Comprehensive Notes on Groups 1 and 7, Atmosphere, Reactivity, and Metal Extraction

THE ALKALI METALS (GROUP 1)

• The Group 1 elements are called the alkali metals. They include Lithium ($Li$), Sodium ($Na$), Potassium ($K$), Rubidium ($Rb$), Caesium ($Cs$), and Francium ($Fr$).

• Francium ($Fr$):

  • Located at the bottom of the group and is radioactive.

  • One isotope is produced during the radioactive decay of Uranium-235 ($U-235$).

  • It is extremely short-lived; scientists estimate only $20-30\,g$ exists in the Earth’s crust at any time.

• Physical Properties:

  • Melting and Boiling Points: These are very low for metals and decrease as you move down the group. Lithium ($181\,^{\circ}C$) vs. Caesium ($29\,^{\circ}C$).

  • Density: Tends to increase down the group (Lithium: $0.53\,g/cm^3$, Sodium: $0.97\,g/cm^3$, Potassium: $0.86\,g/cm^3$, Rubidium: $1.53\,g/cm^3$, Caesium: $1.88\,g/cm^3$). Lithium, sodium, and potassium float on water.

  • Softness: They are soft and easily cut with a knife, becoming softer down the group.

  • Appearance: Shiny and silver when freshly cut, but tarnish quickly when exposed to air.

• Storage and Handling:

  • Lithium, sodium, and potassium are stored under oil to prevent reaction with oxygen and water vapor.

  • Rubidium and caesium are stored in sealed glass tubes.

  • They must not be handled with bare fingers, as sweat can cause a reaction producing heat and corrosive metal hydroxides.

• Similarities as a Family:

  • They all have one electron in their outer shell ($Li$: 2,1; $Na$: 2,8,1; $K$: 2,8,8,1).

  • They react with water to form hydroxides with the formula $MOH$ (e.g., $NaOH$) and hydrogen ($H_2$).

  • They react with oxygen to form oxides with the formula $M_2O$ (e.g., $Na_2O$).

  • They react with halogens ($X$) to form compounds with the formula $MX$ (e.g., $LiCl$).

  • They form ionic compounds containing the $M^{+}$ ion.

• Reactivity with Water:

  • Lithium: Fizzes slowly, moves on the surface, and eventually disappears. It does not melt because its melting point is high and heat is not generated fast enough.

  • $2Li(s) + 2H_2O(l) \rightarrow 2LiOH(aq) + H_2(g)$

  • Sodium: Fizzes more vigorously, melts into a shiny ball that moves rapidly on the surface.

  • $2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)$

  • Potassium: Reacts even more violently. Enough heat is produced to ignite the hydrogen, burning with a lilac flame. The reaction may end with an explosion.

  • $2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g)$

  • Rubidium and Caesium: React explosively.

• Trend in Reactivity (Chemistry Only):

  • Reactivity increases down the group.

  • These metals react by losing their single outer electron to form a $1+$ ion (e.g., $Na(s) \rightarrow Na^{+}(aq) + e^{-}$).

  • As you move down the group, atoms have more electron shells and are larger. The outer electron becomes further from the positive nucleus.

  • This increased distance results in a weaker electrostatic attraction to the nucleus, making the electron more easily lost.

THE HALOGENS (GROUP 7)

• The halogens include Fluorine ($F$), Chlorine ($Cl$), Bromine ($Br$), Iodine ($I$), and Astatine ($At$). The name means "salt-producing."

• They are non-metallic elements that exist as diatomic molecules ($F_2, Cl_2, Br_2, I_2, At_2$).

• Physical Properties and Trends:

  • Fluorine ($F_2$): Yellow gas.

  • Chlorine ($Cl_2$): Green gas.

  • Bromine ($Br_2$): Red-brown liquid with an orange-brown vapor.

  • Iodine ($I_2$): Grey solid with a purple vapor.

  • Trends: Melting and boiling points increase down the group as relative molecular mass increases, strengthening intermolecular forces. Colors get darker down the group.

  • Astatine ($At$): Radioactive and extremely rare. Predicted to be a black solid with a high melting point.

• Safety: Halogens have poisonous vapors (handled in fume cupboards). Liquid bromine is highly corrosive.

• Chemical Reactions:

  • With Hydrogen: Form hydrogen halides ($HX$). These are acidic, poisonous, and covalently bonded gases that dissolve in water to form acids (e.g., $HCl(g) \rightarrow HCl(aq)$, hydrochloric acid).

  • $H_2(g) + Br_2(g) \rightarrow 2HBr(g)$

  • With Alkali Metals: Form ionic salts.

  • $2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$

• Displacement Reactions:

  • A more reactive halogen will displace a less reactive one from its salt solution.

  • Reactivity Trend: Decreases down the group (Cl > Br > I).

  • Examples:

  • Chlorine + Potassium Bromide: $2KBr(aq) + Cl_2(aq) \rightarrow 2KCl(aq) + Br_2(aq)$. Solution turns orange.

  • Bromine + Potassium Iodide: $2KI(aq) + Br_2(aq) \rightarrow 2KBr(aq) + I_2(aq)$. Solution turns brown.

  • Ionic Equations: Only show the particles that change. Potassium ($K^{+}$) is a spectator ion.

  • $2Br^{-}(aq) + Cl_2(aq) \rightarrow 2Cl^{-}(aq) + Br_2(aq)$

• Redox and Reactivity explanation (Chemistry Only):

  • Halogens react by gaining an electron to form a $1-$ ion. This is reduction. The halogen acts as an oxidising agent.

  • Smaller atoms (like Chlorine) attract an incoming electron more strongly because the outer shell is closer to the nucleus.

  • Larger atoms (like Bromine) have less attraction for the incoming electron due to increased distance from the nucleus.

  • Therefore, Chlorine is a stronger oxidising agent than Bromine or Iodine.

GASES IN THE ATMOSPHERE

• Composition of Dry Air:

  • Nitrogen ($N_2$): approx. $78.1\,$ (about $4/5$).

  • Oxygen ($O_2$): approx. $21.0\,$ (about $1/5$).

  • Argon ($Ar$): approx. $0.9\,$.

  • Carbon Dioxide ($CO_2$): approx. $0.04\,$.

• Determining Oxygen Percentage:

  • Using Copper: 100 cm3 of air passed over heated copper. Copper turns black forming Copper(II) oxide ($2Cu + O_2 \rightarrow 2CuO$). The volume decrease represents the oxygen used.

  • Using Iron Rusting: Iron filings in a conical flask connected to a gas syringe. Iron reacts with oxygen and water over a week. The volume change in the syringe measures oxygen consumption.

  • Using Phosphorus: Phosphorus ignited inside a bell jar floating in water. White smoke of phosphorus oxide forms and dissolves. The water level rises by approx. $21\,$.

• Combustion of Elements in Oxygen:

  • Magnesium: Burns with a bright white flame to form white magnesium oxide ($2Mg(s) + O_2(g) \rightarrow 2MgO(s)$). Dissolves slightly to form alkaline $Mg(OH)_2$.

  • Sulfur: Burns with a blue flame to form poisonous sulfur dioxide ($S(s) + O_2(g) \rightarrow SO_2(g)$). Dissolves in water to form acidic sulfurous acid ($H_2SO_3$).

  • Hydrogen: Burns with a pale blue flame to form water ($2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$).

• Properties of Oxides:

  • Metal Oxides: Usually ionic, basic, and insoluble (if soluble, they form alkaline solutions with $OH^{-}$ ions).

  • Non-metal Oxides: Covalent, acidic, and often soluble in water (forming acidic solutions with $H^{+}$ ions).

• Carbon Dioxide ($CO_2$):

  • Prepared via reaction: $CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + CO_2(g) + H_2O(l)$.

  • Obtained via thermal decomposition of metal carbonates: $CuCO_3(s) \xrightarrow{heat} CuO(s) + CO_2(g)$.

  • Greenhouse Effect: $CO_2$ absorbs infrared (IR) radiation radiated by the Earth, trapping heat in the atmosphere. Increasing levels from burning fossil fuels and deforestation lead to global warming/climate change.

  • Why CO₂ is Greenhouse Gas: It is a non-polar linear molecule, but vibration makes it non-symmetrical and polar, allowing it to absorb IR. Oxygen ($O_2$) and Nitrogen ($N_2$) are symmetric and non-polar even when vibrating, so they cannot absorb IR.

THE REACTIVITY SERIES

• The Order of Reactivity: Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium, (Carbon), Zinc, Iron, (Hydrogen), Copper, Silver, Gold.

• Redox Definitions:

  • Oxidation: Gain of oxygen or loss of electrons.

  • Reduction: Loss of oxygen or gain of electrons.

  • Redox Reaction: Both oxidation and reduction occur simultaneously.

  • Oxidising Agent: Gives oxygen or takes electrons; it is reduced.

  • Reducing Agent: Takes oxygen or gives electrons; it is oxidised.

  • Mnemonic: OILRIG (Oxidation Is Loss, Reduction Is Gain).

• Displacement involving Oxides:

  • A more reactive metal displaces a less reactive one (Competition reactions).

  • $Mg(s) + CuO(s) \rightarrow MgO(s) + Cu(s)$ ($Mg$ is oxidised, $CuO$ is reduced).

  • $C(s) + 2CuO(s) \rightarrow CO_2(g) + 2Cu(s)$ (Carbon displaces Copper).

• Displacement in Solutions:

  • $Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$.

  • Ionic Equation: $Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$.

  • Half-equations: $Zn \rightarrow Zn^{2+} + 2e^{-}$ (Oxidation); $Cu^{2+} + 2e^{-} \rightarrow Cu$ (Reduction).

• Reactions with Water/Steam:

  • With Cold Water: Metal + Water → Metal Hydroxide + Hydrogen.

  • Calcium: $Ca(s) + 2H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g)$.

  • With Steam: Metal + Steam → Metal Oxide + Hydrogen.

  • Magnesium: $Mg(s) + H_2O(g) \rightarrow MgO(s) + H_2(g)$.

  • Zinc: $Zn(s) + H_2O(g) \rightarrow ZnO(s) + H_2(g)$.

  • Iron: $3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$.

• Reactions with Dilute Acids (MASH: Metal + Acid → Salt + Hydrogen):

  • Magnesium: Vigorous reaction, heat produced. $Mg + 2HCl \rightarrow MgCl_2 + H_2$.

  • Aluminium: Reacts slowly at first due to a protective aluminium oxide layer; reacts vigorously once layer is removed by heating.

  • Iron: Slow fizzing, pale green solution formed ($FeCl_2$).

  • Copper/Silver/Gold: No reaction.

• Rusting of Iron:

  • Conditions: Requires both oxygen and water. Sped up by salt.

  • Formula: Hydrated iron(III) oxide ($Fe_2O_3 \cdot xH_2O$).

  • Prevention:

  • Barriers: Painting, oil/grease, plastic coating, or tin plating.

  • Galvanising: Coating with Zinc. Zinc is more reactive and reacts in preference to Iron (sacrificial protection), even when scratched.

  • Sacrificial Protection: Attaching blocks of more reactive metal (Zinc, Magnesium, Aluminium) to large structures like ship hulls or pipelines.

EXTRACTION AND USES OF METALS (CHEMISTRY ONLY)

• Ores: Rocks containing enough minerals to be worthwhile for metal extraction.

• Native Metals: Unreactive metals found uncombined (Gold, Silver).

• Extraction Methods:

  • Below Carbon: Extracted by heating with carbon (Reduction).

  • Iron Extraction: $Fe_2O_3 + 3C \rightarrow 2Fe + 3CO$ (Main reducing agent is actually $CO$: $Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$).

  • Above Carbon: Extracted via Electrolysis (requires high electricity costs).

  • Aluminium: Dissolved in molten cryolite to lower melting point from over $2000\,^{\circ}C$ to $1000\,^{\circ}C$.

  • Cathode: $Al^{3+} + 3e^{-} \rightarrow Al$.

  • Anode: $2O^{2-} \rightarrow O_2 + 4e^{-}$.

• Alloys:

  • Mixture of a metal with other elements (e.g., Brass = Copper + Zinc; Steel = Iron + Carbon).

  • Why they are harder: Different atom sizes disrupt the regular lattice, preventing layers from sliding over each other.

• Steel Types:

  • Mild Steel: Up to $0.25\,$ Carbon. Malleable, used for car bodies and bridges. Disadvantage: Rusts.

  • High-carbon Steel: $0.6-1.2\,$ Carbon. Hard, brittle, used for cutting tools.

  • Stainless Steel: Iron + Chromium (+ Nickel). Resists corrosion, used for cutlery and chemical vessels.

ACIDS, ALKALIS AND TITRATIONS

• pH Scale:

  • 0-3: Strongly acidic (e.g., $HCl$).

  • 4-6: Weakly acidic (e.g., ethanoic acid/vinegar).

  • 7: Neutral (e.g., sodium chloride).

  • 8-10: Weakly alkaline (e.g., ammonia).

  • 11-14: Strongly alkaline (e.g., $NaOH$).

• Indicators:

  • Litmus: Red (Acid), Blue (Alkali).

  • Methyl Orange: Red (Acid), Yellow (Alkali).

  • Phenolphthalein: Colourless (Acid), Pink (Alkali).

  • Universal Indicator: Changes color across the whole pH range; used for approximate pH.

• Definitions:

  • Acids: Source of hydrogen ions ($H^{+}$) in aqueous solution ($HCl \rightarrow H^{+} + Cl^{-}$).

  • Alkalis: Source of hydroxide ions ($OH^{-}$) in solution ($NaOH \rightarrow Na^{+} + OH^{-}$).

  • Neutralisation: $H^{+}(aq) + OH^{-}(aq) \rightarrow H_2O(l)$.

• Titration (Chemistry Only):

  • Used to find the volume of acid needed to neutralise an alkali.

  • Method:

  • Use a pipette and pipette filler to measure exactly $25.0\,cm^3$ of alkali into a conical flask.

  • Add indicator (e.g., phenolphthalein).

  • Fill a burette with acid (record initial reading to the nearest $0.05\,cm^3$).

  • Add acid to the flask slowly (dropwise near the endpoint) until color changes permanently.

  • Record final burette reading. Calculate "titre" (volume added).