Lesson Notes By Weeks and Term v4 - SHS 3
MATTER AND ITS PROPERTIES
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MATTER AND ITS PROPERTIES
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Subject: Chemistry
Class: SHS 3
Term: 1st Term
Week: 10
Grade code: 1.1.1.LI.2
Strand code: 1
Sub-strand code: 1
Content standard code: 1.1.1.CS.2
Indicator code: 1.1.1.LI.2
Theme: PHYSICAL CHEMISTRY
Subtheme: MATTER AND ITS PROPERTIES
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This lesson introduces the fundamental concept of the Atomic Mass Unit (AMU). We will explore why scientists needed to create a special unit for measuring the mass of atoms, which are incredibly tiny. Just as we use cedis and pesewas for money instead of just pesewas for large amounts, or kilometres instead of centimetres to measure the distance from Accra to Kumasi, chemists use the AMU to work with the masses of atoms in a convenient way. We will learn how this unit is defined using a specific standard—the Carbon-12 isotope—and understand why the atomic masses you see on the periodic table are often not whole numbers.
2.1 The Problem: Why Can't We Use Grams?
Atoms are the basic building blocks of matter, but they are unimaginably small and light. For example, the mass of a single hydrogen atom is approximately: `0.000 000 000 000 000 000 000 001 67 grams` (or 1.67 x 10⁻²⁴ g).
Using such tiny numbers in everyday calculations is very difficult and impractical. To solve this, scientists created a new unit of mass specifically for atoms and molecules: the Atomic Mass Unit. 2.2 The Solution: A Relative Mass Scale
Instead of using the actual mass (absolute mass) in grams, scientists decided to measure the mass of every atom *relative* to a chosen standard atom. Analogy: Imagine you don't have a weighing scale, but you have a standard stone that you decide weighs "10 units". You can then find the mass of other objects by comparing them to your standard stone. A yam that is half as heavy would be "5 units", and a pineapple that is twice as heavy would be "20 units". This is a *relative* scale.