Ionic Bonds and Ions Quiz

Ions

• Ions are charged particles formed when atoms lose or gain electrons.
• Gaining electrons creates a negative ion.
• Losing electrons creates a positive ion.

Molecular Ions

• Molecular ions are covalently bonded atoms that have lost or gained electrons.

Electrons Lost in Positive Ion Formation

• Electrons in the highest energy levels are lost when an atom becomes a positive ion.

Metals and Electron Gain/Loss

• Metals generally lose electrons to form positive ions.
• Non-metals generally gain electrons to form negative ions.

Elements That Don't Form Ions

• Beryllium, boron, carbon, and silicon don't readily form ions.
• This is because a significant amount of energy is needed to remove their outer shell electrons.

Types of Chemical Bonds

• Ionic bonding: Electrostatic attraction between positive and negative ions.
• Covalent bonding: Electrostatic attraction between a shared pair of electrons and the nuclei.
• Metallic bonding: Electrostatic attraction between positive metal ions and a sea of delocalised electrons.

Ionic Bonding Example

• Sodium chloride (NaCl), also known as salt, is an ionically bonded substance.

Ionic Bond Strength Factors

• Ionic radius and ionic charge determine the strength of an ionic bond.
• Stronger bonds correlate with smaller ions and higher charges.

Dot-and-Cross Diagram for MgO

• Mg²⁺: Has a 2,8 electron arrangement in the core, outer shell has lost 2 electrons.
• O^2-: Has 2,8 electron arrangement in the core, outer shell has gained 2 electrons.

Ionic Radius Trend Down a Group

• Ionic radius increases as you go down a group in the periodic table.
• This is because the ions have more electron shells, and the outermost electron experiences less pull from the positive nucleus.

Isoelectronic Ions Ionic Radius Trend

• The trend of ionic radius in a set of isoelectronic ions (same number of electrons) is determined by the number of protons.
• As the number of protons increases, the nuclear attraction increases, making the ionic radius decrease.

Physical Properties of Ionic Compounds

• High melting and boiling points due to strong electrostatic attractions between ions.
• Non-conductors of electricity in solid state as ions are fixed.
• Conductors in liquid or molten state as ions are mobile and can carry charge.
• Brittle due to repulsion between ions of same charge when the crystal lattice is disturbed.

Electrode Migration in Solution

• In a solution of CuCrO₄ with electrodes, Cu²⁺ ions migrate to the negative electrode, and CrO₄^2- ions migrate to the positive electrode.

Covalent Bond Definitions

• Single bond: Atoms bonded by one shared electron pair.
• Double bond: Atoms bonded by two shared electron pairs.
• Triple bond: Atoms bonded by three shared electron pairs.
• Dative covalent bond: Bond where both shared electrons are supplied by one atom only.

Covalent Structure Types

• Simple molecular lattice
• Giant covalent lattice

Simple Molecular Structures

• Atoms within a molecule are held by strong covalent bonds, and different molecules are held by weak intermolecular forces.

Low Melting/Boiling Points of Simple Molecular Structures

• Simple molecular structures have low melting and boiling points.
• This is because the weak intermolecular forces between molecules require very little energy to overcome.

Conductivity of Simple Molecular Structures

• Simple molecular structures are usually non-conductors of electricity because they have no free charged particles to move around.

Solvents for Simple Molecular Structures

• Simple molecular structures usually dissolve in non-polar solvents.

Giant Covalent Structures Examples

• Diamond, graphite, silicon dioxide (SiO₂)

Giant Covalent Structures Properties

• High melting and boiling points due to strong covalent bonds.
• Non-conductors, except graphite, which conducts because of delocalized electrons.
• Insoluble in various solvents due to the strong covalent bonds.

Graphite Electrical Conductivity

• Graphite conducts electricity due to delocalized electrons between the layers.

High Melting/Boiling Points of Giant Covalent Structures

• Giant covalent structures have high melting and boiling points because a large amount of energy is needed to break the strong covalent bonds within the molecules.
• Delocalized electrons are not freely movable in the solid structure for most giant covalent structures.

Diamond Structure

• Diamond has a 3-dimensional tetrahedral structure where each carbon atom is bonded to four other carbon atoms.

Molecular Shape Factors

• Molecular shape depends on the number of electron pairs and lone pairs in the outer shell of the central atom.

Molecular Shapes & Bond Angles

• BeCl₂: Linear, 180°
• BCl₃: Trigonal planar, 120°
• CH₄: Tetrahedral, 109.5°
• PCl₅: Trigonal bipyramidal, 90° and 120°
• SF₆: Octahedral, 90°
• NH₃: Pyramidal, 107°
• H₂O: Non-linear, 104.5°
• NH₄⁺: Tetrahedral, 109.5°

Lone Pairs and Bond Angles

• Each lone pair reduces the bond angle by approximately 2.5°.

Electronegativity Definition

• Electronegativity is the measure of an atom's ability to attract the pair of electrons in a covalent bond.

Electronegativity Scale

• Electronegativity is measured on the Pauling scale.

Electronegativity Trend

• Electronegativity increases from bottom-left to top-right in the periodic table, generally towards fluorine.

Non-Polar Bonds

• A non-polar bond has evenly distributed electrons.

Most Electronegative Element

• Fluorine.

Polar Bond Formation

• A polar bond forms when atoms involved in bonding have different electronegativities.

Water (H₂O) and Carbon Dioxide (CO₂) Polarity

• H₂O is polar because it's not symmetrical, unlike CO₂, which is a linear and symmetrical molecule. Thus, resulting in an overall dipole moment for water and no overall dipole moment for CO₂.

Intermolecular Force Definition

• An intermolecular force is an attractive force between neighboring molecules.

Intermolecular Force Types

• Hydrogen bonding
• Permanent dipole-dipole (permanent dipoles) interactions
• London dispersion forces (London forces)

London Forces Description

• London forces arise from momentary movements of electrons, resulting in instantaneous dipoles and thus creating induced dipoles in neighboring molecules, which attract.

London Forces and Molecular Size

• London forces are larger in larger molecules due to the increased number of electrons.

Noble Gas Boiling Points

• Boiling points in the noble gas group increase because the number of electrons increases, increasing the strength of the London forces.

Hydrogen Bonding Conditions

• Hydrogen bonding occurs when O-H, N-H, or F-H bonds are present, and there is a lone pair of electrons connected to oxygen, nitrogen, or fluorine atoms.
• The hydrogen nucleus must be left exposed for this interaction.

Ice Density

• Ice is less dense than liquid water.
• This is because the water molecules in ice are arranged in a specific ordered pattern, an open lattice with hydrogen bonds, which results in increased distance between molecules.
• The lattice structure collapses in liquid water resulting in molecules being closer together and more dense.

Water's High Melting/Boiling Point

• Water has a higher-than-expected melting and boiling point due to the stronger hydrogen bonds, which need more energy to break than other intermolecular forces.

Alkanes Intermolecular Forces

• Alkanes experience London forces (induced dipole-dipole) because their bonds are non-polar.

Alkane Boiling Point Increase and Chain Length

• As alkane chain length increases, the boiling point increases.
• This is due to the increased surface area, which leads to more induced dipole-dipole interactions, requiring more energy to overcome intermolecular forces.

Branched vs. Straight-Chain Alkanes

• Branched alkanes have lower boiling points than their straight-chain counterparts of similar chain length.
• This is because branched molecules have a smaller surface area, thus fewer induced dipole-dipole interactions.

Alkane Solubility in Water

• Alkanes are insoluble in water because water's stronger hydrogen bonds outweigh the weaker London forces in alkanes. Polarity differences prevent significant interaction.

Alcohols' Intermolecular Forces

• Alcohols exhibit hydrogen bonding due to the electronegativity difference in the O-H bond.

Alcohols Boiling Points and Hydrocarbons

• Alcohols generally have higher boiling points than hydrocarbons of similar chain lengths because hydrogen bonding.

Alcohol Solubility in Water

• Short-chain alcohols are soluble in water due to the ability of the hydroxyl group to hydrogen bond with water.
• Long-chain alcohols become less soluble as the non-polar hydrocarbon portion increases, exceeding the effect of the hydrogen bonding from the hydroxyl group.

Hydrogen Halide Boiling Points

• A general trend of increasing boiling point from HCl to HI is observed, due to strengthening of London dispersion forces as the number of electrons increases.
• The significant drop in boiling point from HF to HCl is due to the exceptionally strong hydrogen bonding within HF.