Lesson Notes By Weeks and Term v5 - Grade 10

Lesson Video

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Performance objectives

• Define protons, neutrons, and electrons by mass and charge.
• Determine the particle count in atoms and ions.
• Explain isotopes and calculate relative atomic mass.
• Describe the periodic table's arrangement.
• Identify trends like atomic size and metallic character.

Lesson summary

This week, we delve into the fascinating world of the atom and how understanding its structure allows us to organize the elements in the periodic table. Atomic structure is fundamental to understanding all chemical reactions and the properties of matter. The periodic table isn't just a chart; it's a powerful tool that reveals relationships between elements and allows us to predict their behavior. In South Africa, a strong understanding of these concepts is crucial for various industries, including mining (understanding the composition and extraction of minerals), agriculture (understanding fertilizers and soil chemistry), and manufacturing (developing new materials).

Lesson notes

2.1 Subatomic Particles Atoms, the fundamental building blocks of matter, are composed of three primary subatomic particles: Protons: Positively charged particles located in the nucleus (the atom's central core). The number of protons defines the element; it's the atomic number (Z). Each element has a unique atomic number.

Neutrons: Neutral (no charge) particles also located in the nucleus. The number of neutrons can vary within an element, leading to isotopes (explained below).

Electrons: Negatively charged particles orbiting the nucleus in specific energy levels or shells. In a neutral atom, the number of electrons equals the number of protons. Electrons are responsible for chemical bonding.

Relative Mass and Charge: | Particle | Relative Mass | Relative Charge | |---|---|---| | Proton | 1 | +1 | | Neutron | 1 | 0 | | Electron | ~1/1836 (negligible) | -1 | 2.2 Atomic Number, Mass Number, and Isotopes Atomic Number (Z): The number of protons in the nucleus of an atom. This is what identifies the 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 = Z + N (where N is the number of neutrons).

Isotopes: Atoms of the same element (same number of protons) but with different numbers of neutrons. Because they have different numbers of neutrons, isotopes have different mass numbers. For example, Carbon-12 (¹²C) has 6 protons and 6 neutrons, while Carbon-14 (¹⁴C) has 6 protons and 8 neutrons. Both are carbon, but they are different isotopes.

Representing Isotopes: Isotopes are often represented in the form: A X, where: X is the element symbol A is the mass number 2.3 Ions: Ions are formed when atoms gain or lose electrons.

Cations: Positively charged ions formed when an atom loses electrons.

Anions: Negatively charged ions formed when an atom gains electrons. 2.4 Relative Atomic Mass The relative atomic mass (Ar) of an element is the weighted average of the masses of its naturally occurring isotopes, relative to the mass of Carbon-

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2. Because isotopes exist, the atomic mass displayed on the periodic table is rarely a whole number.

Calculating Relative Atomic Mass: Ar = [(% abundance of isotope 1 x mass of isotope 1) + (% abundance of isotope 2 x mass of isotope 2) + ...] / 100 Example 1: Chlorine has two isotopes: Chlorine-35 (³⁵Cl) with an abundance of 75.77% and Chlorine-37 (³⁷Cl) with an abundance of 24.23%. Calculate the relative atomic mass of chlorine. Ar (Cl) = [(75.77 x 35) + (24.23 x 37)] / 100 Ar (Cl) = (2651.95 + 896.51) / 100 Ar (Cl) = 3548.46 / 100 Ar (Cl) = 35.48 Example 2: Magnesium has three isotopes: Magnesium-24 (²⁴Mg) with 78.99% abundance, Magnesium-25 (²⁵Mg) with 10.00% abundance, and Magnesium-26 (²⁶Mg) with 11.01% abundance. Calculate the relative atomic mass of magnesium. Ar (Mg) = [(78.99 x 24) + (10.00 x 25) + (11.01 x 26)] / 100 Ar (Mg) = (1895.76 + 250 + 286.26) / 100 Ar (Mg) = 2432.02 / 100 Ar (Mg) = 24.32 2.5 The Periodic Table The periodic table organizes elements by increasing atomic number and arranges them in a way that reflects their recurring (periodic) 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 (electrons in the outermost shell) and therefore exhibit similar chemical behavior. Electron Configuration and the Periodic Table: The periodic table directly reflects the electronic structure of atoms. The group number (for main group elements) often corresponds to the number of valence electrons. The period number corresponds to the number of electron shells. 2.6 Periodic Trends Several properties of elements exhibit trends as you move across a period or down a group in the periodic table.

Metallic Character: Generally, metallic character decreases across a period (left to right) and increases down a group (top to bottom). Metals lose electrons to form positive ions (cations), while nonmetals gain electrons to form negative ions (anions). In South Africa, where mining is a large industry, understanding metallic characteristics helps identify potentially profitable metals.

Atomic Size (Atomic Radius): Atomic size generally decreases across a period (left to right) because the increasing nuclear charge pulls the electrons closer to the nucleus. Atomic size generally increases down a group (top to bottom) because more electron shells are added.

Ionization Energy: The energy required to remove an electron from an atom in the gaseous phase. Ionization energy generally increases across a period (left to right) because the increasing nuclear charge holds the electrons more tightly. Ionization energy generally decreases down a group (top to bottom) because the outermost electrons are further from the nucleus and therefore easier to remove.

Electronegativity: A measure of an atom's ability to attract electrons in a chemical bond.