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THE PERIODIC TABLE

ELEMENTS & NON-METALS

• Abbreviations of Elements: Cu, Al, Cr, Ac, Ta, Ni, Mg, Bi, Ta, Ti, V, Os, Co, Mo, Ge, S, Si, Pt, Cr, Ga, I, N.

• Catalysts in Chemistry: Reference to enhance potential and reactions.

8.1 ARRANGEMENT OF ELEMENTS

WHAT ARE ATOMS AND ELEMENTS?

• Atoms:

  • Smallest units of matter that retain the properties of an element.

  • Consist of:

  • Nucleus: Contains protons and neutrons.

  • Electrons: Orbit in shells surrounding the nucleus.

• Elements:

  • A substance made up of only one type of atom.

  • Example: Oxygen consists of all atoms containing 8 protons.

• Periodic Table:

  • Organizes all known elements by atomic number (number of protons).

  • Functions:

  • Identify elements by atomic structure.

  • Predict atomic behavior in chemical reactions.

  • Understand trends and similarities among elements.

  • Represents building blocks of all atoms in nature.

STRUCTURE OF THE PERIODIC TABLE

• Organized in horizontal and vertical columns based on increasing atomic number (proton number).

• Period (row): Corresponds to the number of electron shells an atom possesses.

  • Example: Period 2 has 2 shells; Period 3 has 3 shells.

• Group (column): Corresponds to the number of valence electrons (outer shell electrons).

  • Example: Group I has 1 outer electron; Group VII has 7.

  • Group 0 (VIII): Noble gases with a full outer shell of 8 electrons (except He which has 2).

OTHER IMPORTANT TERMS

ATOMIC RADIUS (Atomic Size)

• Distance from the nucleus to the outermost electron shell.

• Across a Period: Atomic radius decreases (more protons pull electrons closer).

• Down a Group: Atomic radius increases (more electron shells are added).

  • Example: Lithium (Li) has a larger atomic radius than Fluorine (F) in the same period. Sodium (Na) has a larger radius than Lithium because it has more shells.

ELECTRON SHELLS (Energy Levels)

• Electrons arranged in shells around the nucleus.

• Period number equals the number of shells.

  • E.g., Magnesium (Mg) in Period 3 has 3 shells.

IONIZATION ENERGY

• Energy needed to remove an electron from an atom.

• Across a Period: Ionization energy increases.

• Down a Group: Ionization energy decreases.

ELECTRONEGATIVITY

• Ability of an atom to attract electrons in a bond.

• Across a Period: Increases.

• Down a Group: Decreases.

METALS, NON-METALS, AND METALLOIDS

• Periodic Table Composition:

  • Over 100 elements divided into:

  • Metals (the majority)

  • Non-metals

  • Metalloids/Semimetals (properties of both).

• Location:

  • Metals on the left and center, non-metals on the right, metalloids zig-zag in between.

• Metallic Character:

  • Increases down a group (as electrons are lost more easily).

  • Decreases across a period from left to right (atoms gain electrons instead).

• Non-metallic Character:

  • Increases across a period and decreases down a group.

PHYSICAL PROPERTIES

• Metals:

  • Conduct heat and electricity.

  • Malleable (can be shaped).

  • Ductile (can be drawn into wires).

  • Lustrous (shiny).

  • High density and high melting points.

  • Form positive ions by electron loss (basic oxides).

• Non-Metals:

  • Poor conductors of heat and electricity.

  • Brittle when solid; dull, not reflective.

  • Low density and melting points (often gases).

  • Form negative ions by electron gain (except H, which may form positive ions).

  • Form acidic oxides.

ELECTRONIC CONFIGURATION AND PREDICTING PROPERTIES

• Electronic Configuration:

  • Determines the group and period:

  • Number of shells => period.

  • Number of outer electrons => group.

• Valence Electrons:

  • Electrons in the outermost shell.

  • Group number reflects number of valence electrons for the main groups.

  • Example: Group I elements all have 1 valence electron; hence, highly reactive.

• Elements in the same group exhibit similar chemical properties due to the same number of valence electrons.

  • Example: Group I metals form 1+ ions and react with water similarly.

• Valency of an Element:

  • Number of electrons lost, gained, or shared to achieve a full outer shell (8 electrons).

  • Example: An oxygen atom with 6 valence electrons needs 2 more to complete its shell (valency 2).

  • Example: A sodium atom with 1 valence electron loses it to achieve a full outer shell like noble gases (valency 1).

TABLE OF GROUP VALENCY AND IONS FORMED

TRENDS ACROSS PERIODS AND DOWN GROUPS

• Across a Period (left → right):

  • Proton number increases.

  • Electrons are added.

  • Metallic character decreases while non-metallic character increases.

• Atomic Radius: Generally decreases across a period.

• Ionization Energy and Electronegativity: Generally increase across a period.

• Down a Group:

  • Each element has one more electron shell.

  • Atomic radius increases.

  • Ionization energy decreases (valence electrons are farther from the nucleus).

  • For metals: Reactivity increases down Group I.

  • For non-metals: Reactivity decreases down Group VII.

PREDICTING CHEMICAL BEHAVIOR FROM THE PERIODIC TABLE

• Group number predicts the type of ion formed and its charge:

  • Group I (e.g., Na) loses 1 electron → Na⁺

  • Group II (e.g., Mg) loses 2 electrons → Mg²⁺

  • Group VII (e.g., Cl) gains 1 electron → Cl⁻

  • Group VI (e.g., O) gains 2 electrons → O²⁻

• Period number indicates the number of electron shells in an atom.

• Elements in the same group form similar compounds (for instance: NaCl, KCl, LiCl — all white salts).

VISUAL TRENDS SUMMARY TABLE

WHY GROUP MEMBERS BEHAVE SIMILARLY

• Outer Electrons:

  • Determine the nature of chemical reactions.

  • Reactions involve electron transfer or sharing.

  • Same valence configuration results in similar reaction patterns.

• Examples:

  • Lithium (configuration 2,1), sodium (configuration 2,8,1), potassium (2,8,8,1) all belong to Group I and lose 1 electron.

  • Fluorine (2,7), chlorine (2,8,7), bromine (2,8,18,7) all belong to Group VII and gain 1 electron.

TIPS TO REMEMBER

• Group Number = Number of Outer Electrons.

• Metallic Behavior: Metals form positive ions; non-metals form negative ions.

• Reactivity Characteristics:

  • Group I: Soft, reactive, low-density metals (alkali metals).

  • Group VII: Diatomic, reactive non-metals (halogens).

  • Members of the same group display similar chemistry due to the same valence electrons.

8.2 NON-METALS

PHYSICAL PROPERTIES

• Characteristics of Non-Metals:

  • Poor conductors of heat and electricity.

  • Brittle when solid.

  • Dull and not reflective.

  • Low density and melting points (often gases).

  • Form negative ions through electron gain (except hydrogen).

  • Produce acidic oxides.

GROUP VII PROPERTIES - THE HALOGENS

• Overview:

  • Group VII elements are called halogens: Fluorine (F₂), Chlorine (Cl₂), Bromine (Br₂), Iodine (I₂), Astatine (At₂).

  • Diatomic Property: Exists in molecules of two atoms.

  • All halogens have 7 electrons in the outermost shell, gaining 1 electron to form halide ions (X⁻) achieving a stable octet.

  • Reactivity: Reactivity increases up the group, which is opposite to Group I.

PHYSICAL PROPERTIES OF HALOGENS

• Exam Tip: Iodine transitions through different colors:

  • Solid state: Grey-black.

  • Vapor state: Purple.

  • Aqueous state: Brown.

CHEMICAL REACTIVITY OF HALOGENS

• Characteristics:

  • Halogens react by gaining 1 electron to create halide ions (X⁻).

  • Similar chemical properties are a result of their valence configuration (7 electrons).

• Reactivity Rule: Reactivity increases as one moves up the group:

  • Order: F > Cl > Br > I > At.

• Reasoning: Higher reactivity at the top due to fewer electron shells (outer electrons are closer to the nucleus), leading to a stronger electrostatic attraction, making it easier to gain an extra electron.

  • Example:

  • Chlorine (2,8,7) exhibits more reactivity than Bromine (2,8,18,7) because Chlorine can attract an extra electron more strongly.

HALOGEN DISPLACEMENT REACTIONS

• A more reactive halogen can displace a less reactive halogen from a solution of its halide (X⁻).

• This is attributed to more reactive halogens being better oxidizing agents.

• General Reaction:

  • Halogen A (more reactive) + Halide B⁻ (less reactive) → Halide A⁻ + Halogen B (elemental)

• Displacement Examples:

  1. Chlorine + Potassium Bromide:

  • Cl₂ + 2KBr → 2KCl + Br₂ (Orange solution forms, Br₂ is produced).

  1. Bromine + Magnesium Iodide:

  • Br₂ + MgI₂ → MgBr₂ + I₂ (Brown solution forms, I₂ is produced).

    • Note: Iodine and Bromine (I₂ + KBr) yields no reaction since Iodine is less reactive than Bromine.

DISPLACEMENT REACTION SUMMARY TABLE

TRENDS SUMMARY - GROUP VII HALOGENS

GROUP 0 - THE NOBLE GASES

OVERVIEW

• Elements: He, Ne, Ar, Kr, Xe.

• Characteristics:

  • Monoatomic, inert (unreactive), colorless gases.

  • Exhibit very low boiling/melting points.

  • Possess full outer electron shells, leading to high stability and reluctance to form compounds.

ELECTRON CONFIGURATION OF NOBLE GASES

APPLICATIONS OF NOBLE GASES

• Their nonreactivity makes them safe for various applications:

  • Neon: Used for neon lights that glow under electrical current.

  • Helium: Used in balloons and airships — it's non-flammable and lighter than air.

  • Argon: Employed as an inert shielding gas in welding and metal fabrication to prevent oxidation and contamination of hot metals.

SULFUR DIOXIDE (SO₂)

FORMATION & PROPERTIES

• Formed by the burning of sulfur or sulfur-containing fuels.

• Recognized as an air pollutant that contributes to acid rain.

USES OF SO₂

1. Bleaching wood pulp in the paper industry.

2. Production of sulfuric acid (H₂SO₄) using the contact process (refer to chapter 6).

3. Acts as a food preservative by killing bacteria.

FERTILISERS

COMPOUND FERTILISERS

• Definition: Contain more than one essential element for plant growth.

• Example: NPK fertilisers, which provide nitrogen, phosphorus, and potassium (NPK).

NPK FERTILISERS

• Enhance plant growth by supplying essential nutrients:

  1. Nitrogen (N): Crucial for proteins, leaf, and stem growth as well as fruit formation.

  2. Phosphorus (P): Important for root development.

  3. Potassium (K): Supports general plant health, disease resistance, and overall quality.

CARBONATES

PROPERTIES OF CARBONATES

1. Solubility:

   • Soluble: Sodium, potassium, ammonium carbonates.

   • Insoluble: All others (e.g., calcium carbonate).

2. Reaction with Dilute Acids:

   • Metal carbonate + acid → salt + water + carbon dioxide.

3. Thermal Decomposition:

   • Metal carbonate → (heat) → metal oxide + carbon dioxide.

   • Example: CuCO₃ → CuO + CO₂ (decomposes with heat).

   • Note: Sodium and potassium carbonates are resistant to thermal decomposition.

SOURCES & FORMS OF CALCIUM CARBONATE

• Found naturally in limestone rocks near rivers, also exists as:

  1. Marble: Formed from limestone under high pressure (metamorphic rock).

  2. Chalk: Composed of marine algae shells.

USES OF CALCIUM CARBONATE (CaCO₃)

1. Used in extraction of iron (fluxes away impurities).

2. Manufacture of cement.

3. Production of lime via a lime kiln.

4. Neutralizing acidic soils (acts slowly and effectively; insoluble).

5. Lime (CaO) & Slaked Lime [Ca(OH)₂]:

   • Utilized to neutralize acidic soil.

   • Used in flue gas desulfurization:

     • Example reaction: CaO + SO₂ → CaSO₃ (neutralizes acidic sulfur dioxide in power stations).

GLOSSARY

1. Acid Rain: Precipitation with a low pH, formed when sulfur dioxide (SO₂) and nitrogen oxides dissolve in atmospheric moisture.

2. Alkaline Earth Metals: Group II elements (e.g., Ca, Mg), which form 2+ ions by losing two electrons and are reactive metals.

3. Alkali Metals: Group I elements (e.g., Li, Na, K), soft and reactive, forming 1+ ions by losing one electron.

4. Allotropes: Different structural forms of the same element (e.g., diamond and graphite for carbon).

5. Atomic Number (Proton Number): Number of protons in the nucleus, defining the element's identity.

6. Atomic Radius: Distance from the nucleus to the outermost electron shell; decreases across a period, increases down a group.

7. Boiling Point: Temperature at which a liquid becomes gas; increases down Group VII due to stronger intermolecular forces.

8. Carbonate: A salt containing the CO₃²⁻ ion; decomposes on heating (except Group I carbonates).

9. Chemical Property: How a substance reacts chemically; determined by the number of valence electrons.

10. Colour Intensity: A trend in halogens where the colour darkens down the group.

11. Combining Power (Valency): Number of electrons an atom gains, loses, or shares to obtain a stable electron configuration.

12. Compound Fertiliser: Contains more than one essential nutrient (e.g., NPK fertilisers).

13. Diatomic Molecule: Molecule consisting of two bonded atoms (e.g., Cl₂, O₂).

14. Displacement Reaction (Halogens): More reactive halogen displaces a less reactive one from its halide solution.

15. Electronegativity: Measure of an atom's ability to attract electrons in a bond; increases across a period, decreases down a group.

16. Electronic Configuration: Arrangement of electrons; dictates chemical properties and positioning in the table.

17. Element: Pure substance made of one type of atom.

18. Flue Gas Desulfurization: Removal of sulfur dioxide from emissions, usually using lime.

19. Group: Vertical column in the Periodic Table; elements in the same group share similar outer electrons and chemical properties.

20. Halide Ion: Negatively charged ion formed from halogen gaining an electron.

21. Halogen: Group VII element (e.g., F, Cl, Br, I); diatomic and highly reactive.

22. Ion: Atom or molecule with an electric charge from electron gain or loss.

23. Ionization Energy: Energy to remove one mole of electrons from gaseous atoms; increases across a period, decreases down a group.

24. Lime (Calcium Oxide): Basic oxide used in agriculture/industry to neutralize acidic soil and remove SO₂ emissions.

25. Limestone: Sedimentary rock primarily composed of CaCO₃; used in construction and industry.

26. Marble: Metamorphic rock (limestone under heat/pressure); mainly CaCO₃.

27. Melting Point: Temperature at which a solid becomes liquid; increases down Group VII.

28. Metal: Elements that are malleable, ductile, conduct electricity, and form positive ions.

29. Metallic Character: Tendency of an element to lose electrons; increases down a group, decreases across a period.

30. Metalloid: Elements with intermediate properties between metals and non-metals; situated along the zig-zag line in the Periodic Table.

31. Monoatomic: Composed of single atoms; describes noble gases in Group 0.

32. Neutralisation Reaction: Reaction involving the combining of an acid with a base to produce salt and water.

33. Noble Gas: Group 0 element with a full outer shell; chemically inert.

34. Non-Metal: Typically gains electrons in reactions, brittle, non-conductive, and forms negative ions.

35. Non-Metallic Character: Tendency of gaining electrons; increases across a period, decreases down a group.

36. NPK Fertiliser: Compound fertiliser containing nitrogen, phosphorus, and potassium for plant growth.

37. Octet Rule: Atoms gain, lose, or share electrons to achieve a full outer shell of eight electrons.

38. Oxidising Agent: Substance gaining electrons in a redox reaction, causing another to be oxidized.

39. Period: Horizontal row in the Periodic Table; indicates number of electron shells of an atom.

40. Periodic Table: Arrangement of elements in order of atomic number, demonstrating trends.

41. Physical Property: A feature observed without changing chemical identity (e.g., density, boiling point).

42. Reactivity (Group I Metals): Increases down the group due to decreased nuclear attraction.

43. Reactivity (Group VII Halogens): Decreases down the group due to increased size weakening attraction.

44. Salt: Ionic compound from the reaction between an acid and a base or carbonate.

45. Shell (Electron Shell): Set of orbitals indicating energy levels correlating with the Periodic Table.

46. Slaked Lime (Calcium Hydroxide): Used for neutralizing acidic soils and removing sulfur dioxide.

47. Sulfur Dioxide (SO₂): Gas produced by burning sulfur; causes acid rain and is used in various industries.

48. Thermal Decomposition: Breakdown of a compound by heating; e.g., metal carbonate decomposing.

49. Valency: Number of electrons an atom must gain, lose, or share.

PERIODIC TABLE - MIND MAPS

STRUCTURE OF THE PERIODIC TABLE

• Arrangement:

  • Periods represent the same number of electron shells.

  • Groups represent the same number of outer electrons (valence).

  • Organized by increasing atomic number.

  • Notable groups: Group 0 (Noble Gases) - 8 outer electrons (except He has 2).

  • Categories: Metals, Non-Metals, Metalloids.

ELECTRONIC CONFIGURATION & PROPERTIES

• Valency & Group Patterns:

  • Relationship of period to number of shells.

  • Group reflects number of valence electrons.

  • Elements in the same group possess similar chemical properties.

VALENCY & IONS

TRENDS IN THE PERIODIC TABLE

• Period vs Group Trends:

  • Across a Period: Proton number increases, leading to decreased atomic radius, increased ionization energy, and electronegativity.

  • Down a Group: Increased atomic radius, decreased ionization energy, and electronegativity; varied reactivity trends for Group I and Group VII.

GENERAL CHARACTERISTICS OF GROUPS

• Reactivity: Group I increases down; Group VII decreases down.

• Halogens demonstrate clear trends through color intensity and state changes (gas → liquid → solid).

• Oxidizing capabilities decrease down the group.