Chem pp2

Electronic Configuration and the Periodic Table:​

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● Periods (Rows): Elements in the same period have the same number of occupied electron shells. This means as you move across a period, the number of protons and electrons increases, but the number of electron shells remains constant.​

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● Groups (Columns): Elements in the same group have similar chemical properties due to having the same number of valence electrons (electrons in the outermost shell). This similarity in electron arrangement leads to the repeating patterns, or periodicity, of element properties.

Trends in Atomic Structure and Properties:​

1. Atomic Radius: Atomic radius decreases across a period and increases down a group. This is due to the interplay between the number of electron shells and the effective nuclear charge. ​

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* Across a Period: As the number of protons increases, the positive charge of the nucleus pulls the electrons closer, resulting in a smaller atomic radius. ​

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* Down a Group: As you move down a group, the number of electron shells increases, meaning the valence electrons are further from the nucleus, leading to a larger atomic radius.​ This trend results in a decrease in ionization energy across a period, as the increased nuclear charge makes it easier to remove an outer electron, while ionization energy tends to increase down a group due to the greater distance of the valence electrons from the nucleus.

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Electronegativity:​

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Electronegativity is a measure of an atom's ability to attract electrons towards itself in a chemical bond. It increases across a period and decreases down a group. ​

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* Across a Period: As the effective nuclear charge increases, the stronger attraction for electrons leads to higher electronegativity. ​

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* Down a Group: Increasing atomic radius means the valence electrons are farther from the nucleus, resulting in a weaker attraction and lower electronegativity

First Ionisation Energy: As we move across a period, first ionisation energy generally increases due to the increasing effective nuclear charge, which holds the electrons more tightly. Conversely, down a group, first ionisation energy decreases because the outermost electrons are further from the nucleus and experience more shielding from inner electron shells. ​

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The first ionisation energy is the energy required to remove the most loosely bound electron from a neutral atom. ​

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* Across a Period: The increasing effective nuclear charge makes it more difficult to remove an electron, hence the higher ionisation energy. ​

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* Down a Group: With increasing atomic radius, the outermost electron is less tightly held by the nucleus, requiring less energy to remove, hence the lower ionisation energy.

​​Metallic and Non-Metallic Character: ​

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* Metallic Character: Generally decreases across a period and increases down a group. This is because metals tend to lose electrons easily, and this tendency is related to lower ionisation energies and electronegativities. ​

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* Non-Metallic Character: Generally increases across a period and decreases down a group. Non-metals tend to gain electrons, a trend associated with higher electronegativities and ionisation energies ​

Reactivity: Reactivity refers to the tendency of an element to undergo chemical reactions. Reactivity trends vary depending on whether you are considering metals or non-metals.​

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Metals: Reactivity increases down a group (as it becomes easier to lose electrons) and generally decreases across a period.​

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Non-Metals: Reactivity decreases down a group (as it becomes harder to gain electrons) and generally increases across a period. The exception to this is the noble gases, which are very unreactive due to having a full outer shell of electrons.​

Trends in Reactivity:

• Alkali metals (Group 1) are highly reactive, especially with water, and their reactivity increases down the group.

• Halogens (Group 17) are also very reactive, with reactivity increasing up the group as they seek to gain electrons to achieve a full outer shell.

The Bohr Model: ​

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This model is a simplified representation of electron configuration, where electrons are arranged in energy levels called shells, which are numbered 1, 2, 3, and so on, starting from the shell closest to the nucleus. ​

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The maximum number of electrons each shell can hold is given by the formula 2n², where 'n' is the shell number. For example:​

● Shell 1 (n = 1): Can hold a maximum of 2 electrons (2 x 1² = 2)​

● Shell 2 (n = 2): Can hold a maximum of 8 electrons (2 x 2² = 8)

The Schrödinger Model (Quantum Mechanical Model): ​

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This model provides a more detailed and accurate description of electron configuration. It introduces the concept of subshells within each shell. Subshells are designated by the letters s, p, d, and f, and each subshell has a specific number of orbitals:​

● s subshells: 1 orbital​

● p subshells: 3 orbitals​

● d subshells: 5 orbitals​

● f subshells: 7 orbitals​

Each orbital can hold a maximum of two electrons.

The energy levels of subshells don't always follow a simple numerical order. For example, the 4s subshell is filled before the 3d subshell because it has a lower energy level.

Shell Electron Configuration: This notation simply lists the number of electrons in each shell, separated by commas. For example, oxygen (8 electrons) has a shell configuration of 2, 6.​

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Subshell Electron Configuration: This notation uses the shell number, subshell letter, and a superscript to indicate the number of electrons in each subshell. For example:​

● Carbon: 1s²2s²2p²​

● Sodium: 1s²2s²2p⁶3s¹​

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Condensed Subshell Electron Configuration: This notation uses the symbol of the previous noble gas in brackets to represent the filled inner shells, followed by the configuration of the outer shells. For example:​

● Aluminium: [Ne]3s²3p¹​

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