Chapter 8: Chemical Reactivity

Metals: Na, Mg, Ca, Fe and Cu.

Non-metals: H, C, N, O, S, Cl and Br.

Common metalloids: B, Si, Ge, As, Sb and Te.

Scientists sometimes classify borderline elements differently.

Air is about 78% nitrogen, 21% oxygen and 0.93% argon by volume. Oxygen, silicon and aluminium are the most common elements by mass in Earth’s crust. Hydrogen and helium are the most common in the universe.

Look out for exceptions. Sodium is a soft metal. Iodine is a shiny non-metal. Graphite is a non-metal that conducts electricity.

An electropositive element tends to lose electrons. A metal atom becomes a positively charged cation on losing electron(s).

Many metals hold their outer electrons fairly weakly. Only a small amount of energy is needed to remove them.

Group 1 metals have 1 outer electron, ns¹, and usually form M⁺. Group 2 metals have 2 outer electrons, ns², and usually form M²⁺.

Remember: Group 1 metals are alkali metals. Group 2 metals are alkaline earth metals.

Down Group 1, atoms get larger and have more inner shells. Inner electrons reduce the nucleus’s pull on the outer electron. This is shielding. The outer electron is usually easier to remove, so reactivity generally increases.

For common reactions with water: Li < Na < K < Rb < Cs. Francium is too rare and radioactive for ordinary comparisons.

Reactivity usually increases from Be to Ba. Removing electrons becomes easier. Surface coatings and reaction conditions can also affect the result.

First ionization energy is the energy needed to remove one electron from each atom in one mole of isolated gaseous atoms: M(g) → M⁺(g) + e⁻.

The first ionization energy is about 590 kJ/mol for calcium and 503 kJ/mol for barium.

2Mg + O₂ → 2MgO | Ca + Cl₂ → CaCl₂ | 2Na + Cl₂ → 2NaCl

A protective surface layer slows Mg’s reaction with cold water. It reacts more readily with hot water or steam. Ca reacts noticeably with cold water.

Some Group 1 metals form carbides. A carbide contains carbon and a less electronegative element. Lithium forms Li₂C₂.

A Group 1 atom needs to lose only 1 outer electron. Removing it usually takes little energy. This helps explain its high reactivity. Compare specific reactions carefully.

Heating can raise electrons to higher energy levels. When they fall to lower levels, they can give out visible light. Different elements can give different flame colours.

Be and Mg do not give reliable colours in ordinary school flame tests. Burning Mg gives bright white light, but that is different from testing Mg²⁺ ions in a flame.

Flame tests need trained staff and proper laboratory equipment. Never taste or inhale salts.

Na has atomic number 11. It is in Group 1, Period 3. Its electron arrangement is 2,8,1. Sodium is soft and silvery. It melts near 98 °C and reacts very strongly with water.

Water: 2Na + 2H₂O → 2NaOH + H₂↑.

Bromine: 2Na + Br₂ → 2NaBr.

Dilute sulfuric acid: 2Na + H₂SO₄ → Na₂SO₄ + H₂↑ (theoretical reaction; not a safe demonstration).

Oxygen: 4Na + O₂ → 2Na₂O under oxygen-limited conditions; sodium can also form sodium peroxide, 2Na + O₂ → Na₂O₂, depending on conditions.

Sodium reacts quickly with water and may catch fire. Oil keeps air and moisture away.

Sodium metal is used in some heat-transfer systems and as a reducing agent. Sodium vapour lamps give yellow light. Their light comes from the vapour.

Magnesium: atomic number 12, Group 2, Period 3. Electron arrangement 2,8,2; usually forms Mg²⁺.

Mg is a light, silvery metal. It melts near 650 °C. Burning Mg gives bright white light:

2Mg + O₂ → 2MgO.

Mg reacts very slowly with cold water. Hot water can give Mg(OH)₂ and H₂. With steam:

Mg + H₂O(g) → MgO + H₂.

Light magnesium alloys are used in vehicles, aircraft and portable devices.

Mg(OH)₂ is used in some antacid medicines.

Speciality flares make use of its bright white light.

Magnesium can corrode instead of steel. It loses electrons and supplies them to the steel. In this role, Mg is a sacrificial anode.

Do not use water on burning magnesium. A trained person must choose the correct method for a metal fire.

Calcium: atomic number 20, Group 2, Period 4; electronic arrangement 2,8,8,2. It forms Ca²⁺ and reacts with water: Ca + 2H₂O → Ca(OH)₂ + H₂.

In the past, strongly heating CaO in an oxygen–hydrogen flame gave a bright light. This “limelight” was used on stages.

Gypsum is CaSO₄·2H₂O. Heating removes some of its water of crystallization to make plaster of Paris, CaSO₄·½H₂O. Water of crystallization is water present in a fixed ratio in the crystal.

Hardness is resistance to scratching or denting. The Mohs scale compares how easily one material scratches another. It does not measure resistance to being pulled apart.

Strong metallic bonding can help make a metal hard and give it a high melting point. The crystal arrangement, defects and alloy composition also affect hardness.

Noble metals have low tendency to react or corrode under many ordinary conditions. Examples include Au, Ag, Pt, Pd, Rh and Ir. 'Noble' does not mean they can never react.

Silver is shiny. It conducts electricity very well and reflects light strongly. Uses include contacts, mirrors, coatings and jewellery. AgNO₃ is called lunar caustic. AgBr and AgI were widely used in photography.

Gold can form thin sheets and wires. It conducts heat and electricity and resists corrosion. Uses include jewellery, electrical contacts and thin reflective coatings. Some special equipment uses gold coatings to control radiant heat.

Platinum resists corrosion. It also speeds up many reactions as a catalyst. Uses include catalytic converters, chemical industries and some electrical and medical devices.

Gold and platinum prices change. Density alone does not decide which costs more.

Most non-metals conduct heat and electricity poorly. Solid non-metals often break rather than bend. Many form covalent bonds.

Gases at room temperature: H₂, N₂, O₂, F₂, Cl₂ and the noble gases.

Liquid at room temperature: Br₂ (bromine).

Solids at room temperature: C, P, S, Se and I₂.

There are exceptions. Graphite conducts electricity. Iodine is shiny. Diamond is extremely hard.

Group 18 contains He, Ne, Ar, Kr, Xe, Rn and Og. The familiar noble gases have full outer shells and usually react very little.

Group 17 contains F, Cl, Br, I, At and Ts. F₂ is pale-yellow gas. Cl₂ is yellow-green gas. Br₂ is red-brown liquid. I₂ is a dark solid that can give violet vapour.

A non-metal can gain electron(s) to form a negative ion: Cl + e⁻ → Cl⁻. Oxygen can form O²⁻ in suitable ionic compounds.

Electron affinity describes the energy change when a gaseous atom gains an electron. Electronegativity is the ability of a bonded atom to attract a shared electron pair. They are related but not the same.

Electronegativity usually increases across a period and decreases down a group. Fluorine has the highest value.

Many non-metal oxides form acids in water.

CO₂ + H₂O ⇌ H₂CO₃.

SO₂ + H₂O ⇌ H₂SO₃.

Further oxidation of SO₂ in air can produce sulfuric acid and contribute to acid rain.

Some non-metal oxides are neutral rather than acidic. Examples are CO, NO and N₂O.

Halogens often gain electrons in redox reactions. They therefore often act as oxidizing agents.

Oxidizing ability of common halogens decreases down Group 17: F₂ > Cl₂ > Br₂ > I₂.

A more reactive halogen can replace a less reactive halogen from its halide solution.

Hydrogen combines with a halogen:

H₂ + X₂ → 2HX.

In dilute water solutions, acid strength usually increases:

HF < HCl < HBr < HI.

HF is a weak acid in water, while HCl, HBr and HI are strong acids. All hydrogen halides are hazardous; hydrofluoric acid is especially dangerous.

Fluoride helps tooth enamel resist acids. It helps minerals return to the enamel and form a surface that acids attack less easily. Enamel is mainly hydroxyapatite, a calcium phosphate mineral.

Qualitative analysis identifies the ions present using colours, precipitates or gases. A precipitate is an insoluble solid formed in a solution. Sometimes an extra test is needed to confirm the ion.

Zn²⁺ + 2OH⁻ → Zn(OH)₂(s). In excess OH⁻: Zn(OH)₂ + 2OH⁻ → [Zn(OH)₄]²⁻.

Mg²⁺ + 2OH⁻ → Mg(OH)₂(s).

NH₄⁺ + OH⁻ → NH₃ + H₂O (when warmed under suitable lab conditions).

Ammonia irritates the body, and some barium salts are toxic. These tests need a trained teacher and a suitable laboratory.

White AgCl dissolves in dilute aqueous ammonia. Yellow AgI does not dissolve readily. This difference helps tell the ions apart.

Ca(OH)₂ + CO₂ → CaCO₃(s) + H₂O. Insoluble CaCO₃ makes limewater milky.

BaSO₄ dissolves very little. It stays as a white solid under suitable acidic test conditions.

Adding acid to nitrite can release toxic gases. Such tests need a properly equipped laboratory and trained supervision.

Metals: Ca, Sr, K, Ba. Non-metals: C, S, P. Metalloids: Sb, Si, Ge.

Group 1 metals with one valence electron: Li, Na, K, Rb, Cs and Fr. Hydrogen is in Group 1 but is not an alkali metal.

Group 2 metals with two valence electrons: Be, Mg, Ca, Sr, Ba and Ra.

Group 17 non-metals: F, Cl, Br, I, At; tennessine (Ts) is also in Group 17, although its properties are not well established.

Boron (B), silicon (Si) and germanium (Ge).

A cation is a positive ion formed by loss of electron(s), e.g. Na⁺. An anion is a negative ion formed by gain of electron(s), e.g. Cl⁻.

Their high reactivity makes them react readily with water, oxygen and other substances to form compounds.

They resist oxidation and corrosion in many ordinary environments due to their relatively low chemical reactivity.

Gold (Au), silver (Ag), platinum (Pt), palladium (Pd), rhodium (Rh) and iridium (Ir).

Non-metal atoms can attract electrons strongly; in covalent bonds their electronegativity measures attraction for the shared electrons. Some also gain electrons to become anions.

Silver is lustrous, reflects visible light well, and is an excellent electrical conductor. Uses: mirrors, electrical contacts, jewellery, coatings and traditional light-sensitive photographic compounds (AgBr, AgI).

Gold is highly resistant to corrosion and is malleable, ductile and electrically conductive. It is used in jewellery, electronic contacts, medical devices and thin reflective coatings.

In school-level qualitative analysis, Ag⁺ gives a white precipitate with Cl⁻ (AgCl) and a yellow precipitate with I⁻ (AgI). AgCl dissolves in dilute ammonia, while AgI does not. Observations and reagents should be interpreted in a properly supervised laboratory.

Most metals readily lose outer electrons, especially those with low ionization energy. Na → Na⁺ + e⁻, Mg → Mg²⁺ + 2e⁻. The resulting positively charged ions are cations.

Magnesium is in Group 2 and Period 3. It forms Mg²⁺ and burns with a bright white flame. Its alloys are lightweight, Mg(OH)₂ is used in antacids, and it is used as a sacrificial anode for corrosion protection.

Sodium is in Group 1 and Period 3. It forms Na⁺ and reacts vigorously with water, so it is stored away from moisture. Its compounds include NaCl (food and industry), Na₂CO₃ (glass and softening water) and NaHCO₃ (baking).

In water: HF < HCl < HBr < HI. HF is weak; HCl, HBr and HI are strong acids. The H–X bond weakens down the group, making proton release easier.

It is the ability of an atom in a bond to attract shared electrons. It generally increases across periods and decreases down groups. Fluorine has the highest electronegativity.

Sodium is a soft Group 1 metal, Mohs hardness about 0.5, low density and low melting point (~98°C). Iron is a harder Group 8 transition metal, Mohs hardness about 4, denser, and melts at ~1538°C.

Fluorine, chlorine, bromine and iodine are oxidizing non-metals with decreasing oxidizing power F₂ > Cl₂ > Br₂ > I₂. Cl₂ displaces Br₂ from bromides and I₂ from iodides. Example: Cl₂ + 2KI → 2KCl + I₂.

Metals lose electrons → cations; non-metals can gain electrons → anions.

Alkali metals: Group 1, ns¹, generally M⁺. Alkaline earth metals: Group 2, ns², generally M²⁺.

Metallic reactivity usually increases down Groups 1 and 2.

Halogen oxidizing power: F₂ > Cl₂ > Br₂ > I₂.

Hydrohalic acid strength in water: HI > HBr > HCl > HF.

Bromine is a liquid non-metal; mercury is a liquid metal at ordinary room temperature.

Noble metals such as Au, Ag and Pt resist corrosion.

Carbonate → CO₂; chloride → white AgCl; iodide → yellow AgI; sulfate → white BaSO₄.

Group 1 metals are alkali metals; Group 2 metals are alkaline earth metals. Magnesium reacts slowly with cold water. Gypsum is CaSO₄·2H₂O. Electron affinity and electronegativity describe different properties.

Metal fires, reactive sodium, toxic nitrogen oxides, strong acids, barium salts and other reagents require trained laboratory supervision. Chemical tests here are for understanding, not unsupervised practice.