Temperature, Units and Moles

Temperature Scales

• There are three common temperature scales: Celsius (°C), Fahrenheit (°F), and Kelvin (K).
• Formula to convert Celsius to Kelvin: °C + 273.15
• Formula to convert Celsius to Fahrenheit: (9/5)°C + 32
• Formula to convert Fahrenheit to Celsius: (5/9)(°F - 32)

Base Physical Quantities and Their Units

• Length is measured in metres (m).
• Mass is measured in kilograms (kg).
• Time is measured in seconds (s).
• Electric current is measured in amperes (A).
• Temperature is measured in kelvin (K).
• Amount of substance is measured in moles (mol).
• Luminous intensity is measured in candelas (cd).

Mole Concept

• A mole is Avogadro's number (6.02 x 10^23) of particles.
• Formula for calculating number of moles (n) = mass of substance (W) / molar mass (M)
• Formula for calculating number of moles (n) = number of atoms/molecules / Avogadro's number
• Formula for calculating number of moles (n) (for gas only) = volume in litres at STP / 22.4

Dalton's Atomic Theory

• Matter consists of indivisible particles called atoms.
• All atoms of a specific element are identical.
• Atoms of different elements are different.
• Atoms are the fundamental units involved in chemical reactions.
• Atoms unite in simple whole-number ratios to form compound atoms.
• Atoms are indestructible and cannot be created or destroyed.

Limiting Reagent

• Limiting reagent is the substance fully consumed during a chemical reaction.

Atomic Mass Unit (amu)

• 1 amu = 1.66 x 10^-24 g = 1.66 x 10^-27 kg

Laws of Chemical Combination

• Law of conservation of mass (Lavoisier): Total mass of reactants equals total mass of products.
• Law of definite proportions (Proust): A chemical compound contains the same elements in the same proportions by mass.
• Law of multiple proportions (Dalton): When two elements form multiple compounds, the masses of one element that combine with the fixed mass of the other are in a ratio of small whole numbers.
• Gay-Lussac's Law of Gaseous Volumes: When gases react, the volumes of reactants and products are in a simple whole number ratio at constant temperature and pressure.
• Avogadro's Law: Equal volumes of all gases under the same conditions contain equal number of molecules.

Percentage Composition

• Formula for calculating mass percentage of an element: (Mass of that element in the compound / molar mass of the compound) * 100

Empirical and Molecular Formulas

• Empirical formula: The simplest whole number ratio of atoms in a compound.
• Molecular formula: The exact number of different types of atoms in a molecule.
• Formula: Molecular formula = Empirical formula x n
• Formula for Calculating n: n = molecular mass / empirical formula mass

Reactions in Solutions

• Mass percent: (Mass of solute / mass of solution) * 100
• Mole fraction (χ): Ratio of moles of a component to total moles in solution
• Mole fraction of A: nA/(nA + nB)
• Mole fraction of B: nB/(nA + nB)
• Molarity (M): Moles of solute per litre of solution
• Formula: Molarity = No. of moles of solute / Volume of solution in litres
• Molality (m): Moles of solute per kilogram of solvent
• Formula: Molality = No. of moles of solute / Mass of solvent in kg
• Molarity is temperature-dependent, molality is temperature independent.

Structure of Atom: Cathode and Anode Rays

• Cathode rays (electrons) travel in straight lines, consist of negatively charged particles, and their nature doesn't depend on the gas or cathode material used.
• Charge of an electron is 1.602 x 10^-19 C from the Millikan oil drop experiment.
• Mass of an electron is 9.10 x 10^-31 kg.
• Anode/canal rays travel in straight lines, their nature depends on the gas, consist of positively charged particles, and the smallest one from hydrogen is called a proton.
• Discovery of the neutron by Chadwick: 4Be + 2He -> 6C + 0n

Models of Atom

• Thomson’s model: Plum pudding or watermelon model.
• Rutherford’s nuclear model: Used alpha-ray scattering to bombard a thin gold sheet.
• Observations:
  • Most alpha particles passed through undeflected.
  • Some deflected at small angles.
  • Few deflected back.
• Conclusions:
  • Atom is mostly empty space.
  • A heavy, positively charged core (nucleus) exists.
  • Nucleus volume is very small compared to the atom.

Rutherford’s Nuclear Model of the Atom

• Most mass and positive charge are in a small region called the nucleus.
• Electrons move around the nucleus at high speed.
• Electrostatic force of attraction holds electrons and nucleus together.
• Rutherford's model failed because it couldn't explain atom stability and hydrogen spectrum using Maxwell's theory of electromagnetic radiation.

Composition of Atom

• Atomic number (z): Number of protons = number of electrons.
• Mass number (A): Number of protons + number of neutrons.
• Number of neutrons is mass number minus atomic number.

Electromagnetic Waves

• Wavelength (λ): The distance between two consecutive crests or troughs, measured in nanometers (10^-9 m), angstroms (10^-10 m), or picometers (10^-12 m).
• Frequency (ν): Number of wavelengths passing a point in one second (Hz).
• Wave number (ν̄): Number of wavelengths per centimeter or meter (cm^-1 or m^-1).
• Relationship between speed of light (c), frequency (ν), and wavelength (λ): c = νλ
• Electromagnetic Spectrum's arrangement by increasing wavelength and decreasing frequency: Cosmic rays < γ-rays < Ultra violet < visible < infrared < microwaves < radiowaves.

Planck's Quantum Theory

• Radiant energy is emitted/absorbed discontinuously as quanta or photons.
• Quantum energy is directly proportional to frequency: E = hν = hc/λ.
• Plank's constant (h) = 6.626 x 10^-34 Joules sec.
• Total energy emitted/absorbed is an integer multiple of quantum: E = nhν.

Black Body Radiation

• A perfect absorber and emitter of radiation across all frequencies.

Photoelectric Effect

• Electrons ejected from a metal surface when light of suitable frequency strikes it.
• Electrons are ejected instantly.
• Number of electrons is proportional to light intensity.
• Threshold frequency (ν0) exists below which no effect occurs.
• Kinetic energy of ejected electrons is proportional to incident light frequency.

Photoelectric Energy Equation

• Energy of Incident Photon = Work Function + Kinetic Energy
• Equation: hν = hν0+ 1/2mev²

Atomic Spectrum of Hydrogen

• Rydberg Formula: 1/λ = RH(1/n1² - 1/n2²).
• RH = Rydberg's constant = 109677 cm⁻¹.
• n1 and n2 are integers where n2 > n1.
• Lyman series: n1 = 1, n2 = 2, 3, 4, 5...
• Balmer series: n1 = 2, n2 = 3, 4, 5, 6...
• Paschen series: n1 = 3, n2 = 4, 5, 6, 7...
• Brackett series: n1 = 4, n2 = 5, 6, 7, 8...
• Pfund series: n1 = 5, n2 = 6, 7, 8, 9...

Bohr Model of Atom

• Electrons revolve around nucleus in definite orbits.
• Orbits have specific energy levels (energy shells).
• Shells are numbered 1, 2, 3, 4... or K, L, M, N...
• Permitted orbitals have angular momentum as a whole number multiple of h/2π.
• Angular momentum of electron = nh/2π (n = 1, 2, 3, 4...).
• Electrons neither absorb nor lose energy while present in a certain orbit.
• Electrons absorb energy to jump to higher energy states.

Bohr's Model Merits

• Could explain atom stability
• Could calculate electron energy in hydrogen-like species (He+, Li2+): En = -2.18 x 10^-18 x Z^2/n^2 J/atom
• Could explain the atomic spectrum of hydrogen

Bohr's Theory Limitations

• Could not explain atoms with more than one electron.
• Failed to explain the fine spectrum.
• Could not explain Zeeman and Stark effects.

Dual Nature of Radiations

• All material particles exhibit dual behavior, acting as both waves and particles; de Broglie Equation = λ = h/mv = h/p

Heisenberg's Uncertainty Principle

• It is impossible to determine the position and momentum of microscopic particles precisely and simultaneously.
• Formula: Δx⋅Δp > h/4π

Schrödinger Wave Equation

• Simplified form: Ĥψ = Eψ
• ψ = wave function, E = permissible total energy of electron, Ĥ = Hamiltonian operator

Quantum Numbers

• Atomic electron address is given by numbers.
• Principle quantum number (n): Determines the shell/orbit and its size
• Azimuthal quantum number (l): Describes the subshells and orbital shapes, values are 0 to (n-1).
• Magnetic quantum number (m): Describes the orbital and its orientation.

Pauli Exclusion Principle

• No two electrons in an atom can possess the same values for all four quantum numbers
• Therefore, an orbital cannot have more than two electrons.
• Electrons are filled in different orbitals while increasing their energy

Hund's Rule of Maximum Multiplicity

• Pairing of electrons in degenerate orbitals happens only once all orbitals are singly occupied.

Shapes of Atomic Orbitals

• s orbital - spherical
• p orbital - dumb-bell

Nodal Surfaces

• Regions where probability to find an electron is zero.
• Number of radial nodes: n - l - 1
• Number of angular nodes: l
• Total number of nodes: n - 1

Electronic Configurations of Chromium and Copper

• Chromium (Cr, Z=24): [Ar] 3d5 4s1 (due to half-filled stability of 3d orbitals).
• Copper (Cu, Z=29): [Ar] 3d10 4s1 (due to fully filled stability of 3d orbitals).

Dobereiner’s Triads

• Certain similar elements exist in groups of three elements called triads.
• Elements are arranged in order of increasing atomic mass; the atomic mass of the middle elements is the arithmetic mean of the atomic weights of the other two elements in the triad.
• Exmaples; Li, Na, K, and Ca, Sr, Ba and Cl, Br, I

Newland’s Law of Octaves

• Elements are arranged in the order of increasing atomic weights, and every eighth element has similar properties.

Mendeleev’s Periodic Law

• The properties of the elements are a periodic function of their atomic weights.

Defects in Mendeleev's Periodic Table

• Position of hydrogen is not justified.
• Isotopes are placed in the same position in the table.

Mendeleev’s periodic table, merits

• It was the first comprehensive classification of elements.
• He corrected the wrong atomic weights of some elements and placed them in correct position in the periodic table.
• He left vacant places for undiscovered elements and predicted some of their properties. Eka-Aluminium(=Gallium) and Eka-Silicon(=Germanium)

Modern Periodic Law

• The physical and chemical properties of the elements are periodic functions of their atomic numbers. (by Moseley)

IUPAC Nomenclature

• IUPAC Nomenclature for Elements with Atomic Number >100
• s-Block Elements
  • Group 1 (alkali metals) - ns¹ (outermost electronic configuration)
  • Group 2 (alkaline earth metals) - ns² (outermost electronic configuration)
  • Alkali metals form +1 ion and alkaline earth metals form +2 ion.
  • They are never found in pure state in nature. (Reason - they are highly reactive)
• p-Block Elements
  • Elements belonging to Groups 13 to 18
  • Outermost electronic configuration varies from ns²np¹ to ns²np⁶ in each period.
  • Group 18 (ns²np⁶) - noble gases
  • Group 17 - halogen
  • Group 16 - chalcogens
  • Non-metallic character increases from left to right across a period.
• d-Block Elements (Transition Elements)
  • Elements of group 3 to group 12
  • General electronic configuration is (n − 1) d¹⁻¹⁰ ns⁰⁻².
  • Called transition elements
  • All are metals. They form coloured ions, exhibit variable oxidation states, paramagnetism, and are used as catalysts.
• f-Block Elements
  • Lanthanoids → Ce (Z = 58) to Lu (Z = 71)
  • Actinoids → Th (Z = 90) to Lr (Z = 103)
  • Outer electronic configuration → (n - 2)f¹⁻¹⁴ (n-1)d⁰⁻¹ns²
  • They are called inner-transition elements.
• All are metals.
• Elements after uranium are called Transuranium elements.

Periodic Trends

• Atomic radius diminishes while expanding the atomic number in a period.
• An example include, atomic radii decrease from Li to F in the second period.
• In the move, the atomic number expands progressively and in a single unit. Therefore, the electron cloud is pulled closer to the core by the effective nuclear charge and decreases the atomic size.
• Atomic radius increases from peak to bottom within a group of the periodic table.

Ionic Radius

• A cation is diminished than its parent atom. This is because a cation has few electrons, but the nuclear charge lives same as compared to the parent atom. For example, Nat is diminished than Na atom.
• The anion size is better than its parent atom. As an example, Cl− is better than Cl atom. This is because the addition of one or more electrons will give an expanded electronic repulsion and reduce the effective nuclear charge.

Isoelectronic Species

• Atoms and ions that possess the same amounts of electrons will be addressed as isoelectronic species. For example, O2-, F-, Ne, Na+, Mg2+. Also, all possess 10 electrons.

Ionization Enthalpy

• It's stated as the capacity needed to eradicate an electron from the outer shell of an solitary gaseous atom in its ground state X(g) + ΔiH → X⁺ (g) + e⁻

Factors Affecting Ionization Enthalpy

• Important influences that affect ionization enthalpy includes:
  • Atomic size: Greater the atomic size, the reduced is the ionization enthalpy
  • Nuclear charge: The value of ionization enthalpy expands with nuclear charge.
  • Shielding effect: Because the shielding effect expands, the electrons could be eliminated easily and the ionization enthalpy diminishes.
  • Presence of half-filled or completely filled orbitals expands ionization enthalpy. Along a time frame, ionization enthalpy expands from left to right. That's the decline in atomic radius and development in nuclear charge. Therefore alkali metals bear the least AiH and noble gases possess the most. Down a group, AiH diminishes to expand in atomic radius and shielding effect. Thus, among alkali metals, lithium possesses the least ∆iH, and francium possesses the most.

Electron Gain Enthalpy

• It is stated as the enthalpy change coming into place as an separated gaseous atom acknowledges an electron to create a monovalent gaseous anion X(g)+℮ → X¯(g)
• Larger the value of electron gain enthalpy, greater is the tendency of an atom to admit an electron.
• Variation across time frame leads to be more negative as to go from left to right across a time frame
• Variation down a group lead to be less negative on going down the group

Electron Gain Enthalpy

• The electron gain enthalpy of fluorine exists in less negative than chlorine because an electron is attached to F, it enters into the diminished 2nd shell. Therefore, the electrons endure to be more repulsion from the some electrons. But by Cl, the incoming electron goes to the better 3rd shell. Therefore, the electrons endure low repulsion, and Cl admits electron more easily than F

Electronegativity

• Stated as the inclination of an atom in a molecule to capture the paired pairs of electrons near itself.
• Greater the practical nuclear charge, the better is electronegativity.
• Diminished the atomic radius, the better is the electronegativity.
• In a period is expanded in expanding to move from left to right.
• In a group is diminishes in moving down a group.

Lewis Symbols

• Each element contains its symbol, containing number of valence electrons. H., He.Li., Be., B· , C·, N:, O:, F:, Ne:

Structure of Lewis

• Every bond and electron configuration that shows how atoms are structured and bonded Cl-Cl; O=O; NitripleN; O=C=O

VSEPR Theory

• The shape of a molecule depends upon the valence shell electron pairs around the centrel atom
• Pairs of electrons in valence shell repel one another since their electron clouds is negatively charged
• To reduce repulsion, the electron pairs stay very far apart as possible and multiple bond is treated as single super pair.
• The repulsive relations of decreasing electron pairs in the stated order: lp-lp>lp-bp>bp-bp

Chemical Bonding and Molecular Structure

• Shape of NH3 is five electrons in N, and bonded to three hydrogen atoms making the three electrons
• Shape of H2O two electrons in H, and two electron pairs
• Formsigma bonds axial overlapping of partially filled atomic orbitals and inter nuclear axis
• Form Pi bonds through sidewise overlapping
• SP3 hybridization , one S-Orbital mix with three P-Orbitals to form SP3 hybridization
• In SP3, direction is four Sp3 hybrised to the corner of tetrahedrons

Sp2 Hybridization

• Sp2 hybridization is when one S-Orbital combines with to P-Orbitals, and there is 3 sp2 hybrid orbital directions to the corners and corners of Triangles.
• In SP hybridization, combination of one S-orbital with one P-orbital forms two sp hybridized orbital and are directionally oposite to each other (liner).
• In hybridization for one S-Orbital with three P-Orbitals is where the one (D-Orbital) forms five sp3d hybridized Orbitals and this hybridation has direct corners to Triangle.

MOT

These have Key Features with theory on what’s offered below:

• Molecules must have electrons and has orbitals in these theories.
• Molecules must have (symmetry in their) and proper energies which must be conbimed to form a orbit
• Molecules that should have polycentric orbitals.
• Molecular orbitals are having 2 keys (bonding orbital) and it must be (anti-bonding orbit).
• Stable bonds (low energy) and while antibonding orbits have high and low of energy

Bond Orders (b.o)

• It has a limited difference between the amount of present electrons (the number) in bonding and anti-bonding orbit

Applications for Bond Order

• Its positive so Its stable and negative then you cant use unstable and molecule exist
• Bond Order is directly relational for the stability of its molecule.

Hydro Bonding

• Between Molecules and atoms and also molecule that forms and has electronegative
• Types of Bonds includes Intermolecular(the form between the 2 molecules is intermolecular).
• Also Intramolecular( is made within a molecule it Is Intramolecular)
• The H2O is and H2S is because that it contains intermolecular then it will be in the higher temperature
• This is a due for the inter bonding having higher point over bonding.