Understanding Atomic Structure

Atomic Structure

• An atom stands as the smallest independent particle of an element.
• Atoms consist of a nucleus containing protons and neutrons, orbited by electrons.
• Protons carry a positive charge, neutrons are neutral, and electrons are negatively charged.
• The nucleus has an overall positive charge, while an atom has an overall neutral charge.

Subatomic Particles

• Protons, neutrons, and electrons are subatomic particles.
• Their masses and charges are measured in relation to each other using Relative Atomic Mass and Relative Atomic Charge.
• Proton: Relative Atomic Charge = +1, Relative Atomic Mass = 1
• Neutron: Relative Atomic Charge = 0, Relative Atomic Mass = 1
• Electron: Relative Atomic Charge = -1, Relative Atomic Mass = 1/1836

Atomic Number and Mass Number

• Atomic Number/Proton Number equals the number of protons in the nucleus.
• For a neutral atom, the atomic number equals the number of electrons.
• Mass Number/Nucleon Number equals the total number of protons and neutrons in the atom.
• Neutrons can be calculated by subtracting the Atomic Number from the Mass Number.

Atomic Radius

• The atomic radius of an element is half the distance between two nuclei of covalently bonded atoms of the same type.
• Atomic radius decreases across a period and increases down a group.
• Across a period, atomic radius decreases due to an increase in the nucleus charge. More electrons are added in same quantum shell with approximately constant shielding, resulting in greater attraction between nucleus and electrons, pulling them closer.
• Down a group, the quantum shells increase, increasing shielding, resulting in greater atomic radius.

Ionic Radius

• A cation's size is always smaller than its parent atom.
• An anion's size is always greater than its parent atom.
• A cation is always smaller than an anion if they are in the same period.

Ionisation Energy

• Ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous ions.
• It is measured under standard conditions.
• Units are Kilo joule per mole (KJmol-1).
• Ionisation energy is always endothermic because it breaks electrostatic force between proton and electron.
• The first ionisation energy is the energy required to remove an electron from one mole of isolated gaseous atoms to form one mole of 1+ ions.
• The first I.E of an element X can be represented by the equation: X(g) -> X+(g) + e-
• Second I.E is larger than the first I.E because radius becomes smaller upon removing electrons, causing less repulsion, greater nuclear force of attraction, and requiring more I.E for the removal of the second electron.

Factors Affecting Ionisation Energy

• Ionisation energy increases across a period and decreases down a group.
• Atomic Radius: Electrons further from the nucleus experience less attraction and are easier to remove.
• Nuclear Charge: Greater nuclear charge increases the force of attraction and makes electrons more difficult to remove.
• Shielding Effect: Inner quantum shell electrons push the outer quantum shells away from the nucleus, reducing nuclear attraction and making it easier to remove electrons.
• Spin pair repulsion: Paired electrons in the same orbital repel each other, making it easier to remove one electron.
• Across a period, nuclear charge increases, strengthening the attraction between the nucleus and electrons, decreasing the atomic radius, and keeping the shielding effect constant.
• Down the group, atomic radius and shielding effect increases, decreasing I.E

Ionisation Energy Trends (Exceptions)

• An increase in subshell will decrease the I.E as outer subshell electrons are further away from the nucleus such as between Beryllium and Boron
• An example would be Magnesium and Aluminium
• There is a decrease in I.E when a second electron enters the px subshell as it experiences spin-spin repulsion making it easier to remove an electron such between Nitrogen and Oxygen
• An example would be Phosphorus and Sulfur
• Ionisation energy decreases between the last element in a period and the first element of the next due to the change in quantum shell.

Successive Ionisation Energies of an Element

• It increases.
• Removing an electron from an ion is harder than from its neutral atom.
• This is due to the decreased shielding effect and increased nuclear attraction.
• A change in quantum shell causes a large increase in successive ionisation energy.
• The large jump can be used to deduce the group number of an element.
• A large increase between the third and fourth I.E indicates that the element belongs to group 3.

Isotopes

• Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
• An Isotope is represented by element name, a dash, and mass number.
• Example: Carbon-12
• Isotopes have similar chemical properties but different physical properties.
• Isotopes react in the same manner due to the same number of electrons.
• Due to different numbers of neutrons, isotopes have different masses and densities.

Electron Shells

• Electronic configuration arranges electrons in an atom
• Electrons are arranged around the nucleus in principal energy levels/quantum shells
• The energy level/quantum shell is represented by the Principal Quantum Number (n)
• A higher n means the shell is further from the nucleus
• Each quantum shell can hold a fixed number of electrons.

Subshells

• The principal quantum shells are split into subshells
• Subshells are represented by s, p, and d
• The energy of the subshells increases in the order s<p<d
• s subshell contains 2 electrons
• p subshell contains 6 electrons
• d subshell contains 10 electrons

Orbitals

• Subshells contain one or more atomic orbitals.
• Orbitals exist at specific energy levels; electrons can only be found at these levels.
• Each orbital can have a maximum of 2 electrons.
• The s subshell has 1 orbital, the p subshell has 3, and the d subshell has 5.
• Each orbital has a specific shape
• The three orbitals of p subshell are px, py and pz
• n=1, Electrons (2n²)=2, Subshells=1s²
• n=2, Electrons (2n²)=8, Subshells=2s²,2p⁶
• n=3, Electrons (2n²)=18, Subshells=3s²,3p⁶,3d¹⁰
• n=4, Electrons (2n²)=32, Subshells=4s²,4p⁶,4d¹⁰,4f¹⁴

Ground State

• It is the most stable electronic configuration of an atom.
• It has the lowest amount of energy.
• Subshells with lower energy levels are filled first.
• The pattern is disrupted at quantum shells n=3 and above.
• 4s has a lower energy level than 3d.
• Orbitals in the same subshell have the same energy; they are known as degenerate.
• px, py, and pz all have the same energy levels.

S & p orbitals

• s orbitals are spherical in shape.
• The s orbitals are bigger at a higher principal quantum number.
• p orbitals are dumbbell shaped.
• Every p subshell has 3 orbitals except for when n=1.
• The p orbitals become larger and longer with increasing quantum number.

Electronic Configuration

• Subshells are filled in increasing order of energy levels.
• Electrons are spinning charges rotating clockwise or anticlockwise about their own axis.
• Electrons with similar spin repel each other, causing spin-pair repulsion.
• Electrons occupy different orbitals in the same subshell first to avoid spin-pair repulsion.
• They are paired when there are no more empty orbitals.
• Paired electrons spin in opposite directions to minimise repulsion.
• If there are three electrons in the p subshell, all three orbitals will have one electron each.
• A fourth electron in the p subshell will pair in px subshell.

Periodic Table Blocks

• Group 1 and 2 elements are known as s block elements.
• s block elements have their electrons in an s subshell.
• Group 13 to 18 elements are p block elements.
• Their valence electrons are located in the p subshell.
• Transition block elements are d block elements.
• Their valence electrons are in the d subshell.

Exceptions

• Copper has configuration 3d¹⁰,4s¹ instead of 3d⁹,4s².
• Chromium has configuration 3d⁵,4s¹ instead of 3d⁴, 4s².
• A configuration is more stable when the d subshell is half full or full, so one electron moves from 4s subshell to 3d subshell.

Moles & Stoichiometry

• Relative Atomic Mass is the average mass of all the isotopes relative to 1/12th of the mass of Carbon-12.
• Relative Molecular Mass is the average mass of a molecule of an element or compound relative to 1/12th the mass of Carbon-12.
• Relative Formula Mass is the average mass of a formula unit of an ionic compound relative to 1/12th the mass of Carbon-12.
• Relative Isotopic Mass is the mass of an isotope of an element relative to 1/12th the mass of Carbon-12.

Mole and Avogadro Constant

• One mole of any substance contains 6.02 x10²³ atoms.
• This number is also known as Avogadro's constant.
• One mole of a substance has mass equal to its Relative Atomic or Molecular Mass.
• One mole of a gas at room temperature and pressure occupies 24dm³ volume.

Formulae

• Moles = Mass/A, or M
• Moles = Volume x Concentration/1000
• Moles = Volume of gas at r.t.p/24
• Percentage Yield = actual mass / predicted mass x 100

Empirical and Molecular Formula

• Empirical formula is the simplest whole number ratio of atoms in a molecule.
• Molecular formula is the actual number of atoms of all the elements present in a molecule.

Steps to Calculate Empirical Formula

• Divide the mass of elements by the atomic mass to get moles.
• Divide all the values calculated by the smallest value calculated to get ratio.
• The ratio corresponds to the number of atoms of each element in the Empirical Formula.

Combustion

• Combustion is the burning of a substance in air (O₂).
• General equation for combustion of gaseous Hydrocarbons: CxHy(g) + x+y/2O2(g) → XCO2(g) + y/2H2O(1)

Chemical Bonding

• Electronegativity is the ability of an atom to attract or gain electrons.
• Fluorine is the most electronegative atom.
• Hydrogen is the least electronegative non-metal.

Electronegativity Factors & Trends

• Electronegativity increases left to right across a period.
• Electronegativity decreases down the group.
• Increased nuclear charge results in increased electronegativity.
• Increased atomic radius results in decreased electronegativity.
• Increased shielding through increased quantum shells or subshell will decrease electronegativity.
• Metals are less electronegative than nonmetals.

Electronegativity and Bonding

• A large difference in electronegativities between bonded atoms means the bond is ionic.
• Small or no difference means the bond is covalent.

Bond Energy

• It is the energy required to break one mole of covalent bonds in the gaseous state.
• Bond Energy has units KJmol⁻¹.
• Bond Energy is always endothermic.
• A higher bond energy means the covalent bond is stronger.

Bond Length

• Distance between the nuclei of two covalently bonded atoms.
• The greater the nuclear attraction, the lower the bond length.
• The lower the bond length, the stronger the covalent bond.
• Therefore, the lower the bond length, the higher the Bond Energy.

Shapes of Covalent Molecules

• Predicted by the Valence Shell Electron Pair Repulsion Theory (VSEPR).
• Electrons are negatively charged and repel each other when close.
• The repulsion in bonding pair of electrons in a molecule causes it to adopt a shape to minimize the repulsion.
• Lone pairs repel each other more than bond pairs.
• Lone pair-lone pair > lone pair-bond pair > bond pair-bond pair repulsion.
• Bond Pairs = 2, Lone Pairs = 0, Shape Name = Linear, Bond Angles = 180
• Bond Pairs = 3, Lone Pairs = 0, Shape Name = Triangular Planar, Bond Angles = 120
• Bond Pairs = 4, Lone Pairs = 0, Shape Name = Tetrahedral, Bond Angles = 109.5
• Bond Pairs = 5, Lone Pairs = 0, Shape Name = Triangular Bipyramidal, Bond Angles = 90 & 120
• Bond Pairs = 6, Lone Pairs = 0, Shape Name = Octahedral, Bond Angles = 90
• Bond Pairs = 3, Lone Pairs = 1, Shape Name = Triangular Pyramidal, Bond Angles = 107
• Bond Pairs = 2, Lone Pairs = 2, Shape Name = Bent non-linear, Bond Angles = 104.5

Hydrogen Bonding

• It is strongest form of intermolecular bonding.
• It is a type of a permanent dipole-permanent dipole bonding.
• It is present in molecules where Hydrogen is covalently bonded to small, highly electronegative atoms: F, O, N only.
• Due to the electronegativity difference, the bond becomes highly polarised.
• The H becomes so partial positive charged that it can bond with the lone pair of an O or N atom of another molecule.
• Hydrogen bonding causes high melting and boiling points, such as in water and causes high surface tension in water.

Polarity

• Bond polarity is the charge separation in a covalent molecule due to a difference in electronegativities between bonded atoms.
• When two covalently bonded atoms have the same electronegativity, the bond is nonpolar.
• The lesser electronegativity atom gains a partial positive charge (δ+).
• The higher electronegativity atom gains a partial negative charge (δ-).
• The greater the difference in electronegativity, the more polar the bond becomes.

Dipole

• The dipole moment is a measure of how polar a bond is.
• It is represented by an arrow pointing towards the partial negative end of the dipole.

Polarity in Molecules

• Determined by the polarity of each bond and the arrangement of bonds.
• Equal and opposite dipoles cancel each other out.
• Symmetrical molecules like linear, planar, or tetrahedral are non-polar because dipoles are equal and opposite.
• Molecules having lone pairs are polar because they distort symmetry.

Van der Waals Forces

• Intermolecular forces between covalent molecules are known as Van der Waals forces.
• There are two types: temporary/induced dipole - induced dipole forces, and permanent dipole - permanent dipole forces.

Instantaneous dipole - Induced dipole forces

• Present in non-polar covalent molecules
• Occurs when two nonpolar molecules come close together, electrons repel each other
• Repulsion causes electrons to be unevenly distributed, inducing a dipole
• The partial positive end of a molecule attracts the partial negative end of another
• This attraction is induced dipole-induced dipole force
• Induced dipole-induced dipole forces increase with increasing electrons and greater surface area to allow more contact points.

Permanent dipole - permanent dipole forces

• Polar molecules have permanent dipoles.
• The molecules always have partial positive and partial negative ends.
• The forces between two polar molecules are permanent dipole-permanent dipole forces.
• The partial positive end of one molecule attracts the partial negative end of another.
• Permanent dipole-permanent dipole forces are stronger than induced dipole-induced dipole forces in molecules having the same number of electrons.
• Polar molecules therefore have higher melting and boiling points.

Coordinate/Dative bonding

• A dative bond forms when one atom provides both electrons for a covalent bond.
• Sharing is not mutual.
• Represented by an arrow pointing away from the lone pair of electrons that form the bond.
• Al₂Cl₆ is a dimer formed by dative bonding; a chlorine atom of one AlCl₃ gives two electrons to the Aluminum atom of another.

Incomplete and Expanded Octet

• Some species can have more or less than eight electrons in their outer shell.
• GaCl₃, AICI₃, BeCl₂, BF₃, BCI₃ are examples of molecules with incomplete octets.
• PCI₅, SO₂, SO₃, SF₆, SeF₆ are examples of expanded octet.

Sigma and Pi Bonding

• Occurs with the overlapping of two half filled atomic orbitals.
• A greater atomic orbital overlap means a stronger bond.

Sigma Bonds

• Formed by the end to end overlapping of atomic orbitals.
• Both s and p orbitals overlap this way.
• The pair of electrons is found between the two nuclei.
• The force between the electrons and nuclei bonds the atoms together.
• All single covalent bonds are sigma bonds.

Pi Bonds

• Formed from the sideways overlap of adjacent p orbitals.
• Double covalent bonds contain 1 sigma and 1 pi bond.
• Triple covalent bonds contain 1 sigma and 2 pi bonds.

States of Matter

• No intermolecular forces are present between gas molecules
• The actual volume of gas molecules is negligible compared to the volume occupied
• All collisions between gas molecules are elastic, no energy is lost
• The kinetic energy of gas molecules is directly proportional to temperature

Ideal Gases & Real Gases

• Gases that follow the kinetic theory are known as ideal gases, however, gases, in reality, do not follow this theory exactly even though they come close, they are known as real gases
• Under conditions of low temperature and high pressure, real gases deviate from ideal behavior
• Under conditions of high temperature and low pressure, real gases show ideal behavior
• Non-polar and small size gases behave more ideally; e.g He, H2, N2, O2
• Polar with Hydrogen Bonding and large molecule gases deviate from ideal behaviour; NH3

Ideal Gas Equation

• pV = nRT
• p = pressure (Pa)
• V = volume (m³)
• n = number of moles of gas (mol)
• R = gas constant (8.31 JK-1mol-1)
• T = temperature (Kelvin)
• It's necessary to convert to the correct units

Lattice Structures

• Most ionic, covalent and metallic structures are lattice structures
• The ions, atoms or molecules are arranged in a regular and repeating arrangement

Giant Ionic Lattices

• An ionic bond is formed by the transfer of electrons from a metal to a nonmetal atom
• The ions have an electrostatic force of attraction between them
• Ionic compounds are arranged in giant ionic lattices
• The positive and negative ions are arranged in an alternating order
• For example, NaCl and MgO
• They are strong but brittle
• They have high melting and boiling points due to strong electrostatic forces
• Soluble in water as they can form ion-dipole bonds
• Only conduct electricity in molten or aqueous states as ions can move around

Covalent Lattices

• Covalent bonds form via sharing of electrons between nonmetals.
• Covalent compounds can form simple molecular or giant molecular lattices.
• Iodine and Ice have simple molecular lattices, while Sand, Graphite, and Diamond have giant molecular lattices.
• Simple covalent structures have low melting and boiling points.
• They are mostly insoluble in water unless they are polar or can form Hydrogen bonds.
• They do not conduct electricity in solid or liquid state.
• Giant covalent compounds have high melting and boiling points.
• They can be hard or soft, depending on structure.
• They are mostly insoluble in water.
• They mostly do not conduct electricity unless free electrons are available.

Metallic Lattices

• Metals form giant metallic lattices where metal ions are surrounded by a sea of delocalized electrons.
• They are often packed in hexagonal layers or cubic arrangement.
• Metallic compounds are malleable; their layers can slide.
• Metals are strong and hard because of strong forces between ions and electrons.
• Metals have high melting and boiling points.
• Pure metals are insoluble in water.
• They can conduct electricity in solid or liquid states due to delocalized electrons. Giant Ionic: MPs are high, conducts when Molten or aqueous, soluble, hard and brittle, attractive force is electrostatic attraction e.g. NaCI
• Giant Metallic: MPs are high, solid or liquid, insoluble, hard and malleable, attractive force is attraction between electrons and ions e.g. Copper
• Giant Covalent: MPs are V.high except Graphite, no conductivity except graphite, insoluble hard except graphite attractive force is strong covalent bonds e.g. SiO2
• Simple Covalent: MPs are low except Graphite, no conductivity, insoluble unless polar then soluble, attractive force is weak intermolecular e.g. Cl2

Chemical Energetics

• Total chemical energy inside a substance is called enthalpy.
• Enthalpy change is the change in chemical energy during a chemical reaction.
• Enthalpy change is represented by ΔH.
• Enthalpy change can be positive or negative.
• Activation energy is the minimum energy needed for a reaction to take place.
• Activation energy is represented by Ea.

Exothermic Reactions

• Reactions where products have less energy than the reactants.
• Heat energy is given off to the surroundings.
• The enthalpy decreases so ∆H is negative.
• Lower Ea than endothermic reactions.
• Bond making is exothermic.
• If more energy is released when new bonds are formed than required to break bonds, the reaction is exothermic.

Endothermic Reactions

• Reactions where products have greater energy than the reactants.
• Heat energy is absorbed from the surroundings.
• The enthalpy increases so ΔH is positive.
• Have a higher Ea than exothermic reactions.
• Bond breaking is endothermic.
• If more energy is required to break bonds than is released when new bonds are formed, the reaction is endothermic.

Energy Level Diagrams

• It is a graph of the energies of reactants and products against time.

Enthalpy Changes at Standard Conditions

• Enthalpy changes are measured at standard conditions for fair comparison between reactions.
• Pressure 101kPa
• Temperature 298 K (25 degrees)
• Aqueous solutions should be at a concentration of 1.0mol dm-1
• Enthalpy change under standard conditions is represented by ΔH°.

Enthalpy Change of Formation(ΔH°f)

• The enthalpy change when one mole of a compound is formed from its constituent elements under standard conditions.

Enthalpy Change of Combustion(ΔH°c)

• The enthalpy released when one mole of a substance is burnt in air completely under standard conditions

Enthalpy Change of Neutralisation(ΔH°neut)

• The enthalpy released when one mole of water is formed by the reaction between an acid and an alkali under standard conditions.

Bond Energy

• Exact bond energy is the amount of energy required to break one mole of a specific covalent bond Average bond energy is the average of bon energies in different environments.
• Average bond energy calculated through enthalpy changes.
• Average bond energy = total bond energy/number of bonds

Calculating Enthalpy Change with Bond Energies

• ΔH = Total bond energy of reactants - total bond energy of products

Measuring Enthalpy Change

• Q=mcAT
• Q = energy transferred (J)
• m = mass (g)
• c = specific heat capacity (Jg¯¹K-1)
• ΔT = change in temperature

Hess's Law

• The total enthalpy change in a reaction is independent of the route by which the chemical reaction takes place as long as the initial and final conditions are the same.
• Whatever route the reaction takes, the enthalpy change will be the same.
• It is used to calculate enthalpy changes that can't be measured.
• It can be calculated through the enthalpy change of formation and combustion.

Calculating ΔH from Standard Enthalpy Change of Formation

• ΔH = (Total standard enthalpy change of formation of products) - (total enthalpy change of formation of reactants)
• Steps:
  1. Write down balanced equation.
  2. Write down all the elements that the compounds involved in the reaction form from.
  3. Draw arrows correctly from elements to compounds
  4. Apply Hess Law; energy change of direct route should equal indirect route; subtract the enthalpy going the opposite direction of an arrow and add when going in the same direction

Calculating ΔH from Standard Enthalpy Change of Combustion

• ΔH = (Total standard enthalpy change of combustion of reactants) - (total standard enthalpy change of combustion of products)
• Steps:
  1. Write down balanced equation
  2. Write down all the products of combustion of reactants and products
  3. Draw arrows correctly from reactants and products to combustion compounds
  4. Apply Hess Law; energy change of direct route should equal indirect route; subtract the enthalpy going the opposite direction of an arrow and add when going in the same direction

Calculating standard enthalpy change of formation from standard enthalpy change of combustion

• ΔH₁ = (total standard enthalpy change of combustion of reactants) - (total standard enthalpy change of combustion of products)

Electrochemistry

• Oxidation number represents the number of electrons in an atom or ion.
  • Loss of electrons: positive oxidation number.
  • Gain of electrons: negative oxidation number.
  • Total oxidation state of a neutral compound: Zero.
  • Total change in oxidation state in a reaction: equal.

Redox

• Reduction is the gain of electrons or loss of oxidation state.
• Oxidation is the loss of electrons or gain of oxidation state.
• Reactions in which reduction and oxidation take place are known as redox.
• An oxidising agent oxidises another substance and reduces itself.
• A reducing agent reduces another substance and oxidises itself.
• A redox reaction must have both a reducing and oxidising agent.

Disproportionation Reaction

• A reaction where one substance is both oxidised and reduced.

Equilibria

• In a reversible reaction, the products can react to form the reactants back.
• The reaction can proceed in both the forward and backward directions.
• Reversible reactions are represented by a double arrow.

Dynamic Equilibrium

• The rate of the forward reaction equals the rate of the backward direction in a closed system.

Le Chatelier's Principle

• When any dynamic equilibrium is disturbed, it will shift the equilibrium in the forward or backward direction to undo the disturbance and restore the equilibrium.
• Temperature: Increasing temperature shifts the equilibrium towards the endothermic reaction; decreasing shifts towards exothermic.
• Pressure: Increasing pressure shifts equilibrium towards a lesser number of moles of gas.
• Concentration: Increasing the concentration of reactants or decreasing the concentration of products will shift equilibrium towards forward reaction. Decreasing concentration of reactants or increasing concentration of products shifts equilibrium towards backward reaction.

Equilibrium Constant - Concentrations

• The equilibrium constant K is the ratio of concentrations of products in a reaction to the concentrations of the reactants.
• Square brackets [] represent concentration.
• K = [products] / [reactants]
• Concentrations of different products and reactants are multiplied with each other; moles of each substance are taken as a power in the calculation.
• The unit of K is deduced from its calculation.
• A change in temperature affects the value of K.
• Calculation example: 2SO2(g) + O2 -> 2SO3 (g); Kc = [SO3]^2 / [SO2]^2[O2]

Mole Fraction

• It is the number of moles of one gas divided by the total number of moles of gas at equilibrium.

Partial Pressure

• It is the pressure exerted by one gas in a mixture of gases.
  • Sum of partial pressures = total pressure.
  • Partial pressure = mole fraction x total pressure.

Equilibrium Constant - Partial Pressures

• Equilibrium constant Kp is the ratio of the partial pressures of the products to the partial pressure of the reactants.
• Kp = partial pressure of products/partial pressure of reactants
• 2SO2(g) + O2 ⇄ 2SO3 (g); Kp = (p(SO3)²) / (p(SO2)²*p(O2))

Conditions used in Haber's Process

• N₂(g) + 3H₂(g) ⇄ 2NH₃(g)
• Carried out under these conditions: 450-500 degrees temperature, 200-250 atm, and Fe or Fe₂O₃ as a catalyst.
• Although increasing temperature increases the rate of reaction, shifts equilibrium backward because the backward reaction is endothermic which decreases the yield of ammonia.
• Compromising temperature of 450-500 degrees is used.
• Increasing pressure increases both the rate of reaction but also shifts equilibrium forward and increases yield.
• Pressure is kept at 200-250 atm; maintaining higher pressures is expensive.

Conditions used in Contact Process

• 2SO₂(g) + O2 (g) ⇄ 2SO₃(g)
• Increasing pressure shifts the equilibrium forward.
• The reaction is carried out at 1 atm; Kp is already high; higher pressures are expensive and unnecessary.
• Although temperature increases the rate of reaction, it shifts equilibrium backward since the backward reaction is endothermic, decreasing the yield of ammonia.
• Compromising temperature of 450 degrees is used.

Lowry-Bronsted Acid-Base Theory

• According to this theory, acids are proton donors (H+ ions) and bases are proton acceptors.
• A lowry-bronsted acid gives away H+ ions, while a Lowry-Bronsted base accepts H+ ions.
• CH₃CO₂H + H₂O ⇄ H₃O+ + CH₃CO₂⁻
• CH₃CO₂H is the Bronsted Acid, H₂O is the Conjugate Acid, H₃O+ is the Conjugate Acid, and CH₃CO₂⁻ is the Conjugate Base.

pH Scale

• A scale that shows how acidic or alkaline a substance is.
  • Acids have pH below 7.
  • Alkalis have pH above 7.
  • pH 7 is neutral, e.g., water has ph7.
  • The lower the pH, the more acidic a substance.
  • The higher the pH, the more alkaline a substance.
• The most accurate way to measure pH is with a pH meter.
• It can also be measured by a universal indicator which changes colour according to the pH of the solution.
• Acids are red/orange, neutral solutions are green, bases are blue/purple.

Strong & Weak Acids

• Strong acids fully dissociate H+ ion in water.
• The greater the concentration of the H+ ions, the more acidic the substance.
• Weaker acids only partially dissociate H+ ions.
• Stronger acids conduct electricity better due to the concentration of H+ ions.
• Stronger acids are more reactive.

Neutralisation Reaction

• A reaction between an acid and an alkali to produce a salt and water.
• The ionic equation of any neutralisation reaction is the same: H⁺(aq) + OH⁻(aq) → H₂O(I)

pH Titration Curve

• Titration is a technique used to carry out neutralisation reactions.
• Adding a titrant from a burette to a solution in a conical flask with an indicator.
• The titrant is added till the endpoint, which is when the solution changes colour.
• Neutralisation takes place at the end point.
• It is also known as the equivalence point.
• A curve is drawn with the volume of titrant in cm³ on the x-axis and pH on the y-axis. It startsfrom the pH of the solution in the conical flask and goes to the pH of the titrant.
• At the equivalence point, the line is vertical.
• The shape of the curve changes according the pH of the solution and titrant.

Rate of Reaction

• The speed at which a reaction takes place
• rate of reaction = change in reactants or products (mol dm-3)/time

Collision Theory

• For a reaction to take place, the particles need to collide with each other in correct orientation and with enough energy
• The minimum energy required for successful collisions is known as the activation energy
• When particles don't have enough energy or correct orientation, the collision is ineffective and the particles bounce off each other

Factors of Rate of Reaction

• The collision frequency is the number of collisions per unit time
• When collision frequency increases, more particles have energy above Ea
• Thus, increasing the collision frequency, increases the rate of reaction

Catalyst

• A substance that speeds up the rate of reaction without taking part in the reaction
• It increases the rate by lowering the energy of activation

Concentration

• The greater the concentration of a fluid, the more the particles in a given volume
• This increases the chances of collisions and hence the collision frequency
• Increased collision frequency thus means increased rate of reaction
• Reducing concentration reduces collision frequency and rate of reaction
• The rate of reaction keeps changing through a reaction because the concentration of reactants and products keeps changing as more products are formed

Pressure

• Increasing the pressure pushes particles closer and there are more particles in a volume
• Thus the collision frequency increases and the rate of reaction

Temperature

• Increasing the temperature gives the particles more kinetic energy
• The particles then move faster, increasing the collision frequency and rate of reaction

Boltzmann Distribution Curve

• Graph of the distribution of energies in particles in a given temperature.
• Few particles have very low and high energy; most particles are in between.
• Only a small amount of particles will have enough energy for successful collisions.

Changes in Temperature

• When the temperature is increased, particles gain more energy.
• The collision frequency is increased and the rate of reaction increases.
• The Boltzmann distribution curve flattens more and shifts peak towards higher energy.
• The area under the curve remains the same and represents the total number of particles.

Catalysts

• Catalysts lower the activation energy; the point marked Ea on the curve shifts to the left.
• Hence, more particles have a higher energy than Ea, increasing reaction's rate.

Homogeneous and Heterogeneous Catalysts

• Homogeneous catalysts are catalysts that are in the same state as the reactants.
• Heterogeneous catalysts are in a different state than the reactants.

Inorganic Chemistry - Period 3

Properties of Elements of the 3rd Period

• Atomic Radius: Decreases across the period. Greater attraction due to increased nuclear charge/ number of electrons, results in smaller atomic radius. Smaller atomic radius and more valence electrons means stronger metallic bonding.
• Ionic Radius: Ionic radius decreases when nuclear charge increases (for isoelectronic species).
• Melting Point & Boiling Point: MP/BP increases from left to right, reaches maximum at silicon, then sharply decreases to minimum at Argon.
• Na, Mg, and Al are metallic elements forming giant ionic lattices; metallic bonding strength increases from Na to Al due to valence electrons.
• Si has a giant covalent structure (network of strong covalent bonds) and the highest melting point.
• S8 molecules have stronger Vander Waals force, resulting in hight MP compared to P4.
• Electrical Conductivity: Metals are conductors due to free valence electrons. Metal electrical conductivity depends on valence electrons and increases from from Na to Al. From Si onward, are nonmetals and non-conductors.

Reactions of Period 3 elements with water or steam

• Metal(s) + Water (I) → Metal Hydroxide(aq) + H2(g)
• Metal(s) + Steam (g) → Metal Oxide(s) + H2(g)
• Na reacts with cold water vigorously to form NaOH and H2 and Na dissolves quickly
• Mg reacts with cold water extremely slowly, solution formed is weak alkaline
• Mg reacts with steam vigorously to make MgO and H2, white solid left behind
• Cl will undergo disproportionation reaction to give acidic solution

Reactions of Period 3 elements with Oxygen

• Na: 4Na(s) + O2(g) → 2Na2O(s) heat at a vigorous rate with a bright yellow flame and White solid Product
• Mg: 2Mg(s) + O2(g) → 2MgO(s) heat at a vigorous rate with a bright white flame and White solid Product
• Al: 4Al(s) + 3O2(g) → 2Al2O3(s) uses powdered material at a fast rate with a bright white flame and White solid Product
• Si: Si(s) + O2(g) → SiO2(s) uses powdered material + heat at a slow rate with a bright white flame and White solid Product
• P: 4P(s) + 5O2(g) → P4O10(s) heat at a vigorous rate with a yellow/white flame and white clouds Product
• S: S(s) + O2(g) → SO2(g) uses powdered + heat at a gentle rate with a blue flame at Toxic fumes Product

Reactions of Period 3 elements with Chlorine

• Na: 2Na(s) + Cl2(g) → 2NaCl(s) heat, vigorous