Ionic Bonding

Ionic Bonding

Introducing Bonding

• Each element consists of one type of atom, often bonded to other atoms of the same element, forming molecules (e.g., chlorine Cl_2).

• Compounds like water (H_2O) have simple structures with few atoms bonded.

• Atoms form bonds to achieve a full outer electron shell, resembling stable noble gases.

• Noble gases are stable due to their full outer electron shells and are typically unreactive.

• Atoms of other elements have incomplete outer shells and are unstable.

• Bonding involves changes in the number of electrons in the outer shells of atoms.

Making Ions

• An ion is an atom or group of atoms with an electrical charge (positive or negative).

• Atoms have an equal number of protons and electrons, resulting in no overall charge.

• Atoms with incomplete outer electron shells are unstable.

• By gaining or losing electrons, atoms achieve full outer electron shells and become stable, resulting in an overall charge and the formation of ions.

• Atoms with nearly empty outer shells lose electrons to achieve a full outer shell and become positive ions.

• Atoms with nearly full outer shells gain electrons to achieve a full outer shell and become negative ions.

• Metal atoms lose one or more electrons to form positive ions. They are represented with a small “+” symbol and a number indicating the number of electrons lost (e.g., Na^+, Mg^{2+}).

  • Example: Sodium (Na) loses 1 electron to form Na^+. A sodium atom has 11 protons (+11) and 11 electrons (-11), resulting in a total charge of 0. A sodium ion has 11 protons (+11) and 10 electrons (-10), resulting in a total charge of +1. Electron configuration changes from 2.8.1 to [2.8]+

  • Example: Magnesium (Mg) loses 2 electrons to form Mg^{2+}. A magnesium atom has 12 protons (+12) and 12 electrons (-12), resulting in a total charge of 0. A magnesium ion has 12 protons (+12) and 10 electrons (-10), resulting in a total charge of +2. Electron configuration changes from 2.8.2 to [2.8]2+.

• Non-metal atoms gain one or more electrons to form negative ions. They are represented with a small “-” symbol and a number indicating the number of electrons gained.

  • Example: Fluorine (F) gains 1 electron to form F^-. A fluorine atom has 9 protons (+9) and 9 electrons (-9), resulting in a total charge of 0. A fluoride ion has 9 protons (+9) and 10 electrons (-10), resulting in a total charge of -1. Electron configuration changes from 2.7 to [2.8]-

  • Example: Sulfur (S) gains 2 electrons to form S^{2-}. A sulfur atom has 16 protons (+16) and 16 electrons (-16), resulting in a total charge of 0. A sulfide ion has 16 protons (+16) and 18 electrons (-18), resulting in a total charge of -2. Electron configuration changes from 2.8.6 to [2.8.8]2-.

• Calculating Ion Charges:

  • Calcium (Ca): Electron shells 2.8.8.2, loses 2 electrons, charge on ion: 2+

  • Hydrogen (H): Electron shell 1, loses 1 electron, charge on ion: 1+

  • Phosphorus (P): Electron shells 2.8.5, gains 3 electrons, charge on ion: 3-

  • Fluorine (F): Electron shells 2.7, gains 1 electron, charge on ion: 1-

  • Beryllium (Be): Electron shells 2.2, loses 2 electrons, charge on ion: 2+

Ionic Bonding

• Ionic compounds contain ions and are formed by a reaction between a metal and a non-metal.

• Metal and non-metal atoms have incomplete outer electron shells and are unstable.

• Electrons are transferred from metal atoms to non-metal atoms.

• The metal and non-metal atoms form ions with completely full outer shells and become stable.

• An ionic bond is the strong electrostatic attraction between oppositely charged ions.

• Dot-and-cross diagrams represent the formation of ionic compounds.

• Ionic bond is formed when electrons are transferred from one atom (or group of atoms) to another.

  • Example: Sodium (Na) + Chlorine (Cl) → Na^+ + Cl^-; Formula: NaCl

  • Sodium (Na) (2.8.1) + Chlorine (Cl) (2.8.7) → Na^+ (2.8) + Cl^- (2.8.8)

  • Example: Potassium (K) + Oxygen (O) → K^+ + O^{2-}; Formula: K_2O

  • Potassium (K) (2.8.8.1) + Oxygen (O) (2.6) → K^+ (2.8.8) + O^{2-} (2.8)

Formulae of Ionic Compounds

• In simple ionic compounds (e.g., NaCl, MgO), the metal loses the same number of electrons that the non-metal gains.

• Sodium chloride (NaCl): 1:1 ratio of Na^+ to Cl^-

• Magnesium oxide (MgO): 1:1 ratio of Mg^{2+} to O^{2-}

• Sodium Oxide (Na_2O):

  • Sodium (2.8.1) loses 1 electron to form Na^+.

  • Oxygen (2.6) gains 2 electrons to form O^{2-}.

  • Two sodium atoms are required for each oxygen atom; the ratio of sodium ions to oxide ions is 2:1.

• Magnesium Chloride (MgCl_2):

  • Magnesium (2.8.2) loses 2 electrons to form Mg^{2+}.

  • Chlorine (2.8.7) gains 1 electron to form Cl^-.

  • Two chlorine atoms are required for each magnesium atom; the ratio of magnesium ions to chloride ions is 1:2.

• Writing the formula of an ionic compound:

  1. Write down the symbol for each element (metal first).

  2. Calculate the charge for each type of ion.

  3. Balance the number of ions so that the positive and negative charges are balanced and equal zero. This gives the ratio of ions.

  4. Use the ratio to write down the formula of the ionic compound.

• Example: Aluminium Bromide (AlBr_3)

  • Aluminium(Al): Symbol Al, Ion charge +3

  • Bromide (Br): Symbol Br, Ion charge -1

  • 3 bromide ions are needed for each aluminium ion. Ratio 1:3.

• Example: Aluminium Oxide (Al2O3)

  • Aluminium: Symbol Al, Ion charge +3

  • Oxygen: Symbol O, Ion charge -2

  • 2 aluminium ions are needed for 3 oxide ions. Ratio 2:3.

• Examples of Ionic Compounds from Combinations of Metals and Non-metals:

  • Metals: Li, Ca, Na, Mg, Al, K

  • Non-metals: F, O, N, Br, S, Cl

  • LiF, CaF2, NaF, MgF2, AlF_3, KF

  • Li2O, CaO, Na2O, MgO, Al2O3, K_2O

  • Li3N, Ca3N2, Na3N, Mg3N2, AlN, K_3N

  • LiBr, CaBr2, NaBr, MgBr2, AlBr_3, KBr

  • Li2S, CaS, Na2S, MgS, Al2S3, K_2S

  • LiCl, CaCl2, NaCl, MgCl2, AlCl_3, KCl

• Lithium Nitrate (LiNO_3)

  • Lithium(Li): Symbol Li, Ion charge 1+

    • Nitrate (NO3): Symbol NO_3, Ion Charge 1-

    • 1 Lithium ion is needed for each nitrate ion, Ratio 1:1

Properties of Ionic Compounds

• Ionic compounds form an ionic lattice, a giant 3D structure.

• Ionic compounds form crystals when solid due to the structure of the ionic lattice.

• Ionic compounds have high melting and boiling points because ionic bonds are strong and require a lot of heat to break.

• Larger ionic charges produce stronger ionic bonds, requiring more heat to break (e.g., magnesium oxide vs. sodium chloride).

• Ionic compounds do not conduct electricity as solids because ions are bonded in the lattice and electrons are held within the bonds.

• Molten (liquid) ionic compounds can conduct electricity because ions are free to move and carry an electric current.

• Ionic compounds are brittle and shatter when hit.

• Summary of True or False Statements:

  1. True: The giant 3D structure of an ionic compound is called an ionic lattice.

  2. False: Only some ionic compounds form lattices and crystals when solid.

  3. False: Ionic bonds are quite weak and will break easily.

  4. True: Larger ionic charges create stronger bonds and the melting point is higher.

  5. False: Ionic compounds always conduct electricity, even when solid.

  6. True: Ionic compounds are brittle and shatter when hit.

summarised

Introducing Bonding

• Elements form molecules by bonding atoms.

• Compounds have simple, bonded structures (e.g., \text{H}_2\text{O}).

• Atoms bond to achieve full outer electron shells like noble gases.

• Noble gases are stable due to full outer shells.

• Other elements' atoms are unstable with incomplete outer shells.

• Bonding changes the number of outer-shell electrons.

Making Ions

• An ion is a charged atom or group of atoms.

• Atoms have equal protons and electrons (no charge).

• Atoms with incomplete outer shells are unstable.

• Gaining/losing electrons leads to full outer shells, creating stable ions.

• Nearly empty outer shells: lose electrons → positive ions.

• Nearly full outer shells: gain electrons → negative ions.

• Metals lose electrons → positive ions (e.g., \text{Na}^+, \text{Mg}^{2+}).

  • Example: Na loses 1 electron to form \text{Na}^+ (2.8.1 → 2.8).

  • Example: Mg loses 2 electrons to form \text{Mg}^{2+} (2.8.2 → 2.8).

• Non-metals gain electrons → negative ions.

  • Example: F gains 1 electron to form \text{F}^- (2.7 → 2.8).

  • Example: S gains 2 electrons to form \text{S}^{2-} (2.8.6 → 2.8.8).

Ionic Bonding

• Ionic compounds: ions formed by metal + non-metal reaction.

• Metal/non-metal atoms are unstable with incomplete outer shells.

• Electrons transfer from metal to non-metal atoms.

• Ions form with full outer shells, becoming stable.

• Ionic bond: strong electrostatic attraction between oppositely charged ions.

• Dot-and-cross diagrams show ionic compound formation.

• Example: Na + Cl → \text{Na}^+ + \text{Cl}^-; Formula: NaCl.

• Example: K + O → \text{K}^+ + \text{O}^{2-}; Formula: \text{K}_2\text{O}.

Formulae of Ionic Compounds

• Simple ionic compounds (e.g., NaCl, MgO): metal loses what non-metal gains.

• Sodium chloride (NaCl): 1:1 ratio of \text{Na}^+ to \text{Cl}^-.

• Magnesium oxide (MgO): 1:1 ratio of \text{Mg}^{2+} to \text{O}^{2-}.

• Sodium Oxide (\text{Na}_2\text{O}): 2:1 ratio of \text{Na}^+ to \text{O}^{2-}.

• Magnesium Chloride (\text{MgCl}_2): 1:2 ratio of \text{Mg}^{2+} to \text{Cl}^-.

• Writing ionic formula:

  1. Write element symbols (metal first).

  2. Calculate ion charges.

  3. Balance ions to achieve zero charge, determining the ratio.

  4. Write the formula using the ratio.

• Examples: \text{AlBr}3 (1:3 ratio), \text{Al}2\text{O}_3 (2:3 ratio).

• Examples of compounds with metals (Li, Ca, Na, Mg, Al, K) and non-metals (F, O, N, Br, S, Cl).

• Lithium Nitrate (\text{LiNO}_3): 1:1 ratio.

Properties of Ionic Compounds

• Form ionic lattice: giant 3D structure

• High Melting and Boiling Points: Because the attraction between the ions is super strong, it takes a lot of energy (heat) to break them apart.

• Don't Conduct Electricity (usually): As solids, the ions are stuck in place. But if you melt them, the ions can move around and carry an electric current.

• Brittle: If you hit them, they shatter because the ions get pushed out of alignment, and then the similarly charged ions repel each other.

In Simple Terms:

Ionic bonding is like a strong glue that holds together atoms that have gained or lost electrons. These atoms then form structures with high melting points that don't conduct electricity unless melted.