\documentclass[12pt]{report} \setcounter{tocdepth}{5} \setcounter{secnumdepth}{5} \usepackage{geometry} \geometry{ a4paper, total={170mm,257mm}, left=20mm, top=20mm, } \setlength{\footskip}{45pt} \renewcommand{\familydefault}{\sfdefault} \usepackage{euler} \usepackage{parskip} \usepackage{graphicx} \usepackage{amsmath} \everymath{\displaystyle} \begin{document} \tableofcontents \chapter{Chemistry} \section{Basic Chemistry} \subsection{Laws of Chemical Combination} \subsubsection{Lavoisier’s Law} “that mass can neither be created nor destroyed, although it may be rearranged in space, or the entities associated with it may be changed in form” \subsubsection{Proust’s Law} “given chemical compound always contains its component elements in fixed ratio (by mass) and does not depend on its source and method of preparation” \subsubsection{Dalton’s Law} “if two elements form more than one compound between them, then the ratios of the masses of the second element which combine with a fixed mass of the first element will always be ratios of small whole numbers” \subsubsection{Gay Lusaac’s Law} “when gases react together they do so in volume which bears simple whole number ratio provided that the temperature and pressure of the reacting gases and their products remain constant” \subsubsection{Dalton’s Atomic Theory} “matter consists of indivisible atom”\par “compounds form when atoms combine”\par “atoms of different element differ by mass” \subsubsection{Avogadro’s Law} “equal volumes of all gases, at the same temperature and pressure, have the same number of atoms or molecules” \subsection{Molar Theory} \subsubsection{Masses} \begin{itemize} \item Atomic Weight \item Molecular Weight \item Equivalent Weight \begin{enumerate} \item Elements:\par combine with or displace from a compound, 1.008 grams of hydrogen or its equivalent.\par [equivalents= 8 grams of oxygen or 35.5 grams of chlorine] \item Acids and Bases: \par contains 1.008 grams of replaceable hydrogen ions or neutralizes EW grams of an acid. \par \item Oxidizing and Reducing Agents:\par supplies or reacts with one mole of electrons. \end{enumerate} \end{itemize} \subsubsection{Mole Concept} \subsubsection{Solution Concentration} \begin{enumerate} \item Mass Percent \item Molality \item MGram Per Litre \item Molarity \item Normality
\end{enumerate}
\section{Structure of Atoms}
\subsection{Initial Theories}
\subsubsection{Thompson’s Model} \ldots
\subsubsection{Rutherford’s Model} \ldots
\subsubsection{Bohr’s Model} \ldots
\subsection{Quantum Model}
\includegraphics[width=\textwidth]{schr.jpg}
{\bfseries Quantum Numbers:}
\begin{enumerate}
\item Principal Quantum Number (shell, size and energy):[0 to ][no: 2n]
\item Azimuthal Quantum Number (sub-shell, shape and momentum):[0 to (n-1)][no: 2(2l+1)]
\item Magnetic Quantum Number (orbital, orientation):[-l to l][no: 2]
\item Spin Quantum Number (spin):[][no: 1]
\end{enumerate}
{\bfseries Rules For Filling Electrons:}
\begin{enumerate}
\item Afbau Principle: Electrons are filled up on atomic orbitals on the order of their energies. \par
Exception:
\item (n+l) rule: The shell with greater (n+l) is filled first.
\item Hund’s Maximum Multiplicity: The orbitals are first filled singly with parallel spin electrons.
\item Pauli’s Exclusion Principle: No two electrons in an atom can have same set of quantum numbers.
\end{enumerate}
\section{Periodic Table}
\includegraphics[width=0.95\textwidth]{vv.png}
\subsection{Periodic Trends}
\begin{enumerate}
\item Atomic Radius [nuclear charge, screening effects] (, )
\item Ionization Enthalpy, Electron Gain Enthalpy, Electronegativity (, )
\end{enumerate}
\begin{itemize}
\item Application Of Ionization Enthalpy:
\begin{enumerate}
\item Metallic Character
\item ie 16eV, lower state is stable; ie 11eV, higher state is stable
\end{enumerate}
\item Application Of Electronegativity:
\begin{enumerate}
\item Non-Metallic Character
\item Bond Energy en
\end{enumerate}
\end{itemize}
\subsection{s-block}
\begin{enumerate}
\item Physical: Colorless
\item Air: tarnish in air due forming oxides which in turn reacts with mositures to form hydroxides.
\item Water: burn in water giving hydrodixes and dihydrogen.
\item Reducer: strong reducting agents
\end{enumerate}
\subsection{p-block}
\begin{enumerate}
\item Physical: Colorless
\item Oxidizer: good oxidizing agents
\end{enumerate}
\subsection{d-block}
\begin{enumerate}
\item Physical:
\begin{itemize}
\item Color Forming [in ionic or molecules, d orbitals are split (CFT), energy is in visible range]
\item Alloy Formation [comparable atomic sizes, so atoms replace each other]
\item Magnetic Properties: [presence of unpaired electrons in d orbitals]
\end{itemize}
\item Disperancies: \par Atomic Radius ($\leftarrow$, $\rightarrow$) [dominating screening effect] \par Ionization Energies ($\rightarrow,\leftarrow,\rightarrow,\leftarrow$) [shielding to electrons]
\item Variable Oxidation States: [participation of s and d orbitals as d core is unstable]
\item Complex Forming: [vacant d-orbitals for accommodation]
\item Catalysation: [form unstable intermediate due to variable oxidations.]
\end{enumerate}
\subsection{General Trends:}
\begin{enumerate}
\item Thermal Decomposition: \par
For having polyatomic ions CO$_3^{-2}$, SO$_4^{-2}$, OH$^{-}$, O$_2^{-2}$, O$_2^{-}$,
$$\text{Thermal Stability}\propto \frac{\text{size of cation}}{\text{charge of cation}}$$
For having hallide, hydride or normal oxide,
$$\text{Thermal Stability}\propto \frac{1}{\text{size}} ~\text{[for a group]},~\propto \text{en [for a period]}$$
Note: \par LiHCO$_3$ and IIA group do not exist in solid state.\par
Carbonates, sulphates and hydroxides of Na, K, Rb and Cs do not decompose. \par
BeCO$_3$ kept in CO$_2$ due to less thermal stability.
\item Heating Effect:\par
\begin{center}
Metal Carbonates $\rightarrow$ Metal Oxide + Carbondioxide Gas\par
Metal Hydroxide $\rightarrow$ Metal Oxide + Water\par
Metal Bicarbonate $\rightarrow$ Metal Oxide + Carbondioxide Gas+ Water\par
Metal Nitrate $\rightarrow$ Metal Oxide+ NO$_2$+O$_2$\par
(except: Na, K, Rb, Cs [low temp]$\rightarrow$ MNO$_2$+$\frac{1}{2}$O$_2$ [high temp]$\rightarrow$ M$_2$O+N$_2$+O$_2$)
\end{center}
\begin{enumerate}
\item NH$_4$ salts having CO$_3^{-2}$, PO$_4^{-2}$, SO$_4^{-2}$, X$^{-}$ [weak oxidizing] give NH$_3$ gas.
\item NH$_4$ salts having Cr$_2$O$_7^{-2}$, ClO$_3^{-}$, NO$_2^{-}$, NO$_3^{-}$ [strong oxidizing] give N$_2$O or N$_2$ gas.
\item Ag$_2$O and HgO further decompose into metal and oxygen gas.
\item Metal salts with high percentage of O$_2$, KMnO$_4$, K$_2$Cr$_2$O$_7$ give O$_2$ gas.
\end{enumerate}
\item Color Of Compounds:
\begin{center}
Color Of Intensity $\propto$ Covalent Character
\end{center}
\end{enumerate}
\section{Chemical Bonding}
\subsection{Intrachemical Bonds}
\subsubsection{Ionic Bonds}
[shared electron is localized to an atom due to electronegativity differences]
\subsubsection{Covalent Bonds}
[two atoms share valence electrons more or less equally]
\subsubsection{Coordinate Covalent Bonds}
[pair of shared bonding electrons are from the same one of the atoms involved in the bond]
\subsubsection{Metallic Bonds}
[bonding electrons are delocalized over a lattice of atoms]
\subsubsection{Other Covalent Bonds}
[includes one-three electron bonds, three center two electrons, four center three electrons]
\subsection{Interchemical Bonds}
[bonds between two or more (otherwise non-associated) molecules, ions or atoms due to dipoles]
\subsection{Bonding Theories}
\subsubsection{Lewis Model}
[uses initial theories to describe bonding and molecules]
\subsubsection{Valence Bond Theory}
[describes a covalent bond between two atoms by the overlap [] of half filled valence atomic orbitals of each atom containing one unpaired electron, complements molecular bond theory]\par
\begin{itemize}
\item Hybridization: mixing atomic orbitals into new hybrid orbitals (with different energies, shapes, etc, than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds.
\begin{center}
\begin{tabular}{|c|c|}\hline
Hybridisation & Shapes \\hline
sp & Linear (180)\
sp & Trigonal (120)\
sp & Tetrahedral (109.5)\
dsp & Square Planar\
spd & Trigonal Bipyramidal\
dsp & Square Pyramidal\
dsp/spd & Octahedral\\hline
\end{tabular}
\end{center}
\item Isovalent Hybridization: if lone pairs involve in hybridisation, shape of hybrid molecules deviates from the obvious.
\item Crystal Field Splitting: breaking of degeneracies of electron orbital states, usually d or f orbitals, due to a static electric field produced by a surrounding charge distribution
\end{itemize}
\subsubsection{Molecular Bond Theory}\ldots
\section{Properties of Reactions} \subsection{Equilibrium} Chemical equilibrium is the state in which both reactants and products are present in concentrations which have no further tendency to change with time, so that there is no observable change in the properties of the system. The equilibrium concentration position of a reaction is said to lie far to the right if, at equilibrium, nearly all the reactants are consumed. Conversely the equilibrium position is said to be far to the left if hardly any product is formed from the reactant. The law of mass action for concerted single step reactions, where [*] is active mass, is equivalent to concentration while constant for solids. The ratios, called equilibrium constant is indeed a constant, independent of the activities of the various species involved, though it does depend on temperature.\par {\bfseries Le-Chatelier’s Principle: }\par “If a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to partially reverse the change.”\par [is like a seesaw that counters the powers of a kid on one side by moving to another] \begin{enumerate} \item Concentration: [moves to the kid whose weight is low] \item Temperature: [moves to the kid whose hands are less on fire] \item Volume: [moves to the kid who cant fit well] \end{enumerate} {\bfseries Special:}\par Ionization of weak binary electrolytes (Ostwald’s Dilution Law) \subsection{Kinetics} \subsubsection{ Factors affecting rate of reactions:} \begin{enumerate} \item Nature of reactants: [ionics are faster than covalents] \item Surface area of reactants: [more surface area, more chances of contact] \item Effect of concentration: [high concentration, likeliness of contact] \item Effect of temperature: [high temperature, active reactants] \item Effect of catalyst: [boosts the reactions] \end{enumerate} so, at a constant temperature, the rate of reaction depends upon the concentration of reactants i.e where, p and q are experimentally determined such that (p+q) is called order of he reaction while K is the rate constant whose value does not depend on the concentration of the reactants.\par {\bfseries First Order Reactions:} \subsubsection{ Activation Energy:}\par Activation energy is the magnitude of the energy barrier separating minima of the potential energy surface pertaining to the initial and final thermodynamic state. For a chemical reaction to proceed, the temperature of the system should be high enough such that there exists an appreciable number of molecules with translational energy equal to or greater than the activation energy. \begin{itemize} \item Cataylst: [provides alternate path for reaction by lowering activation energy] \item Temperature: [roughly increase in 10 degree doubles rate of reaction or] \end{itemize} \subsection{Electricity} \begin{center}\includegraphics[width=0.8\textwidth]{cell.jpg}\end{center} \subsubsection{Electrolytic Cell} [uses electrical energy to drive a non-spontaneous redox reaction]\par {\bfseries Faraday’s Laws} \begin{enumerate} \item First Law: the mass of substance deposited at an electrode in during electrolysis is directly proportional to the amount of charges passed through the solution i.e \item Second Law: when same amount of electricity is passed through the electrolytes in series, the mass of substance deposited at the electrodes is directly proportional to their equivalent weight. \end{enumerate} {\bfseries NOTE:}\par 1 Faraday = quantity of electricity that liberate 1 equivalent weight = 1 mole of e = 96500 C \subsubsection{Galvanic Cell} [derives electrical energy from spontaneous redox reactions]\par Salt Bridge: maintains electrical neutrality within the internal circuit, preventing the cell from rapidly running its reaction to equilibrium else the solution in one half cell would accumulate negative charge and the solution in the other half cell would accumulate positive charge as the reaction proceeded, quickly preventing further reaction.\par Cell Notation: {\bfseries Electrode Potential:}\par when a metal is in contact with a solution of its own ions, the metal shows tendency to go into the solution by losing electrons while the metal ions show tendency to get deposited as metal atoms, finally at equilibrium the metal electrode is either positively or negatively charged in each cases there develops a charged layer at metal-solution interface which attracts the oppositely charged layer hence forming a electrical double layer, such tendency is the measure of single electrode potential.\par {\bfseries Standard Potentials}\par The standard single electrode potential is measured with respect to a primary reference electrode of hydrogen gas at 1 atm in contact with 1M H ions at 25 C on platinum as electrode. The potential of each half cells are different due to which current flows from higher cathode to lower anode giving rise to the emf, {\bfseries Nerst Equation:}\par For a cell reaction, \subsection{Thermodynamics} \section{Chemical Reactions} \subsection{Basic Reactions} [synthesis, decomposition, displacements, physicals] \subsection{Acid-Base Reactions} \begin{center} \begin{tabular}{|c|c|c|} \hline Concept & Acids & Bases \[0.3cm] \hline &&\ Arrhenius&gives H on water&gives OH on water\[0.3cm] Bronsted Lowry& proton donar species & proton accepter species\[0.3cm] Lewis&elecron pair acceptor & electron pair donar\[0.3cm]\hline \end{tabular}\end{center}~\par
{\bfseries Relative Strength of Acids and Bases}
\begin{enumerate}
\item Same Concentration:\par [larger the value of K, the larger is the degree of ionization and vice versa]
\item The p-Values:
$$\text{pH}=-\log~[H^+], ~~\text{pOH}=-\log~[OH^-]$$
[in water] $$K=10^{-14}\rightarrow \text{pH+pOH}=14$$
\end{enumerate}
In Terms Of Nature:
\begin{enumerate}
\item acidic strength increases with increase in atomic size and electronegativity of atom with H.\par
[if acid contains oxygen]
\item acidic strength increases with decrease in atomic size and increase in ON of atom with oxygen.
\end{enumerate}
{\bfseries Buffer Solution:}\par
[solution that resists the change of pH when a small amount of strong acid or base is added to it]
\begin{itemize}
\item Acidic Buffer: [equimolar concentration of weak acid and its salt with strong base]
$$\text{pH}=\text{pK}_a+\log\frac{[\text{salt}]}{\text{[acid]}}$$
\item Basic Buffer: [equimolar concentration of weak base and its salt with strong acid]
$$\text{pOH}=\text{pK}_b+\log\frac{[\text{salt}]}{\text{[base]}}$$
\end{itemize}
\subsection{Redox Reactions}
\begin{enumerate}
\item Oxidation: [loss of electrons] addition of electronegative atom [Oxygen]
\item Reduction: [gain of electrons] removal of electropositive atom [Hyrdrogen]
\end{enumerate}
{\bfseries NOTE:} \par
Oxidation state of an atom is the charge of this atom after ionic approximation of its heteronuclear bonds such that the increase in oxidation state of an atom, through a chemical reaction, is known as an oxidation; a decrease in oxidation state is known as a reduction. The overall electrochemical reaction for a redox process requires a balancing of the component half-reactions for oxidation and reduction which in general for reactions in aqueous solution involves adding H$^+$, OH$^-$, H$_2$O, and electrons to compensate for the oxidation changes.
\subsection{Complex Formations}
[when simple salts are crystallised thus formed molecule either is a double salt that completely dissociates into simple ions or is a complex salt that does not in water]\par
[a coordination complex consists of a central atom or ion, which is usually metallic and is called the coordination centre, and a surrounding array of bound molecules or ions [coordination number], that are in turn known as ligands (monodentate, polydentate; chelates)]\par
{\bfseries Werner's Theory: } [describes the complex cobalt salts]\par
The metal atoms exihit two valencies:
\begin{enumerate}
\item Principle Valency: [ionizable, satisfied by anions, non directional]
\item Secondary Valency: [non-ionizable, satisfied by anions or neutrals, directional geometry, fixed]
\end{enumerate}
\end{document}