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Solid State

• Solids have definite shape and volume.
• The particles are closely packed with strong attraction forces.

Classification of Solids (Based on Structural Features)

• Crystalline solids exhibit long-range order in arrangement.
• Crystalline solids have sharp melting points and heat of fusion.
• Crystalline solids are anisotropic, meaning physical properties differ in different directions (e.g., Quartz, Diamond).
• Amorphous solids show short-range order in arrangement.
• Amorphous solids lack sharp melting points and heat of fusion.
• Amorphous solids are isotropic, with uniform physical properties in all directions (e.g., Plastic, Rubber).

Classification of Solids (Based on Nature of Particles)

• Molecular solids contain molecules (e.g., HCl, SO2).
• Ionic solids contain ions (e.g., NaCl, MgO).
• Metallic solids contain metal atoms (e.g., Fe, Cu, Ag).
• Covalent solids contain atoms linked by covalent bonds (e.g., SiO2, C (diamond)).

Crystal Lattice

• Crystal lattice is the regular 3D arrangement of constituent particles
• Crystal lattice are also called space lattice.

Unit Cell

• The smallest repeating unit of a crystal lattice.
• The unit determines the crystal's structure.

Types of Unit Cells

• Primitive Unit Cells: Constituent particles are present only at the corners.
• Centered Unit Cells: Constituent particles are present at corners and other positions; three types exist.
• Body-centered unit cells: Particles at body center and corners.
• Face-centered unit cells: Particles at the center of each face and corners.
• End-centered unit cells: Particles at the center of two opposite faces and corners.

Number of Atoms in Different Cubic Unit Cells

• Primitive unit cell contains 1 atom (8 corners x 1/8).
• Body-centered cubic (bcc) unit cell contains 2 atoms [(8 corners x 1/8) + 1 at center].
• Face-centered cubic (fcc) unit cell contains 4 atoms [(8 corners x 1/8) + (6 faces x 1/2)].

Close Packing in Solids

• Particles are closely packed in solids
• Particles are hard spheres in these solids

Packing Efficiency

• Packing efficiency calculates percentage of total space occupied by spheres.
• Packing Efficiency = (Volume occupied by spheres in unit cell / Total volume of unit cell) x 100
• fcc/ccp/hcp packing efficiency is 74%.
• bcc packing efficiency is 68%.
• Primitive/simple cubic packing efficiency is 52%.

Imperfections in Solids (Crystal Defects)

• Crystal defects are deviations from orderly arrangement of particles in a solid.
• Point defects are imperfections around a point (atom).
• Line defects are imperfections along a row.

Stoichiometric Defects

• Stoichiometric defects do not disturb the stoichiometry of the solid.
• Ionic solids show these defects, including Schottky and Frenkel defects.
• Schottky defect involves missing equal numbers of cations and anions from lattice sites.
• Schottky defect decreases density due to the missing ions.
• Schottky defect: Cations and anions have similar size
• Schottky defect: Higher coordination number.
• Schottky defect Eg: NaCl, KCl, CsCl, AgBr
• Frenkel defect involves dislocation of ions from lattice to interstitial site.
• Frenkel defect maintains constant density because ions remain within the crystal.
• Frenkel defect: Ionic size difference is large.
• Frenkel defect: Lower coordination number.
• Frenkel defect Eg: ZnS, AgCl, AgBr, AgI

Non-Stoichiometric Point Defects

• Non-stoichiometric point defects: disturb the stoichiometry of the solid Metal excess defect (anion vacancies): anions missing from lattice sites
• To maintain electrical neutrality in Metal excess defect (anion vacancies), electrons occupy anionic vacancies
• This Metal excess defect (anion vacancies) creates F-centers (color centers).
• F-centres colour the crystals.
• Heating NaCl in Na atmosphere produces yellow color. excess Na gives NaCl yellow.
• Heating KCL in K atmosphere excess K gives KCI violet.
• Heating Li in LiCl gives LiCl pink.
• Metal excess defect (extra cations): extra cations present at interstitial sites
• To maintain Metal excess defect (extra cations) electrical neutrality, neighboring interstitial sites also contain electrons
• When ZnO heated, loses oxygen, turns yellow. Excess Zn at interstitial site has equivalent electrons at neighboring sites
• Metal deficiency defect: cations missing; adjacent metal atoms acquire extra positive charge.
• Metal deficiency defect Examples: FeO (often composition Fe.95O).

Impurity Defects

• Impurity Defects: arise from foreign particles in a crystal
• Crystallizing fused NaCl with SrCl2 causes some Na+ ions to be replaced by Sr2+ ions
• The number of resulting cationic vacancies equals number of occupied Sr 2+ ions.
• Solid solutions of CdCl2/AgCl also exhibit Impurity Defects

Properties of Solids: Electrical Properties

• Conduction of electricity related to band theory.
• Valence band: lower energy electron occupied band.
• Conduction band: higher energy unoccupied band.
• Metals: valence band is partially filled or overlaps with the conduction band allowing easy electron flow
• Semiconductors: Small gap between valence and conduction bands; limited electron entry for partial conduction
• Insulators: Large gap between valence and conduction bands, preventing electrical conduction properties

Properties of Solids: Magnetic Properties

• Paramagnetic substances are weakly attracted by magnetic field
• Paramagnetism results from unpaired electrons
• Paramagnetism absent in the magnetic field's absence
• Paramagnetism Examples: O2, Cu, Fe, Cr
• Diamagnetic substances are weakly repelled by magnetic fields
• Diamagnetism is due to paired electrons
• Diamagnetism Examples: NaCl, H2O, benzene
• Ferromagnetic substances are strongly attracted by magnetic fields
• Ferromagnetism: retains magnetism without the application of a field
• Ferromagnetism: spontaneous alignment of magnetic moments (domains) aligned in same direction
• Ferromagnetism Examples: Fe, Co, Ni, CrO2, Gd (Gadolinium)

Magnetic Substances

• Antiferromagnetic substances have expected high magnetic moments but possess zero magnetic moments.
• Antiferromagnetic substance: magnetic moments align in opposite directions alternatively
• Antiferromagnetic substance Example: MnO
• Ferrimagnetic substances show smaller magnetic moments where magnetic moments are cancelled each other
• Ferrimagnetic substances Examples: Fe3O4 (magnetite), ZnFe2O4 (Zinc ferrite), MgFe2O4 (Magnesium ferrite)

Solutions

• Solutions are homogenous mixtures of two or more substances
• Solute: Substance in lesser quantity
• Solvent :: Substance in higher quantity
• Solubility: the maximum amount of a substance that can be dissolved in specific amount of solvent at a temperature

Solubility of a Gas in a Liquid

• Solubility is affected by applied pressure, explained by Henry's law: higher pressure increases gas solubility
• Henry's law: the partial pressure of gas in the vapor phase is directly proportional to its mole fraction in the solution
• Henry's law: P=KHX (KH is Henry's constant) KH value is inversely proportional to gas solubility in a liquid.
• Henry's law Application: To increase CO2 solubility in soft drinks.
• Henry's law Application: Scuba divers Bends happen due to N2 solubility in blood
• To avoid bends Scuba divers dilute oxygen with less soluble Helium
• Henry's law Application: Anoxia at high altitudes due to low oxygen partial pressure
• KH value increases with temperature
• Solubility decreases with temperature (the value above is increases), aquatic life is more comfortable in cold water

Vapour Pressure of Liquid-Liquid Solutions

• Raoult's Law: partial vapor pressure of each solution component is directly proportional to its mole fraction For component 1: P1 is directly proportional to X1, P1=P1°X1 For component 2: P2 is directly proportional to X2 or P2 =P2°X2

Ideal Solutions

• Ideal solutions obey Raoult's law.
• In Ideal solutions ΔH mix = 0 and ΔV mix = 0.
• Ideal binary solutions have A-A (solvent-solvent) and B-B (solute-solute) interactions equal to A-B (solvent-solute) interactions.
• Ideal solution Examples: n-hexane and n-heptane, bromoethane and chloroethane, benzene and toluene.

Non-Ideal Solutions

• Non-Ideal solutions do not obey Raoult's law.
• In Non-Ideal solutions ΔH mix = not equal to 0 and ΔV mix = not equal to 0.
• Non-Ideal solution behavior is either higher or lower than Raoult's law predicts divided into two deviations.
• Positive deviation behavior non-Ideal solutions: observed vapor pressure is greater than expected
• Positive deviation: A-B interaction is less than A-A and B-B
• Positive deviation: Change in volume or enthalpy greater than zero and shows Positive deviation from Roults law: examples Ethanol and acetone
• Negative deviation behavior non-Ideal solutions: partial vapor pressure is less than theoretical
• Negative deviation: A-B > A-A and B-B
• Negative deviation: volume and enthalpy change less than zero .
• Negative deviation Examples: Chloroform and acetone

Azeotropic Mixtures

• Azeotropic mixtures are binary mixtures with same composition in liquid/vapor phases, boil at constant temperature.
• Separation via fractional distillation is not possible for azeotropes.
• Two Azeotropic mixture types exist.
• Minimum boiling azeotropes showing large positive deviation from Raoult's law.

Azeotropic Mixtures

• Maximum boiling azeotropes show a large negative deviation from Raoult's law.

Colligative Properties

• Colligative properties depend on the number of solute particles but not their nature.
• Relative lowering of vapor pressure Colligative properties
• Elevation of boiling point Colligative properties
• Depression of freezing point Colligative properties
• Osmotic pressure Colligative properties -

Colligative Property Applications

• Used to determine molar mass of non-volatile solutes (M2) using equations.
• Relative lowering of VP: (P°-Ps)/Ps = Xsolute, M2 = (P1.w2.M1)/W1 (P°-P1)
• Elevation of BP: ΔTb = Kb.m, M2 = (Kb.w2.1000)/ΔTb.W1
• Depression of FP: ΔTf = Kf.m, M2 = (Kf.w2.1000)/ΔTf.W
• Osmotic pressure: π = CRT, M2 = (w2.R.T)/πV

Osmosis

• Osmosis: solvent molecules move from low to high concentration through semi-permeable membrane.
• Examples of semi-permeable membranes: egg membrane, animal/plant membranes, cellulose acetate.

Osmotic Pressure

• Osmotic pressure is excess pressure needed to be applied on a solution to stop osmosis.

Reverse Osmosis

• Reverse osmosis: reversing osmosis by applying larger pressure on solution, purifying solvent through a semi-permeable membrane.
• Reverse osmosis Applications: desalination of sea water and water purifiers.

Isotonic Solutions

• Isotonic solutions have same osmotic pressure at given temperature.
• Isotonic Example: blood cells and 0.9% NaCl solution.

Hypertonic And Hypotonic Solution

• Hypertonic solution has higher osmotic pressure.
• Hypotonic solution has lower osmotic pressure.

Importance of Colligative Properties

• Osmotic pressure: measures very high molecular mass of proteins, polymers for two reasons:
• Osmotic pressure is measured at room temperature
• Osmotic pressure Molarity is used instead of molality.
• antifreeze solutions (glycol) are used in automobile radiators prevent water the freezing.
• NaCl, CaCl2, is used to remove ice from road : freezing point lowering
• Blood cell fluid's osmotic pressure balances 0.9% mass/volume NaCl i.e. isotonic: , therefore, are injected.

Abnormal Molecular Mass

• Abnormal Molecular Mass: calculated molecular mass via colligative properties appears higher or lower than normal due to two reasons:
• Particles dissociate: number increases, boosting colligative properties, lowering mass.
• Particles associate: number decreases, and colligative properties decrease increasing mass.
• Van't Hoff factor (i) is a the ratio Normal molecular mass/Abnormal molecular mass
• Molecular mass For association, i < 1 . For dissociation, i >1.

Electrochemistry

• Electrochemistry deals with the relationship between Chemical energy and electrical energy and their inter conversions.
• Daniel cell consists of Zn rod dipped in ZnSO4 solution and Cu rod in CuSO4 solution.

Galvanic/Voltaic cell (Daniel cell)

• The two solutions in a Galvanic/Voltaic cell connected externally by a metallic wire (voltmeter/switch) and internally by a salt bridge

Galvanic/Voltaic cell (Daniel cell)

• (i) Reduction: Cu2+ + 2 e→ Cu(s) the right half cell
• (ii) Oxidation: Zn(s) → Zn2+ + 2 e- left half cell-
• Left half cell-Oxidation-Anode-Negative
• Right half cell-Reduction-Cathode-Positive.
• Electron flow from Zinc rod to Copper electrode; current flow from Copper electrode to Zinc electrode.

Electrode Potential (Eel)

• Electrode Potential is the tendency for a metal to loss or gain when connected with its own solution
• Electrode potential is called Standard electrode potential(Eºel)

Cell Potential

• Cell potential is difference between cathode and anode electrode potentials.
• Ecell = E cathode - E anode = E right - E left
• For the Daniel cell, cell is symbolically represented is Zn(s)/Zn2+(aq)//Cu2+(aq)/Cu(s)

Measurement of Electrode Potential

• Measurement of Electrode Potential with Standard Hydrogen Electrode/Normal Hydrogen Electrode

Measurement of Electrode Potential

• Measurement of Electrode Potential. SHE can be represented as Pt(s)/H2(g)/H+(aq) when it acts as anode or H+(aq)/H2(g)/Pt(s) when it acts as cathode.

Electrochemical Series (ECS)

• Electrochemical Series is electrodes arranged by their Electrode Potential.
• Electrochemical Series applications:
• comparing metal reactivity
• predicting displacement reactions
• predicting liberation of H2 gas from acids

Nernst equation

For a general electrode reaction M +(aq) + ne¯ → M(s)

Nernst equation - Various forms

• Electrode potential: Esub(Mⁿ+/M)sub = E°sub(Mⁿ+/M)sub - (RT/nF) ln ([M]/[Mⁿ+])sub

• Cell potential for Daniel cell: Zn + Cu²⁺ → Zn²⁺ + Cu Esubcell sub= E°subcell - (0.059/2) log([Zn²⁺]/[Cu²⁺])

• Free energy and EMF of cell:

• ΔG = -nFEsub(cell)

Equilibrium Constant and EMF of cell: E°subcell sub=((2.303RT)/nF) log Ksubc sub

Molar Conductivity (Ʌm)

• Molar conductivity is the conductivity of 1 mole of an electrolytic solution between two electrodes of unit length
• Molar conductivity decreases with increasing electrolyte concentration but conductivity increases concentration
• For strong and weak electrolyte, molar conductivity increases with dilution.
• For weak electrolytes, molar conductivity rises during dilution versus slight electrolyte increase.
• With maximum dilution and concentration approaches zero; molar conductivity becomes maximum/limiting, noted subɅm°sub.

Kohlrausch's Law

• The law states which the limiting molar conductivity of an electrolyte will be represented as the total molar ionic conductances.
• Let the molar ionic conductances of anion and cation at infinite dilutions are λ + and λ -,Then subɅm°sub= V+λ°sub+ subV-λ°sub
• Where 0+and v. represents the total number of cations and anions

Applications for the Determination

• ^0m of weak electrolytes. By knowing the ^0m values of strong electrolytes, we can calculate the ^0m of weak electrolytes. For e.g. we can determine the ^0mº of acetic acid (CH3COOH) by knowing the ^0m of CH3COONa, NaCl and HCl as follows:

• ^0m* (CH3COONa) +^0m (HCI) - ^0m * (NaCl) = ^0mCH3COOH Determination of degree of dissociation of weak electrolytes

• By knowing the molar conductivity at a particular concentration (^m) and limiting molar conductivity (^mº), we can calculate the degree of dissociation (α) as, α = sub(Ʌ)/(Ʌ°)!sub

Faraday's Laws of Electrolysis

• Faraday’s first law states Amount of deposited substance/liberated as directly equal to electricity quantity transferred through electrolyte.
• Faraday's second law states different substance, electricity is transferred in solutions, substance or liberated at different quantities
• For product of Electrolysis Electrolyte Aq.NaCl, Product at Anode Chlorine, At cathode Hydrogen.
• For product of Electrolysis Electrolyte Molten NaCl, Product at Anode Chlorine, At cathode Sodium.

Batteries/ Commercial Cells

• Non-rechargeable cells Primary cell Drycell, Mercury cell
• Rechargeable cells Secondary cells Lead storage cell,Ni-Cd cells.
• Components:
  • Dry Cell (Leclanche cell): Anode: Zn container, Cathode: Graphite+MnO2, Electrolye: Ammonium chloride
  • Overall Reaction: 2Zn(s) + MnO2 + NHsub4sup+ → 2Zn2+ + MnO(OH) + NH3
  • Mercury cell MercuryAnode:Zinc(Zn) mercury amalgam, Cathode:Paste of HgO and carbon, Electrolye: Paste of KOH and ZnO
  • Overall Reaction: Zn(Hg) + HgO(s) → ZnO(s) + Hg(1)
  • Lead storage cell: 38%H anode 3: Grid of lead Cathode : Grid of lead packed with lead packed lead with dioxide(PbO2). Overall Reaction: Pb(s) +PbO2(s) +2H2SO4(aq) ← 2PbSO4(s) + to 2Hsub2supO(1)
• Overall reaction: 2Zn(s) + MnO2 + NH4+ → 2Zn2+ + MnO(OH) + NH3
• Mercury cell has 1.35 constant potential V since in solution doesn’t involve with any ion with mercury cell mercury cell solution
• In Lead, battery has battery: which was the was in automobile and inverters.

Types of Cells

• Fuel cells are galvanic cells for electricity convert energy of fuel combustion fuels like hydrogen gas, liquid methane, liquid methanol,etc.
• An example fuel cell Hydrogen – Oxygen fuel cell: . which the Apollo’s Space Programme.

Fuel Cell process

• Fuel Cell process, hydrogen gas, and oxygen fuel cell bubbled across aqueous sodium into cathode anode which is anode (-) (+)
• 4OHsub(aq) the cell Reactions into solution are:- -Cathode O2(g) + 2H2O the cell water 2H2(g) + O2(g) → 2 H2O(1)
• Overall reaction It H higher and it. and the is has it. to from water of Water obtained Hsubsub2+2H2O +O fuel cell from be used for from is pollution and is is which
• Water obtained from the electrolysis of heavy hydrogen and can used for.

Corrosion

• Corrosion is slow destruction of metal by it the various Some are of silver. some are The some
• Rusting of Iron : of chemically : an electrochemical the equation. and
• Chemically rust is hydrated ferric that, some