?
Solution: a) Atomic number (Z) = number of protons = 17 b) Mass number (A) = number of protons + number of neutrons = 17 + 18 = 35 c) The element with atomic number 17 is Chlorine (Cl). d) In a neutral atom, the number of electrons equals the number of protons.
Therefore, it has 17 electrons.
Question 2: Calculate the relative atomic mass of Magnesium (Mg) if it has the following isotopic abundances: ²⁴Mg (79.0%), ²⁵Mg (10.0%), ²⁶Mg (11.0%).