L6 - Ionic/Metallic Bonding

Electron Configuration Recap

• Building upon previous concepts gradually.

• Yesterday's lecture: Where do electrons sit? How to quantify them to understand the electronic structure around a nucleus?

• Bohr's model: Concept of multiple shells holds up (n = 1, n = 2, n = 3).

• These are the numbers before the orbitals or subshells.

• Within any particular shell, different subshells can be present.

  • n = 1 shell: only s subshell (maximum of 2 electrons).

  • Second shell: s (2s) and p (2p, holds 6) subshells.

  • Third shell: Includes d subshell.

• Comfortable describing shells and subshells within those shells.

  • s: Maximum capacity of 2 electrons.

  • d: Maximum capacity of 10 electrons.

  • n = 3 shell can technically hold 18 electrons.

• Filling of subshells: Order of filling based on energy levels.

  • Electrons prefer the lowest possible state (1s).

  • Occupy higher energy levels as we go.

• Predicting valence electrons:

  • Valence electrons: Outermost shell, the highest energy shell.

  • Even though n = 3 shell can hold 18 electrons, valence electrons are defined by the outermost shell.

  • Before reaching 3d, n = 4 shell may already be occupied, indicating a new shell is starting.

  • Number of valence electrons before that point would have been 8.

Bonding

• Three types of bonds: ionic, metallic, and covalent (mentioned but not covered today).

• Different types of bonding lead to different properties.

Metallic Bonding

• Explains why metals can be drawn into wires (ductile).

• Allows movement of atoms without breaking the bond.

Ionic Bonding

• Explains why salt (sodium chloride) dissolves in water.

• Consider the ionic components in solution.

• Translate understanding to the world around you based on atom-to-atom level interactions.

• Learning outcomes:

  • Cations and anions: Predict the charge using the periodic table and electronic configuration.

  • Discuss and articulate what ionic and metallic bonding are.

  • Based on elements, predict if they form an ionic or metallic bond.

  • Work out the charges of ions formed and how many electrons are passed between them.

Ions

• Definition: Imbalance of protons and electrons.

• Neutral Species: Equal number of protons and electrons (net zero charge).

• Ions: Formed by gaining or losing electrons, resulting in an overall charge.

• Positively charged species attract negatively charged species (crux of an ionic bond).

Notation

• Gaining an electron: Forms an anion (negatively charged).

• Losing electrons: Forms a cation (positively charged).

• Important: The number of protons doesn't change; the element remains the same.

Fluorine Example

• Fluorine (F) from the periodic table: Group 17, second row.

• Electronic configuration: 1s^2 2s^2 2p^5.

• Achieving stability: Elements try to find a stable form like the noble gases.

• Fluorine gains one electron.

• Notation: F- (charge in the top right).

• Mass number (top left), atomic number (bottom left), charge (top right).

Lithium Example

• Loses an electron.

• The electron is a product (right-hand side) unlike fluorine where it was a reactant (left-hand side).

• Cation losing an electron with the electron going somewhere else, which is presented on the right.

• Charge goes in the top right.

Equations

• Balanced equation including the electron is important to consider to balance the equation.

• Total ionic equation includes species gaining and losing electrons; electrons balance out on both sides.

• Electrons move from one atom to another; number of electrons the same on both sides.

Noble Gases

• Highly stable.

• Monoatomic: Present themselves as individual atoms.

• Require harsh conditions to react.

• Electronic configuration: Neon (Ne) - 1s^2 2s^2 2p^6.

• Valence shell: n = 2, eight valence electrons.

• Adding an additional electron would lose stability.

• Full shell has stability.

Why Ions Occur

• To achieve a full complement shell of electrons.

Sodium Example

• Shorthand electronic configuration: [Ne] 3s^1.

• Valence shell: n = 3.

• Losing one electron achieves the electron configuration of neon, achieving stability.

Quantum Numbers

• Principal quantum number: n value.

• Describes a location where that electron sits, like an address.

• Hierarchy: Street (principal quantum number) to house to bedroom to side of the room.

Periodic Table

• Predicting charges: Implications in electrochemistry.

• Electrochemistry: Electrical charge, reduction and oxidation, plating metals.

• Mixed-metal systems: Copper-coated wire, gold-plated jewelry, circuitry.

• Knowing what kinds of ions form tells how many electrons participate.

• Periodic table arrangement: s-block, d-block, p-block, f-block (lanthanides and actinides).

• Whether to gain or lose electrons: Red line divides; above and to the right are nonmetals (accept electrons, form anions); below and to the left are metals.

• Not a perfect rule, a continuous spectrum.

• Area around the dividing line has both metal and nonmetal characteristics.

• General Trend:

  • Above and right: Nonmetals.

  • Below and left: Metals.

Magnesium Example

• Magnesium (Element 12):

  • Electronic configuration: [Ne] 3s^2.

  • Two valence electrons.

  • Metal.

  • Loses two electrons, like neon.

  • Magnesium 2+.

• Common Ions: Consistent in Group 1, Group 2 (represented in blue in the slides).

Transition Metals

• The middle block has variability.

• Iron (Fe) can form Fe3+ and Fe2+.

Worksheet Examples

• Potassium (K): Group one, loses one electron, K+.

• Aluminum: Valence shell is 3s^2 3p^1. More likely to lose three electrons. Resulting electron configuration is 1s^2 2s^2 2p^6.

• Chlorine: Valence shell 3s^2 3p^5. Gain one. Results in six electrons on the 3p.

• Practice with equations: Electrons as products or reactants.

Isoelectronic Definition

• Isoelectronic: Same electron configuration.

  • F^- is isoelectronic with neon.

  • Li^+ is isoelectronic with helium.

Demos

• Reaction between sodium metal and chlorine gas.

• Observations: Record what you see.

• Ions formed.

• Formula of sodium chloride based on relative charge.

Sodium and Chlorine Reaction

• Stoppered flask with green chlorine gas.

• Sand: Absorbs heat.

• Water: Reacts with sodium, generating heat.

• Video Observations: Evolution of gas, different substance, condensation, color.

• Sodium plus chloride being formed is a ratio of Na^+ and Cl^-.

• One plus charge, one minus charge, one to one.

• Molecular formula: NaCl.

• Balanced equation: Represents the actual form of reactants at the beginning.

• Chlorine is diatomic, meaning two sodium atoms for every one chlorine molecule to form two NaCl.

• Count atoms on both sides.

Aluminum and Bromine

• Aluminum foil and bromine liquid.

• Bromine: Liquid at room temperature, vapor pressure creates bromine gas.

• Added aluminum foil.

• Sodium and Chlorine: The reaction was so quick and vigorous.

• Magnesium and Bromine: Starts about a minute later.

• Rate of reaction, how vigorously something might react, and what is required to get that reaction to start must be considered.

Ionic Bonding

• Relationship between cations and anions.

• Ionic bond: Electrons transferred from one atom to another.

• Requires a cation and an anion.

• Held together by electrostatic attraction.

• Variations based on the magnitude of charge.

• Size of ions: Varies.

• Crystal structure: Balances electrostatic attraction and physical space.

• Lattice-like structure: Alternating sodium and chloride.

Electron Transfer

• Magnesium: Two valence electrons.

• Oxygen: Six valence electrons.

• Two electrons are transferred, the same electrostatic attraction between Magnesium 2+ and Oxygen 2-.

• Two electrons from 2 Sodium are needed to complement 1 Oxygen since each sodium can give one electron to yield 2Na^+, O^{2-}.

• Coming back down to the lattice like structure, it's that balance. Want to minimize repulsion, occupation of physical space, but maximize electrostatic attraction.

Directionality

• Non-directional: If Na^+ is a sphere, its positive charge is equally distributed.

• Attracts in any direction.

• Lattice formation: Each sodium surrounded by chlorides; each chloride surrounded by sodiums.

Packing

• Cubic unit: Shows space occupied by sodiums and chlorides.

• Zinc sulfide: arrangement is quite different from sodium.

Metallic Bonding

• Deals with transition metals.

• Atoms have a tenuous grasp on their electrons.

• Delocalized sea of electrons.

• Electrons move freely around in the lattice structure.

• Because periodically there is an atom in its nucleus that that electron is attracted to, you don't have it locked to specific atoms.

• Explains why things are conductive (copper wire).

• Explains ductility: Stretching doesn't break metallic bonds.

• Sea of electrons allows movement for atoms within it.