Ions are atoms with unequal numbers of protons and electrons

An ion is a charged particle formed when an atom or group of atoms has lost or gained electrons. In ordinary chemical ion formation, protons do not move from one particle to another, so the nuclear charge stays the same. The electron number changes.

A cation has fewer electrons than protons, giving it an overall positive charge. Metal atoms often form cations by losing electrons. A sodium atom, for example, has 11 protons and 11 electrons, but $Na^+$ has 11 protons and 10 electrons, so its charge is $1^+$.

An anion has more electrons than protons, so it carries an overall negative charge. Non-metal atoms often form anions by gaining electrons. For example, a sulfur atom can gain two electrons to form $S^{2-}$, the sulfide ion.

A monatomic ion is an ion made from a single atom. The names of monatomic anions usually end in -ide: chloride, oxide, sulfide, nitride and so on. Cation names usually stay the same as the element name. If a metal can form more than one charge, Roman numerals are used in compound names, such as iron(II) and iron(III).

Predicting charge from electron configuration

An electron configuration is a notation that shows how electrons occupy shells, subshells and orbitals in an atom or ion. At this level, the useful rule is simple: main-group atoms tend to form ions by reaching a noble gas electron configuration, as long as they can do it by losing or gaining only a small number of electrons.

For cations, remove the outer-shell electrons. Sodium is $1s^{2}2s^{2}2p^{6}3s^{1}$, so losing the $3s$ electron gives $Na^+$ with $1s^{2}2s^{2}2p^{6}$. Calcium is $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}$, so losing two $4s$ electrons gives $Ca^{2+}$ with $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}$.

For anions, add electrons until the outer shell is complete. Chlorine is $1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}$, so gaining one electron gives $Cl^-$ with $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}$. Oxygen is $1s^{2}2s^{2}2p^{4}$, so gaining two electrons gives $O^{2-}$ with $1s^{2}2s^{2}2p^{6}$.

Isoelectronic species are atoms or ions that have the same electron configuration. $Na^+$ is isoelectronic with Ne; $Cl^-$ and $Ca^{2+}$ are isoelectronic with Ar. Do not say they are “the same particle”: they have different proton numbers and different charges.

The periodic table gives a quick prediction for many main-group ions:

Group 14 becomes awkward if you try to force the rule. Carbon would have to lose or gain four electrons to make $C^{4+}$ or $C^{4-}$, which is not usually favourable; carbon more commonly shares electrons in covalent compounds.

The octet rule is a model stating that many atoms form bonds so that their valence shell contains eight electrons. It works well for many main-group ions, but it is still a model, not a law. Hydrogen, for example, can lose one electron to form $H^+$ or gain one electron to form $H^-$, reaching the helium configuration rather than an octet.

Ion formation and redox language

Oxidation is a chemical process in which a species loses electrons. Cation formation is oxidation: $Na \to Na^+ + e^-$.

Reduction is a chemical process in which a species gains electrons. Anion formation is reduction: $Cl + e^- \to Cl^-.$

That is why forming an ionic compound from its elements counts as a redox reaction. In the formation of sodium chloride, sodium atoms lose electrons and chlorine molecules gain electrons:

$2Na(s) \to 2Na^+ + 2e^-$

$Cl_2(g) + 2e^- \to 2Cl^-$

The oxidation-state language comes later in the course, but it fits the same idea: sodium goes from $0$ to $+1$, while chlorine goes from $0$ to $-1$.

Transition elements can form ions with different charges

A transition element is an element that has an atom or common ion with a partially filled d subshell. Unlike most main-group metals, transition elements often form more than one cation. Iron, for example, commonly forms $Fe^{2+}$ and $Fe^{3+}$.

When first-row transition elements ionize, the $4s$ electrons are removed before the $3d$ electrons. So Fe, often written $[Ar]4s^{2}3d^{6}$, forms $Fe^{2+}$ as $[Ar]3d^{6}$. Further electron removal can occur because the $4s$ and $3d$ subshells are close in energy, which helps explain variable oxidation states.

Successive ionization energy data support this. For transition elements, several electrons can often be removed with comparatively gradual increases before a much larger jump appears when an inner-shell electron would have to be removed.

Link this to periodic position: main-group ion charges are strongly linked to group number because atoms gain or lose electrons to reach a noble gas configuration. Transition elements are harder to predict from group number alone because the $ns$ and $(n-1)d$ electrons are similar enough in energy for several cation charges to be possible.

What an ionic bond is

An ionic bond is an electrostatic attraction between oppositely charged ions in a compound. It isn’t a tiny stick joining one cation to one anion. The attraction works between ions across the whole structure, which matters a lot when lattices come up.

A chemical bond is a strong attractive interaction that holds atoms or ions together in a substance. In the ionic model, the positive ion is the cation and the negative ion is the anion.

Ionic character and electronegativity

Electronegativity, $\chi$, is a dimensionless measure of the tendency of an atom in a covalent bond to attract the shared pair of electrons, where $\chi$ is the electronegativity value on the Pauling scale and has no unit. It generally increases across a period and up a group; fluorine has the highest value.

The electronegativity difference,

is often used to estimate ionic character.

A large $\Delta\chi$ points to greater ionic character. As a working guideline, bonding is mainly ionic when $\Delta\chi$ is greater than about 1.8. Treat that as a model, though: bonding sits on a continuum, so “metal + non-metal” is a useful first guess, not a guarantee. Aluminium chloride, $AlCl_3$, is the classic warning sign because its properties are more covalent than the element types alone might suggest.

Naming binary ionic compounds

A binary ionic compound is an ionic compound containing ions of two elements only. Write the cation first, then the anion. For a monatomic anion, take the end of the element name and add -ide.

Examples:

The name doesn’t give the number of ions directly. “Magnesium fluoride” is $MgF_2$, not magnesium difluoride, because ionic compound names do not use covalent prefixes such as mono-, di- and tri-.

For transition metals with variable charges, the Roman numeral gives the oxidation state of the metal ion. In iron(III) oxide, iron is present as $Fe^{3+}$, so the formula is $Fe_2O_3$. In copper(II) chloride, copper is $Cu^{2+}$, so the formula is $CuCl_2$. Write the symbol charge as $2^+$ or $3^+$ after the number; write the oxidation state as $+2$ or $+3$ with the sign first. Keep those formats separate.

Polyatomic ions you must know

A polyatomic ion is an ion composed of two or more covalently bonded atoms with an overall charge. When writing ionic formulas, treat a polyatomic ion as one charged unit: brackets may be needed, but the atoms inside the ion must not be changed.

Required polyatomic ions, with formulas and overall charges.

Polyatomic anions such as nitrate, carbonate, sulfate and phosphate are conjugate bases of acids. Their stability links to the acid dissociation constant, $K_a$, where $K_a$ is the acid dissociation constant ($\mathrm{mol\ dm^{-3}}$ under the usual IB concentration convention). In general, a more stable conjugate base corresponds to a larger $K_a$ for its conjugate acid, because acid dissociation is more favourable.

For sulfate, later work on formal charge helps explain which Lewis structure is preferred. Formal charge is a bookkeeping method for assigning electrons in a Lewis structure; the preferred sulfate representation is the one that gives the most reasonable distribution of formal charges while respecting the bonding model being used.

Deducing formulas from ions

An ionic compound is electrically neutral overall. So the total positive charge must equal the total negative charge. That’s the whole game.

A reliable method is:

1. Write the correct cation and anion, with charges.
2. Choose the smallest whole-number ratio that gives zero net charge.
3. Use brackets if more than one polyatomic ion is needed.
4. Check the charges before you write the final formula.

For magnesium chloride, $Mg^{2+}$ needs two $Cl^-$ ions to balance the charge, so the formula is $MgCl_2$.

For aluminium oxide, $Al^{3+}$ and $O^{2-}$ balance in a 2:3 ratio, because $2 \times 3^+$ gives $6^+$ and $3 \times 2^-$ gives $6^-$. The formula is $Al_2O_3$.

For barium hydroxide, $Ba^{2+}$ needs two $OH^-$ ions. Hydroxide is polyatomic, and there is more than one of it, so the formula is $Ba(OH)_2$. Do not write $BaO_2H_2$; that destroys the ion identity.

For ammonium phosphate, $NH_4^+$ and $PO_4^{3-}$ balance as three ammonium ions for one phosphate ion, giving $(NH_4)_3PO_4$.

Interconverting names and formulas

From formula to name, identify the ions first, then name the cation followed by the anion:

From name to formula, don’t guess subscripts from the words. Use the ion charges:

Formation of an ionic compound from its elements is a redox reaction because a metal atom is oxidized to a cation while a non-metal atom is reduced to an anion. The ionic formula you write is the charge-balanced result of those electron-transfer processes.

Ionic compounds are lattices, not molecules

An ionic lattice is a three-dimensional repeating arrangement of cations and anions held together by electrostatic attractions. In a crystal of sodium chloride, there are no separate $NaCl$ molecules. The whole crystal is one continuous array of $Na^+$ and $Cl^-$ ions.

An empirical formula is a chemical formula showing the simplest whole-number ratio of particles in a substance. For ionic compounds, the formula gives the ratio of ions in the lattice rather than a separate molecular unit. So $NaCl$ shows a 1:1 ratio of $Na^+$ to $Cl^-$. $MgCl_2$ shows a 1:2 ratio of $Mg^{2+}$ to $Cl^-`.

Ionic bonding is non-directional, meaning each ion attracts oppositely charged ions around it in all directions. This is why ionic substances build extended lattices instead of small pairs of ions.

Lattice enthalpy: measuring lattice strength

Lattice enthalpy, $\Delta H^\circ_{\text{lattice}}$, is the standard enthalpy change when one mole of solid ionic lattice is separated into its gaseous ions, where $\Delta H^\circ_{\text{lattice}}$ is the lattice enthalpy ($\mathrm{kJ\ mol^{-1}}$). For a simple 1:1 salt:

$MX(s) \to M^+(g) + X^-(g)\qquad \Delta H^\circ_{\text{lattice}} > 0$

Energy has to be supplied to overcome the attractions between ions, so this process is endothermic. Some sources use lattice enthalpy for the reverse process: forming the lattice from gaseous ions. That value has a negative sign. Always check the definition being used; in IB data booklet work, lattice enthalpy is treated as the energy needed to separate the lattice into gaseous ions.

Lattice enthalpy increases when ionic attractions are stronger. The two main factors are:

• Ionic charge: higher charges give stronger attractions. $MgO$ has much stronger attractions than $NaCl$ because $Mg^{2+}$ and $O^{2-}$ have larger charge magnitudes than $Na^+$ and $Cl^-$.
• Ionic radius: smaller ions can get closer together, so attractions are stronger. $NaF$ has a larger lattice enthalpy than $KF$ because $Na^+$ is smaller than $K^+$ while the anion is the same.

Selected ionic lattices showing how shorter ion distance and larger charge product give higher lattice enthalpy.

One useful investigation is to collect lattice enthalpy and ionic radius data for a series such as the group 1 chlorides, put the values into a spreadsheet, and plot lattice enthalpy against cation radius. The expected pattern is a decrease in lattice enthalpy down the group as the cation radius increases. Calculating charge density shows the same idea from another angle: smaller ions with the same charge have higher charge density and stronger attractions.

Volatility and melting point

Volatility is the tendency of a substance to vaporize. Ionic compounds are usually non-volatile because many strong electrostatic attractions must be overcome before ions can separate into the gas phase. The same reasoning explains why ionic compounds usually have high melting and boiling points.

This links lattice enthalpy directly to properties. A higher lattice enthalpy usually points to stronger ionic bonding and therefore a higher melting point, as long as the bonding type is genuinely comparable.

Across period 3 metal chlorides, charge alone does not explain the trend in melting points. Lattice enthalpies matter, and so does the bonding continuum. Sodium chloride and magnesium chloride are largely ionic, so their strong lattices give relatively high melting points; aluminium chloride has more covalent character, so its melting point is much lower than a purely ionic model would predict. It’s a useful reminder that models overlap rather than sit in separate boxes.

Melting points of selected period 3 metal chlorides compared with bonding character.

Electrical conductivity

Electrical conductivity is the ability of a substance to allow charge to flow. A substance conducts electricity only if it contains charged particles that are free to move.

Solid ionic compounds do not conduct electricity. They do contain ions, but those ions are locked into lattice positions and can only vibrate. The charged particles are present, just not mobile.

Molten ionic compounds do conduct electricity because melting breaks the rigid lattice arrangement and lets the ions move. Aqueous ionic solutions also conduct if the compound dissolves, because the separated hydrated ions can move through the water.

The experimental evidence is simple to collect: compare conductivity readings for a solid salt, the molten salt, pure water and an aqueous salt solution. The solid gives little or no conductivity. The molten and aqueous samples give much higher conductivity because mobile ions are present.

Solubility in water and other solvents

Solubility is the extent to which a solute dissolves in a solvent to form a solution. Many ionic compounds dissolve in polar solvents such as water, but not in non-polar solvents such as hexane.

Water is polar: oxygen carries a partial negative charge and the hydrogen atoms carry partial positive charges. When an ionic crystal is placed in water, water molecules orientate themselves around the ions. The oxygen end points toward cations, and the hydrogen ends point toward anions. If these ion–water attractions are strong enough to overcome the attractions within the lattice, ions leave the lattice and become hydrated.

Not all ionic compounds dissolve in water. Calcium carbonate and silver chloride are familiar examples of low-solubility ionic compounds. In these cases, the lattice attractions are too strong relative to the attractions between ions and water molecules.

A precipitate is an insoluble solid that forms when two solutions are mixed and ions combine to make a low-solubility compound. Precipitation reactions give useful qualitative data for ionic compounds: if an initially clear mixture turns cloudy or forms a solid, an insoluble ionic product has formed.

When collecting solubility data, small-scale drop tests on a clear sheet over a dark background are enough to see precipitates clearly. Safety still matters: wear eye protection, treat unknown ionic solutions as irritants, and collect heavy-metal precipitates for proper disposal rather than washing them down the sink.

Typical experimental evidence for ionic properties includes: