Ionic Bonds and Ions Quiz

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Questions and Answers

What are ions?

Charged particles that are formed when an atom loses or gains electrons

What is the charge of the ion when electrons are gained?

Negative

What are molecular ions?

Covalently bonded atoms that lose or gain electrons

Which electrons are lost when an atom becomes a positive ion?

Electrons in the highest energy levels Signup and view all the answers

Do metals usually gain or lose electrons?

Lose electrons Signup and view all the answers

Which are the 4 elements that don’t tend to form ions and why?

The elements are beryllium, boron, carbon and silicon. Requires a lot of energy to transfer outer shell electrons. Signup and view all the answers

What are the 3 main types of chemical bonds?

Ionic (A), Covalent (B), Metallic (D) Signup and view all the answers

Define ionic bonding

The electrostatic attraction between positive and negative ions Signup and view all the answers

Give an example of an ionically bonded substance

NaCl (Sodium Chloride - salt) Signup and view all the answers

What determines the strength of an ionic bond?

<ul> <li>Ionic radius and ionic charge</li> <li>Ionic bonding is stronger and the melting points higher when the ions are smaller and/ or have higher charges.</li> </ul> Signup and view all the answers

Draw the dot and cross diagram to show ionic bonding in MgO

[Mg] 2+ [2, 8] 2+ [O] 2- [2, 8] 2- Signup and view all the answers

Explain the trend in ionic radius down a group

Ionic radii increases going down the group. This is because down the group the ions have more shells of electrons and thus the outermost electron experience less pull from positive nucleus. Signup and view all the answers

Explain the trend in ionic radius for this set of isoelectronic ions, e.g. N3- to Al3+

There are increasing numbers of protons from N to F and then Na to Al but the same number of electrons. Therefore nuclear attraction between the outermost electrons and nucleus increases and ions get smaller. Signup and view all the answers

What are the physical properties of ionic compounds?

High melting points (A), Conductor of electricity when in solution or molten (B), Brittle (C), Non-conductor of electricity when solid (D) Signup and view all the answers

In a solution of CuCrO4 with connected electrodes, which electrode will the 2 ions migrate to?

Cu2+ - migrates to negative electrode CrO4 2- - migrates to positive electrode Signup and view all the answers

Study Notes

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.

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Description

Test your knowledge on ions, ionic bonding, and the characteristics of ionic compounds. This quiz covers key concepts such as the behavior of electrons in ions, types of chemical bonds, and the unique properties of ionic substances. Dive into the fascinating world of chemistry and check your understanding!