Metals, Non-metals, Ions, and Periodic Trends

Ion Formation and Electron Configuration

Cation Formation (Metals)

• Definition: Metals form cations (positive ions) by losing their valence electrons. This results in more protons than electrons, leading to an overall positive charge.

• Goal: Atoms lose valence electrons to achieve the electron configuration of the nearest noble gas.

• Periodic Table Movement: When electrons are lost, one effectively moves backwards in the periodic table to find the noble gas configuration.

• **Generalizations for A Groups (Main Group Metals):

  • Group 1A (Alkali Metals): Always form a +1 charge. This includes hydrogen (even though it's a non-metal, it often has a +1 charge in compounds).

  • Group 2A (Alkaline Earth Metals): Have two valence electrons; they lose both to form a +2 charge.

  • Group 3A: Primarily aluminum (Al) is discussed, which forms a +3 charge (e.g., 13^{th} group for aluminum).

• Examples:

  • Magnesium (Mg), Beryllium (Be), Calcium (Ca) (all in Group 2A) will form ions with a +2 charge.

  • Aluminum (Al) (in Group 3A) will form an ion with a +3 charge.

Electron Configuration and Cation Formation Example: Sodium (Na)

• Neutral Sodium (Na): Atomic number is 11, meaning 11 protons and 11 electrons. Its electron configuration is 1s^22s^22p^63s^1. The highest shell is n=3, with 1 valence electron (3s^1).

• Sodium Ion (Na^+): Sodium loses its single valence electron (3s^1). The new electron configuration is 1s^22s^22p^6. This configuration is identical to Neon (Ne), which has 10 electrons.

• Mathematical Confirmation: If neutral Na has 11 protons and 11 electrons, then Na^+ has 11 protons and 10 electrons, resulting in a net charge of (+11) + (-10) = +1.

Electron Configuration and Cation Formation Example: Magnesium (Mg)

• Neutral Magnesium (Mg): Atomic number is 12, meaning 12 protons and 12 electrons. Its electron configuration is 1s^22s^22p^63s^2. It has 2 valence electrons (3s^2).

• Magnesium Ion (Mg^{2+}): Magnesium loses its two valence electrons (3s^2). The new electron configuration is 1s^22s^22p^6. This configuration also matches Neon (Ne), which has 10 electrons.

• Mathematical Confirmation: If neutral Mg has 12 protons and 12 electrons, then Mg^{2+} has 12 protons and 10 electrons, resulting in a net charge of (+12) + (-10) = +2.

• Key Reminder: Electrons are negatively charged. Losing electrons means there are relatively more protons, leading to a positive charge.

Anion Formation (Non-metals)

• Definition: Non-metals form anions (negative ions) by gaining electrons.

• Goal: Gaining electrons allows non-metals to achieve a stable octet arrangement or the electron configuration of the nearest noble gas.

• Periodic Table Movement: When electrons are gained, one moves forwards in the periodic table to find the noble gas configuration.

• **Generalizations for A Groups (Main Group Non-metals):

  • Group 7A (Halogens): Most readily gain 1 electron to form a -1 charge.

  • Group 6A: Gain 2 electrons to form a -2 charge.

  • Group 5A: Gain 3 electrons to form a -3 charge.

  • Group 4A: (e.g., Carbon). It's energetically unfavorable to either lose 4 or gain 4 electrons. These elements prefer to form covalent compounds by sharing electrons rather than forming ions.

• Why gain vs. lose? It is much easier for elements like nitrogen (Group 5A, 5 valence electrons) to gain 3 electrons to reach an octet (total 8) than to lose 5 electrons. Similarly, oxygen (Group 6A, 6 valence electrons) prefers to gain 2 rather than lose 6, and chlorine (Group 7A, 7 valence electrons) prefers to gain 1 rather than lose 7.

Electron Configuration and Anion Formation Example: Chlorine (Cl)

• Neutral Chlorine (Cl): Atomic number is 17, meaning 17 protons and 17 electrons. As a Group 7A element, it has 7 valence electrons.

• Chloride Ion (Cl^-): Chlorine gains 1 electron to achieve an octet. The new electron configuration is identical to Argon (Ar), which has 18 electrons.

• Mathematical Confirmation: If neutral Cl has 17 protons and 17 electrons, then Cl^- has 17 protons and 18 electrons, resulting in a net charge of (+17) + (-18) = -1.

• Key Reminder: Electrons are negatively charged. Gaining electrons means there are relatively more electrons, leading to a negative charge.

Transition Metals (D-block)

• Ion Formation: Transition metals are metals, so they form cations by losing electrons.

• Octet Rule: They do not follow the octet rule.

• Electron Loss: They lose both s valence and d valence electrons. The specific d electrons lost can lead to variable charges.

• Charge Determination: The charge of a transition metal ion must often be:

  • Given directly.

  • Distinguished by the name of the metal (using Roman numerals, e.g., Iron(II) vs. Iron(III)).

  • Back-calculated based on the charge of its counter-ion in a compound.

• Complexity: The specific charge depends on the environment and what the transition metal is bonded with.

Writing Electron Configurations for Ions

• General Approach:

  1. Determine the neutral atom's electron configuration.

  2. Calculate the total number of electrons in the ion: subtract electrons for cations (positive charge), add electrons for anions (negative charge).

  3. Write the electron configuration for this new total number of electrons.

• Movement on Periodic Table:

  • For a negatively charged ion (anion), one moves forward on the periodic table to find the noble gas it will resemble.

  • For a positively charged ion (cation), one moves backward on the periodic table to find the noble gas it will resemble.

• **Example: Selenium Ion (Se^{2-})

  • Neutral Selenium (Se): Atomic number 34 (34 protons, 34 electrons).

  • Selenium Ion (Se^{2-}): The -2 charge means it has gained 2 electrons. Total electrons: 34 + 2 = 36. The electron configuration of Se^{2-} will be identical to Krypton (Kr), which has 36 electrons.

• Other Ions Resembling Krypton (Kr):

  • Arsenic (As): Atomic number 33. If it forms As^{3-}, it gains 3 electrons: 33 + 3 = 36. Resembles Kr.

  • Bromine (Br): Atomic number 35. If it forms Br^-, it gains 1 electron: 35 + 1 = 36. Resembles Kr.

  • Rubidium (Rb): Atomic number 37. If it forms Rb^+, it loses 1 electron: 37 - 1 = 36. Resembles Kr.

  • Strontium (Sr): Atomic number 38. If it forms Sr^{2+}, it loses 2 electrons: 38 - 2 = 36. Resembles Kr.

• P-block Metals Below the Stair-step: These elements (e.g., Sn, Pb) can also exhibit variable charges due to the involvement of d-orbital electrons, making their behavior similar to transition metals in some aspects.

Determining Most Likely Charge from Valence Electron Configuration

• Given an outer shell valence electron configuration, one can determine the likely ion charge.

• Example 1: ns^2np^5

  • This indicates 2 + 5 = 7 valence electrons.

  • Elements with 7 valence electrons are in Group 7A (halogens).

  • Group 7A elements typically gain 1 electron to achieve an octet, forming a -1 charge.

• Example 2: ns^2

  • This indicates 2 valence electrons.

  • Elements with 2 valence electrons are in Group 2A (alkaline earth metals).

  • Group 2A elements typically lose 2 electrons, forming a +2 charge.

Electron Configuration for a Specific Transition Metal Ion: Titanium (Ti^{4+})

• Given: Ion electron configuration: 1s^22s^22p^63s^23p^6 (which is the configuration of Argon, [Ar]). Atomic number (Z) of Titanium is 22.

• Deduction:

  • Z=22 means 22 protons for Titanium.

  • The given ion configuration has 18 electrons (2+2+6+2+6=18).

  • Since it has 22 protons and 18 electrons, the charge is (+22) + (-18) = +4.

  • Therefore, the ion is Ti^{4+}.

• Neutral Titanium Configuration: [Ar]4s^23d^2.

• Electron Loss in Transition Metals: Transition metals like Titanium lose electrons from the highest principal energy level (n=4 in this case, so 4s^2) first, and then from the d subshell (3d^2). To form Ti^{4+}, Titanium loses its 2 4s electrons and its 2 3d electrons, resulting in the [Ar] configuration.

Why Ions Form: Energy Considerations

• Core Principle: Ion formation is driven by achieving a stable, lower-energy electronic state (like a noble gas configuration). This involves specific energy requirements or releases.

• Ionization Energy (IE):

  • Definition: The energy required to remove an electron from a single atom in the gaseous state.

  • Reaction Representation: Atom{(g)} + Energy ightarrow Cation{(g)} + e^- (Energy is an input, on the reactant side).

  • Interpretation: It measures how easily an atom loses an electron.

  • Favorability: Low ionization energy values are favorable, as less energy is needed to remove an electron.

• Electron Affinity (EA):

  • Definition: The energy released on adding an electron to a single atom in the gaseous state.

  • Reaction Representation: Atom{(g)} + e^- ightarrow Anion{(g)} + Energy (Energy is an output, on the product side).

  • Interpretation: It measures how easily an atom gains an electron.

  • Favorability: High (more negative, meaning greater release) electron affinity values are favorable, as more energy is released when forming an anion.

Periodic Trends in Ionization Energy (IE)

• General Trend:

  • Across a Period (left to right): Ionization energy generally increases. It becomes harder to remove an electron as effective nuclear charge increases and atomic size decreases.

  • Down a Group (top to bottom): Ionization energy generally decreases. As atomic size increases, valence electrons are further from the nucleus and experience more shielding from inner electrons, making them easier to remove.

• **Extremes for IE:

  • Lowest IE (Most Favorable Loss): Found in the bottom-left corner of the periodic table (e.g., Cesium, Cs). Alkali metals (Group 1A) have very low IE, explaining their high reactivity (e.g., with water).

  • Highest IE (Least Favorable Loss): Found in the top-right corner, specifically noble gases (Group 8A). They have extremely high IE because their octet is already stable, requiring a tremendous amount of energy to remove an electron.

• Shielding Effect: Down a group, a larger number of inner-shell electrons