Lesson Notes By Weeks and Term v4 - SHS 3

Lesson Video

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.

Performance objectives

• Explain the periodic law.
• Describe properties like atomic size, ionization energy, and electronegativity.
• Analyze trends across periods and down groups.
• Explain some exceptions to these trends.

Lesson summary

Welcome, future scientists and engineers! Today, we are exploring one of the most powerful tools in chemistry: the Periodic Table. It's more than just a chart of elements; it's a map that reveals predictable patterns in the behaviour of matter. Understanding these patterns, known as periodicity, allows us to predict how an element will react, what kind of materials it will form, and how it behaves in nature without having to memorise facts for all 118 elements. In Ghana, this knowledge is vital. It helps us understand why gold is unreactive and perfect for jewellery, why chlorine is used to treat our water from the Weija Dam, and what nutrients our cocoa plants need from the soil.

Lesson notes

This section breaks down the core ideas you need to master. We will build our understanding step-by-step. 2.1 The Modern Periodic Law The foundation of our topic is the Periodic Law. Statement: The Modern Periodic Law states that the physical and chemical properties of the elements are a periodic function of their atomic number. Explanation: This means if you arrange the elements in order of increasing atomic number (the number of protons), elements with similar properties will appear at regular intervals or "periods". This is why all elements in Group 1 (like Lithium, Sodium, Potassium) are highly reactive metals, and all elements in Group 18 (like Helium, Neon, Argon) are very unreactive noble gases. This regularity is a direct result of the elements' electron configurations, which repeat in a predictable pattern. 2.2 The "Why" Behind the Trends: Core Influencing Factors Almost all periodic trends can be explained by the interplay of two main factors. Understanding these is the key to mastering periodicity. Effective Nuclear Charge (Zeff): Meaning: This is the net positive charge experienced by an outermost (valence) electron. It's the "pull" a valence electron actually *feels* from the nucleus. Why it's not the full nuclear charge: The inner-shell electrons "shield" the valence electrons from the full pull of the positive nucleus. Trend: Across a Period (Left to Right): Zeff *increases*. Why? Because you are adding protons to the nucleus (increasing the positive charge), but the new electrons are being added to the *same principal energy level*. These electrons do not shield each other effectively. So, the pull from the nucleus on the outer electrons gets stronger. Down a Group (Top to Bottom): Zeff *remains relatively constant or increases slightly*. While the number of protons increases significantly, the number of inner shielding electrons also increases, largely cancelling out the effect of the added protons. The dominant factor down a group is the increasing number of shells. Shielding Effect and Principal Quantum Number (n): Meaning: The shielding effect (or screening effect) is the reduction in the nuclear charge experienced by a valence electron due to the presence of inner-shell electrons. Principal Quantum Number (n): This represents the main energy level or "shell" of an electron. A higher 'n' means the electron is, on average, further from the nucleus. Trend: Across a Period: The principal quantum number (n) of the outermost shell remains the *same*. Shielding is relatively constant. Down a Group: The principal quantum number (n) *increases*. A new energy shell is added for each new period. This increases the distance between the nucleus and the valence electrons and also increases the number of inner shells, leading to a much stronger shielding effect.

Summary of a tug-of-war: Think of Zeff as the strength of the "nucleus team" pulling on an electron. Think of shielding and distance (n) as factors helping the "electron team" resist that pull.

2.3 The Periodic Properties

Let's examine each property using our understanding of Zeff and Shielding. A. Atomic Radius Meaning: The atomic radius is typically defined as half the distance between the nuclei of two identical atoms bonded together. It's a measure of the size of an atom. Factors: Zeff and number of electron shells (n). Variation in the Periodic Table: Across a Period (Left to Right): Atomic radius *decreases*. Reasoning: As you move across a period, the number of protons increases, leading to a stronger Effective Nuclear Charge (Zeff). This increased pull draws the electron cloud of the *same energy shell* closer to the nucleus, making the atom smaller. Down a Group (Top to Bottom): Atomic radius *increases*. Reasoning: As you move down a group, a new principal energy level (shell) is added for each element. This new shell is further from the nucleus. This increase in distance and the added shielding from inner electrons far outweighs the increase in nuclear charge. Therefore, the atoms get significantly larger. B. Ionic Radius Meaning: The ionic radius is the radius of an atom's ion. Key Idea: Comparing an ion to its parent atom is crucial. Cations (Positive Ions): Cations are *smaller* than their parent atoms. When an atom loses electrons to form a cation, it often loses its entire valence shell. Furthermore, the remaining electrons are pulled more tightly by the unchanged nuclear charge. (e.g., Na⁺ is much smaller than Na). Anions (Negative Ions): Anions are *larger* than their parent atoms. When an atom gains electrons, the extra electron-electron repulsion causes the electron cloud to expand. The nuclear charge remains the same, but it now has to hold onto more electrons, so its pull on each is weaker. (e.g., Cl⁻ is larger than Cl). Variation in the Periodic Table: The trends for ions follow the same general logic as atomic radii, but you must be careful when comparing. For an *isoelectronic series* (ions with the same number of electrons, e.g., O²⁻, F⁻, Na⁺, Mg²⁺), the radius decreases as the nuclear charge (atomic number) increases. C. First Ionization Energy (IE₁) Meaning: The first ionization energy is the minimum energy required to remove one mole of the most loosely held electrons from one mole of gaseous atoms to form one mole of gaseous ions with a +1 charge. In simple terms: How much energy it takes to "steal" an electron from a neutral atom. `X(g) + Energy → X⁺(g) + e⁻` Factors: High Zeff, small atomic radius, and stable electron configurations make it *harder* to remove an electron (i.e., higher IE). Variation in the Periodic Table: Across a Period (Left to Right): Ionization energy generally *increases*. Reasoning: Zeff increases and atomic radius decreases. The valence electrons are held more tightly, so more energy is needed to remove one. Down a Group (Top to Bottom): Ionization energy *decreases*. Reasoning: Atomic radius increases and shielding increases. The valence electron is further from the nucleus and feels a weaker pull, making it easier to remove. D. Electron Affinity (EA) Meaning: The energy change that occurs when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions with a -1 charge. In simple terms: How much an atom "wants" to gain an electron. `X(g) + e⁻ → X⁻(g) + Energy` (For most elements, this is exothermic, so EA values are often negative). A more negative value means a stronger affinity. Variation in the Periodic Table: Across a Period (Left to Right): Electron affinity generally becomes *more negative* (a stronger attraction for an electron). Reasoning: Increasing Zeff means the nucleus can more effectively attract an incoming electron. Down a Group (Top to Bottom): Electron affinity generally becomes *less negative* (a weaker attraction). Reasoning: The incoming electron is entering a shell further from the nucleus, where the attraction is weaker. E. Electronegativity Meaning: A measure of the ability of an atom in a chemical bond to attract shared electrons towards itself. Factors: It's a calculated value, but it follows the same logic as IE and EA. A small atom with a high Zeff will be highly electronegative. Variation in the Periodic Table: Across a Period (Left to Right): Electronegativity *increases*. (Fluorine is the most electronegative element). Down a Group (Top to Bottom): Electronegativity *decreases*.

Evaluation guide

• Explain how periodic properties change with atomic number and principal
• quantum number.
• Talk for L earning:
• State and explain the periodic law.