Lesson Notes By Weeks and Term v5 - Grade 10

Lesson Video

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Performance objectives

• Describe the structure of an atom (protons, neutrons, electrons).
• Define and use atomic number (Z) and mass number (A).
• Explain isotopes and calculate relative atomic mass.
• Write electronic configurations for elements.
• Relate electronic configuration to the periodic table.

Lesson summary

This week, we delve into the fascinating world of atomic structure and the periodic table. Understanding these concepts is crucial because it forms the bedrock of all chemistry and much of physics. From the fertilizers that support our agriculture to the electronic devices we use daily, our understanding of atomic structure and the periodic table underpins it all. In South Africa, a strong grasp of these concepts is essential for future scientists, engineers, doctors, and anyone involved in industries reliant on chemical processes, mineral resources, and advanced materials.

Lesson notes

2. 1. Atomic Structure An atom is the smallest unit of an element that can exist. It's composed of three primary subatomic particles: Protons: Positively charged particles located in the nucleus of the atom. The number of protons determines the element. Its relative mass is approximately 1 atomic mass unit (amu).

Neutrons: Neutral (no charge) particles located in the nucleus. Neutrons contribute to the mass of the atom but do not affect its chemical properties. Its relative mass is approximately 1 amu.

Electrons: Negatively charged particles that orbit the nucleus in specific energy levels or shells. Their relative mass is negligible compared to protons and neutrons (approximately 1/1836 amu). The nucleus, containing protons and neutrons, accounts for almost all of the atom's mass. 2.

2. Atomic Number (Z) and Mass Number (A)

Atomic Number (Z): The number of protons in the nucleus of an atom. This number uniquely identifies an element. For example, all atoms with Z=6 are carbon atoms.

Mass Number (A): The total number of protons and neutrons in the nucleus of an atom. We can calculate the number of neutrons (N) using the formula: A = Z + N or N = A - Z Example 1: Consider an atom of Sodium (Na) with A = 23 and Z =

1

1. Number of protons = Z = 11 Number of electrons = Z = 11 (for a neutral atom) Number of neutrons = A - Z = 23 - 11 = 12 Example 2: Consider the ion, Mg 2+ with A = 24 and Z =

1

2. Number of protons = Z = 12 Number of neutrons = A - Z = 24 - 12 = 12 Number of electrons = Z - charge = 12 - 2 = 10 (It lost 2 electrons) 2.

3. Isotopes and Relative Atomic Mass Isotopes: Atoms of the same element (same number of protons) that have different numbers of neutrons. This means they have the same atomic number (Z) but different mass numbers (A). For example, Carbon-12 ( 12 C) and Carbon-14 ( 14 C) are isotopes of carbon. Both have 6 protons, but 12 C has 6 neutrons, while 14 C has 8 neutrons.

Relative Atomic Mass (Ar): The weighted average of the masses of all the naturally occurring isotopes of an element, relative to 1/12th the mass of a carbon-12 atom. The relative atomic mass is found on the periodic table.

To calculate the relative atomic mass: Ar = [(% Abundance of Isotope 1 x Mass of Isotope 1) + (% Abundance of Isotope 2 x Mass of Isotope 2) + ... ] / 100 Example 3: Chlorine has two naturally occurring isotopes: Chlorine-35 ( 35 Cl) with an abundance of 75.77% and Chlorine-37 ( 37 Cl) with an abundance of 24.23%. Calculate the relative atomic mass of chlorine. Ar = [(75.77 x 35) + (24.23 x 37)] / 100 Ar = (2651.95 + 896.51) / 100 Ar = 3548.46 / 100 Ar = 35.48 amu (approximately) 2.

4. Electronic Configuration Electrons are arranged in specific energy levels (shells) around the nucleus. These shells are numbered 1, 2, 3, etc., starting from the shell closest to the nucleus. Each shell can hold a maximum number of electrons given by 2n 2 , where n is the shell number. So, shell 1 can hold 2 electrons, shell 2 can hold 8 electrons, shell 3 can hold 18 electrons, and so on. Within each energy level, there are sub-levels (also known as subshells) designated as s, p, d, and f. The s sub-level can hold a maximum of 2 electrons. The p sub-level can hold a maximum of 6 electrons. The d sub-level can hold a maximum of 10 electrons. The f sub-level can hold a maximum of 14 electrons. The order in which these sub-levels are filled follows the Aufbau principle and Hund's rule.

A simplified order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p… Electronic configuration describes how the electrons are arranged in the different energy levels and sub-levels.

Example 4: Write the electronic configuration of Oxygen (O), Z =

8. Oxygen has 8 electrons. 1s can hold 2 electrons: 1s 2 2s can hold 2 electrons: 2s 2 2p can hold the remaining 4 electrons: 2p 4 Therefore, the electronic configuration of Oxygen is 1s 2 2s 2 2p 4 .

Example 5: Write the electronic configuration of Potassium (K), Z =

1

9. Potassium has 19 electrons. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 (

Note: 4s fills before 3d) Therefore, the electronic configuration of Potassium is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 . 2.

5. The Periodic Table The periodic table is an arrangement of elements in order of increasing atomic number, grouped in such a way that elements with similar chemical properties fall in the same column.

Periods: The horizontal rows in the periodic table. The period number corresponds to the highest energy level (shell) occupied by electrons in the element's electronic configuration. For example, elements in the 3rd period have their outermost electrons in the 3rd energy level.

Groups: The vertical columns in the periodic table. Elements in the same group have the same number of valence electrons (electrons in the outermost shell), which determines their similar chemical properties. For example, Group 1 elements (alkali metals) all have one valence electron and are highly reactive. 2.6.