Comprehensive Notes on Atomic Structure, Periodicity, Bonding, and Carbon Chemistry of Salts, and Carbon

Atomic Structure and Chemical Identification

The Atom: Defined as the smallest particle of an element that takes part in a chemical reaction.
Subatomic Particles: The atom consists of three primary particles:
- Protons:
  - Charge: Positively charged ( $+1$ ).
  - Location: Found in the nucleus (center of the atom).
  - Relative Mass: $1$ .
  - Significance: The number of protons determines the Atomic Number of an element.
- Electrons:
  - Charge: Negatively charged ( $-1$ ).
  - Location: Found in fixed regions called energy levels or orbitals surrounding the nucleus.
  - Relative Mass: $\frac{1}{1840}$ .
  - Balance: In a neutral atom, the number of protons and electrons is always equal.
- Neutrons:
  - Charge: Neutral (no charge).
  - Location: Found in the nucleus along with protons.
  - Relative Mass: $1$ .
  - Significance: The sum of protons and neutrons determines the Mass Number of an element.

Data of the First Twenty Elements

Element	Symbol	Protons	Electrons	Neutrons	Atomic Number	Mass Number
Hydrogen	$\text{H}$	$1$	$1$	$0$	$1$	$1$
Helium	$\text{He}$	$2$	$2$	$2$	$2$	$4$
Lithium	$\text{Li}$	$3$	$3$	$4$	$3$	$7$
Beryllium	$\text{Be}$	$4$	$4$	$5$	$4$	$9$
Boron	$\text{B}$	$5$	$5$	$6$	$5$	$11$
Carbon	$\text{C}$	$6$	$6$	$6$	$6$	$12$
Nitrogen	$\text{N}$	$7$	$7$	$7$	$7$	$14$
Oxygen	$\text{O}$	$8$	$8$	$8$	$8$	$16$
Fluorine	$\text{F}$	$9$	$9$	$10$	$9$	$19$
Neon	$\text{Ne}$	$10$	$10$	$10$	$10$	$20$
Sodium	$\text{Na}$	$11$	$11$	$12$	$11$	$23$
Magnesium	$\text{Mg}$	$12$	$12$	$12$	$12$	$24$
Aluminium	$\text{Al}$	$13$	$13$	$14$	$13$	$27$
Silicon	$\text{Si}$	$14$	$14$	$14$	$14$	$28$
Phosphorus	$\text{P}$	$15$	$15$	$16$	$15$	$31$
Sulphur	$\text{S}$	$16$	$16$	$16$	$16$	$32$
Chlorine	$\text{Cl}$	$17$	$17$	$18$	$17$	$35$
Argon	$\text{Ar}$	$18$	$18$	$22$	$18$	$40$
Potassium	$\text{K}$	$19$	$19$	$20$	$19$	$39$
Calcium	$\text{Ca}$	$20$	$20$	$20$	$20$	$40$

Isotopes and Relative Atomic Mass

Isotopes: Atoms of the same element having the same number of protons (Atomic Number) but different numbers of neutrons (and thus different Mass Numbers).
Conventional Notation: Isotopes are written as ${}^{m}_{n}\text{X}$ , where $m$ is the mass number, $n$ is the atomic number, and $\text{X}$ is the chemical symbol.
Common Isotope Examples:
- Hydrogen: ${}^{1}_{1}\text{H}$ (Protium), ${}^{2}_{1}\text{H}$ (Deuterium), ${}^{3}_{1}\text{H}$ (Tritium).
- Chlorine: ${}^{35}_{17}\text{Cl}$ and ${}^{37}_{17}\text{Cl}$ .
- Uranium: ${}^{235}_{92}\text{U}$ and ${}^{238}_{92}\text{U}$ .
Relative Atomic Mass (RAM):
- Atoms are too small to weigh directly ( $10^{-22}\,\text{g}$ ); therefore, masses are compared to a standard.
- Standard: The carbon-12 isotope ( ${}^{12}\text{C}$ ), arbitrarily assigned a mass of $12.000\,\text{a.m.u.}$ (atomic mass units).
- Definition: The mass of an average atom of an element compared to $\frac{1}{12}$ the mass of an atom of ${}^{12}\text{C}$ .
- Formula: $\text{RAM} = \frac{\text{mass of atom of an element}}{\frac{1}{12} \times \text{mass of one atom of } {}^{12}\text{C}}$ .
Measuring RAM: Determined accurately using a mass spectrometer, which identifies isotopes and their relative abundances.
Calculations of RAM:
- Chlorine: Occurs as $75\% \, {}^{35}_{17}\text{Cl}$ and $25\% \, {}^{37}_{17}\text{Cl}$ .
- $\text{RAM} = (\frac{75}{100} \times 35) + (\frac{25}{100} \times 37) = 35.5$
- Potassium: Occurs as $93.1\% \, {}^{39}_{19}\text{K}$ , $0.01\% \, {}^{40}_{19}\text{K}$ , and $6.89\% \, {}^{41}_{19}\text{K}$ .
- $\text{RAM} = (\frac{93.1}{100} \times 39) + (\frac{0.01}{100} \times 40) + (\frac{6.89}{100} \times 41) = 39.1$ (Note: The value in text calculation for potassium indicates closer to $39$ due to high abundance of ${}^{39}\text{K}$ ).
- Neon: Occurs as $90.92\% \, {}^{20}_{10}\text{Ne}$ , $0.26\% \, {}^{21}_{10}\text{Ne}$ , and $8.82\% \, {}^{22}_{10}\text{Ne}$ .
- $\text{RAM} \approx 20.18$ .

Electron Configuration and Energy Levels

Energy Level: A fixed region surrounding the nucleus occupied by electrons of the same potential energy. They are numbered $1, 2, 3 \dots$ outwards from the nucleus.
Occupancy Rules:
- $1^{st}$ Level: Maximum of $2$ electrons.
- $2^{nd}$ Level: Maximum of $8$ electrons.
- $3^{rd}$ Level: Maximum of $8$ (or $18$ if available) electrons.
- $4^{th}$ Level: Maximum of $8$ (or $18/32$ ) electrons.
Electron Configurations (Examples):
- Hydrogen ( $1$ proton): $1$
- Helium ( $2$ protons): $2$
- Lithium ( $3$ protons): $2:1$
- Sodium ( $11$ protons): $2:8:1$
- Calcium ( $20$ protons): $2:8:8:2$

The Periodic Table

History: Formulated by Henry Moseley in $1913$ , building on work by Dmitri Mendeleev.
Organization: Arrangement of elements horizontally and vertically based on atomic numbers.
Periods: The horizontal rows.
- The number of energy levels in an atom's configuration determines its period number.
- Example: ${}^{23}_{11}\text{Na}$ (config: $2:8:1$ ) has $3$ levels, so it is in Period 3.
Groups: The vertical columns.
- The number of electrons in the outermost energy level determines the group number.
- Example: ${}^{23}_{11}\text{Na}$ (config: $2:8:1$ ) has $1$ outer electron, so it is in Group I.
- Groups are named with Roman Numerals (I, II, III…).
Stability: Atoms with a maximum number of electrons in their outer shell are stable (Group $8/18$ , Noble Gases). Others are unstable and react to achieve stability via electron transfer.

Ion Formation and Oxidation States

Ion: Formed when an unstable atom gains or loses electrons to achieve a stable configuration.
Cations: Positively charged ions formed when elements (usually metals) lose electrons.
- Example: $\text{Na} \rightarrow \text{Na}^{+} + e^{-}$
- Aluminium: $\text{Al} \rightarrow \text{Al}^{3+} + 3e^{-}$ ( $2:8:3 \rightarrow 2:8$ )
Anions: Negatively charged ions formed when elements (usually non-metals) gain electrons.
- Example: $\text{O} + 2e^{-} \rightarrow \text{O}^{2-}$ ( $2:6 \rightarrow 2:8$ )
- Chlorine: $\text{Cl} + e^{-} \rightarrow \text{Cl}^{-}$ ( $2:8:7 \rightarrow 2:8:8$ )
Oxidation and Reduction:
- Oxidation: The process of losing electrons.
- Reduction: The process of gaining electrons.
Variable Oxidation States: Some elements can carry different charges.
- Copper: $\text{Cu}^{+}$ (I) and $\text{Cu}^{2+}$ (II).
- Iron: $\text{Fe}^{2+}$ (II) and $\text{Fe}^{3+}$ (III).
- Manganese: $\text{Mn}^{2+}$ up to $\text{Mn}^{7+}$ .

Valency and Chemical Formulas

Valency: The combining power of an element, equivalent to the number of hydrogen atoms it can displace or the number of electrons it must gain/lose to become stable.
Radicals: Groups of atoms that react as a single unit.

Radical Name	Formula	Valency
Ammonium	$\text{NH}_4^{+}$	$1$
Hydroxide	$\text{OH}^{-}$	$1$
Nitrate(V)	$\text{NO}_3^{-}$	$1$
Hydrogen carbonate	$\text{HCO}_3^{-}$	$1$
Sulphate(VI)	$\text{SO}_4^{2-}$	$2$
Carbonate(IV)	$\text{CO}_3^{2-}$	$2$
Phosphate(V)	$\text{PO}_4^{3-}$	$3$

Deriving Formulas: Use the "interchange valency" method.
- Aluminium Oxide: $\text{Al}^{3}$ and $\text{O}^{2} \rightarrow \text{Al}_2\text{O}_3$ .
- Calcium Hydrogen Carbonate: $\text{Ca}^{2}$ and $\text{HCO}_3^{1} \rightarrow \text{Ca}(\text{HCO}_3)_2$ .

Chemical Equations and State Symbols

State Symbols:
- $(s)$ : Solid
- $(l)$ : Liquid
- $(g)$ : Gas
- $(aq)$ : Aqueous (dissolved in water).
Balancing Equations: Ensure the number of atoms of each element on the reactant side equals the product side.
- Example: $\text{Zn}(s) + 2\text{HCl}(aq) \rightarrow \text{ZnCl}_2(aq) + \text{H}_2(g)$ .
- Example: $2\text{H}_2\text{O}_2(aq) \rightarrow 2\text{H}_2\text{O}(l) + \text{O}_2(g)$ .

Group I: Alkali Metals (Li, Na, K, Rb, Cs, Fr)

Characteristics: Monovalent, $1$ electron in the outer shared level.
Trends Down the Group:
- Atomic/Ionic Radius: Increases because the number of energy levels increases.
- Electropositivity: Increases down the group as the nucleus exerts less pull on the outer electron.
- Ionization Energy: Decreases because removing the outer electron becomes easier as it gets further from the nucleus.
- Melting/Boiling Points: Decrease because atomic size increases, weakening the metallic bond.
Reactions:
- With Water: React vigorously to form a metal hydroxide and hydrogen gas.
- $2\text{Na}(s) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{NaOH}(aq) + \text{H}_2(g)$ .
- With Chlorine: Form white metal chloride solids.
- $2\text{Li}(s) + \text{Cl}_2(g) \rightarrow 2\text{LiCl}(s)$ .
Uses:
- Sodium: Gold extraction (sodium cyanide), food seasoning ( $\text{NaCl}$ ), coolant in nuclear reactors, detergent manufacture.

Group II: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)

Characteristics: Divalent, $2$ electrons in the outer level ( $M^{2+}$ ).
Physical Properties: Ductile, malleable, high tensile strength. Harder than Group I because they contribute more delocalized electrons to metallic bonding.
Trends:
- Ionization energy decreases down the group.
- Reactivity with water increases down the group ( $\text{Be}$ is unreactive, $\text{Mg}$ reacts slowly with cold water but burns in steam, $\text{Ca}$ reaches moderately).
Reactions:
- Burning in Air: $\text{Mg}$ burns with a blinding white flame to form $\text{MgO}$ and $\text{Mg}_3\text{N}_2$ .
- Reaction with Steam: $\text{Mg}(s) + \text{H}_2\text{O}(g) \rightarrow \text{MgO}(s) + \text{H}_2(g)$ .
Uses: Magnesium in aircraft bodies (duralumin), Calcium for bone health, cement, and raising soil pH.

Group VII: Halogens (F, Cl, Br, I, At)

Characteristics: Diatomic molecules ( $\text{F}_2, \text{Cl}_2 \dots$ ), $7$ outer electrons, monovalent ( $X^{-}$ ).
States at Room Temp:
- $\text{F}_2, \text{Cl}_2$ : Gases.
- $\text{Br}_2$ : Red liquid.
- $\text{I}_2$ : Grey solid.
Electronegativity: Decreases down the group. Fluorine is the most electronegative element ( $4.0$ on Pauling scale).
Displacement Reactions: A more reactive (more electronegative) halogen will displace a less reactive one from its halide solution.
- $\text{Cl}_2(g) + 2\text{NaBr}(aq) \rightarrow 2\text{NaCl}(aq) + \text{Br}_2(aq)$ .
Uses: Fluorine in toothpaste and etching glass ( $\text{HF}$ ), Bromine in photography ( $\text{AgBr}$ ), Iodine as an antiseptic (tincture).

Group VIII: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)

Characteristics: Chemically inert due to stable octet ( $2:8$ ) or duplet ( $2$ ) configurations. Monatomic gases.
Trends: Boiling points increase down the group as Van der Waals forces increase with atomic size.
Uses:
- Ar: Inert atmosphere for light bulbs and arc welding.
- Ne: Advertisement lights/signs.
- He: Weather balloons, diving mixtures, low-temperature thermometers.

Structure and Bonding

Ionic Bonding: Electrostatic attraction between cations and anions. Formation of giant ionic lattices. Properties: high melting points, conduct electricity when molten/aqueous.
Covalent Bonding: Sharing of electron pairs.
- Simple Molecules: $\text{H}_2, \text{Cl}_2, \text{H}_2\text{O}$ . Low melting points, held by weak Van der Waals forces.
- Giant Covalent Structures:
  - Diamond: Each Carbon bonded to $4$ others ( $3D$ tetrahedral). Extremely hard, non-conductor.
  - Graphite: Each Carbon bonded to $3$ others in layers. Contains free delocalized electrons. Soft, slippery, conducts electricity.
  - Silica ( $\text{SiO}_2$ ): High melting point, non-conductor, used in lenses.
Metallic Bonding: Array of positive metal ions in a "sea" of delocalized mobile electrons.

Salts and Their Properties

Definitions:
- Hygroscopic: Absorbs water but does not form a solution ( $\text{CuSO}_4$ , common salt).
- Deliquescent: Absorbs water to form a solution ( $\text{NaOH}$ , $\text{CaCl}_2$ ).
- Efflorescent: Loses water of crystallization to the atmosphere ( $\text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}$ ).
Salts Preparation:
- Direct Displacement: Metal + Acid.
- Neutralization: Alkali + Acid.
- Precipitation (Double Decomposition): Mixing two soluble salts to get an insoluble one ( $\text{BaCl}_2 + \text{Na}_2\text{SO}_4 \rightarrow \text{BaSO}_4(s) + 2\text{NaCl}$ ).
Effect of Heat:
- Nitrates: Alkali nitrates decompose to nitrite and Oxygen; heavy metal nitrates form metal oxide, Nitrogen(IV) oxide (brown gas), and Oxygen.
- Carbonates: Most decompose to metal oxide and Carbon(IV) oxide ( $\text{CO}_2$ ). $\text{Na}_2\text{CO}_3$ and $\text{K}_2\text{CO}_3$ do not decompose.

Electrolysis

Process: Decomposition of an electrolyte by electric current.
Electrodes:
- Anode: Positive electrode; anions migrate here to lose electrons (oxidation).
- Cathode: Negative electrode; cations migrate here to gain electrons (reduction).
Examples:
- Molten $\text{PbCl}_2$ : Grey beads of Lead at cathode; green Chlorine gas at anode.
Industrial Uses: Extraction of reactive metals ( $\text{Na}, \text{Al}$ ), refining Copper, and electroplating (e.g., silver-plating a spoon).

Carbon and Its Compounds

Allotropes: Diamond and Graphite.
Carbon(IV) Oxide ( $\text{CO}_2$ ):
- Preparation: $\text{CaCO}_3(s) + 2\text{HCl}(aq) \rightarrow \text{CaCl}_2(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$ .
- Tests: Turns lime water ( $\text{Ca(OH)}_2$ ) milky due to white $\text{CaCO}_3$ precipitate.
- Uses: Fire extinguishers, aerated drinks, baking powder, and the Solvay Process.
Solvay Process: Industrial manufacture of Sodium Carbonate (Soda Ash).
- Raw materials: Brine ( $\text{NaCl}$ ), Ammonia, Limestone ( $\text{CaCO}_3$ ).
- Key Intermediate: Sodium Hydrogen Carbonate ( $\text{NaHCO}_3$ ).
Carbon(II) Oxide ( $\text{CO}$ ):
- Preparation: Dehydration of methanoic or oxalic acid by concentrated $\text{H}_2\text{SO}_4$ .
- Properties: Neutral gas, highly toxic (binds to hemoglobin), powerful reducing agent in blast furnaces.
- Combustion: Burns with a characteristic blue flame to form $\text{CO}_2$ .