%% Some of this file was created by LyX 1.5.5, but not all of it. %%%% UPDATE: Feb 11 2009: basically done now, please report any %%%% obnoxious LyX-ities you find \input{../include.tex} \input{../units.tex} \begin{document} \title{IGCSE Chemistry Notes} \date{$2007$--$9$} \author{Edward Lilley} \maketitle \tableofcontents \listoftables \listoffigures \section{Introduction} \subsection{Methods of Purification} \begin{description} \item [Filtration] to remove solid particles (bigger than clay sized ($<\frac{1}{256}\mr{mm}$)). \item [Crystallisation] removes the solvent, to leave the solute. \item [Fractional Distillation] produces a number of substances from the original mixture (e.g. petroleum). \item [Sedimentation] allows an insoluble solid to separate out and sink to the bottom of a container \item [Centrifuging] a spinning motion increases the force of gravity that quickly separates a solid from a suspension. \item [Decanting] pouring off liquid (e.g. pouring off excess water from a pot of peas). \item [Magnetic Separation] a method for separating one solid (usually iron) from a mixture of solids, Very useful for separating Aluminium cans from steel cans. \end{description} \subsubsection{Purification of Substances} Materials that are produced to a high degree of purity give predictable results (e.g. drugs, concrete). Purity can be measured in a number of ways: \begin{itemize} \item melting point/boiling point \item chromatography \end{itemize} \subsection{Measurement} Purity is impossible without measurement. In science we use the MKS system of units. \begin{description} \item [Time] \begin{itemize} \item measured in seconds \item clock, stopwatch, sundial, computer \begin{description} \item[Digital] one step at a time \item[Analogue] continuously changing \end{description} \item atomic clocks \end{itemize} \item [Mass] \begin{itemize} \item measured in grams/kilograms \item scales (electronic), triple beam balance \end{itemize} \item [Distance] \begin{itemize} \item measured in metres \end{itemize} \item [Volume] \begin{itemize} \item cm$^{\text{3}}$, ml \item measuring cylinder, burettes, pipettes \end{itemize} \item [Density] \begin{itemize} \item g/cm$^{\text{3}}$ $(\frac{mass}{volume})$ \end{itemize} \end{description} \subsection{Water} \begin{table} \begin{tabular}{lll} \hline & Air & Water \\ \hline Chemistry & \ce{N2}, \ce{O2}, \ce{Ar}, \ce{CO2}, \ce{H2O} & \ce{H2O}, dissolved salts, \ce{O2}, \ce{CO2} \\ Density & low & high \\ \hline \end{tabular} \caption{Comparison of the chemistry of water and air.} \end{table} \paragraph{Testing for Water} Cobalt Chloride (CoCl$_{2}$) turns pink in the presence of water (it is blue when dry). \paragraph{Uses of Water} \begin{enumerate} \item Home \begin{itemize} \item drinking \item washing \item putting out fires \item garden watering \item cooking \item heating \end{itemize} \item Industry \begin{itemize} \item coolant \item lubricant \item dilutant \item additive \item solvent \item H.E.P. \end{itemize} \end{enumerate} \subsection{Air} \paragraph{Uses of Components of Air} \begin{description} \item [Nitrogen] 79\% of air, used as a foaming agent e.g. packaging. \item [Oxygen] 20\% of air, used in cutting torches (oxy-acetelene). \item [Carbon~Dioxide] $\ll$1\% used in pneumatics, for putting out fires, special effects, coolant (dry ice). \item [Argon] $\lll$1\%, in torch batteries and car headlights. \end{description} Note: The way to obtain any of these gases is by fractional distillation at low temperatures. \subsection{Carbon Dioxide} \begin{itemize} \item A gas at room temperature and pressure (RTP: 25\degc, 1 atm). \item Colourless and odourless. \item The standard test for \ce{CO2} is to pass it through limewater, and observer whether the limewater turns cloudy or not (if it does, then \ce{CO2} is present). \item Industrial uses are: \begin{itemize} \item As a bubble generator in foams (foaming agent) \item Carbonated drinks \item \ce{CO2} enrichment in greenhouses to promote growth \item Fire extinguisher \item Refrigerant (dry ice) \end{itemize} \item Sources of \ce{CO2} \begin{itemize} \item Respiration (cellular) \item Burning fuels (fossil fuels, bio-fuels, natural fires) \item Volcanoes \item Weathering of carbonate rocks. \end{itemize} \end{itemize} \subsection{States of Matter} \begin{description} \item [Solid] at a given temperature has a definite shape and volume. Heating causes this volume to expand and cooling causes contraction (shrinking). \item [Liquid] at a given temperature, a liquid has a set volume but takes on the shape of its container. \item [Gas] at a given temperature has neither a definite shape or a definite volume. \end{description} \begin{figure}[htb!] \scalebox{1}{\includegraphics{states-of-matter}} \caption{The 3 states of matter and associated processes.} \end{figure} \subsection{Kinetic Theory} Kinetic Theory (or Particle Theory) states that all matter is made up of particles. These particles vibrate with solids having the smallest amount of vibration and gases having the most. As materials are cooled, the particle vibration decreases and ceases at $-273.15^{\circ}$C ($0^{\circ}$K), absolute zero. \subsubsection{Evidence for Kinetic Theory} \begin{description} \item [Diffusion] diffusion is used as evidence of the particle theory because gas can be shown to diffuse at a much faster rate than diffusion in liquids e.g. perfume sprayed into one corner of a room will quickly spread through the room. Also, when Potassium Pomanganate (KMnO4) crystals diffusing in warm and cool water, respectively, are compared, it is shown that the KMnO4 diffuses considerably faster in the warm water. This supports kinetic theory in that the warmer particles have more energy, so they can diffuse faster. \item [Brownian Motion] the apparent random movement of small particles. First observed in the early nineteenth century when pollen grains were observed to move in a random pattern on a water droplet. It is explained by the collision of water molecules with the pollen in a random pattern of collisions. \end{description} \subsection{Gas Laws} \begin{itemize} \item There is a relationship between the temperature of a gas and its volume. Charles's law states that the volume of a gas is directly proportional to its temperature: \[V\propto T\] This can be rewritten to: \[\frac{V}{T}=constant\] \item There is an inverse relationship between gas volume and pressure. Boyle's law states that as the pressure on a gas increases its volume decreases: \[V\propto\frac{1}{P}\] This can be rewritten to: \[PV=constant\] \end{itemize} \subsubsection{Combined Gas Laws} \[\frac{P_{1}V_{1}}{T_{1}}=\frac{P_{2}V_{2}}{T_{2}}\] Combining the two laws produces this equation. If you change the Pressure (P) or the Volume (V), it is easy to calculate the new Temperature (T). $T_{1}$is old temperature, $T_{2}$is changed temperature. \section{Atomic Structure} \subsection{Structure of the Atom} \begin{itemize} \item The atom is thought to be made up of a central nucleus that contains protons, and usually neutrons. The nucleus is surrounded by a ``cloud'' of electrons. \begin{itemize} \item Protons (p$^{+}$) have a positive charge. \item Neutrons (n$^{0}$) have no charge. \item Electrons (e$^{-}$) have a negative charge. \end{itemize} \item The atomic number of an element defines which element you are looking at. This is determined by the number of protons (this is also the number of electrons). The atomic mass is the sum of the number of protons + neutrons in the nucleus. \item It is possible to summarize this data using symbols. e.g. \begin{itemize} \item $_{1}^{1}$H = Hydrogen with 1p$^{+}$ and 1e$^{-}$ and no n$^{0}$. \item $_{6}^{12}$C = Carbon with 6p$^{+}$ and 6e$^{-}$ and 6n$^{0}$. \end{itemize} \item The electrons and the protons balance each other so that an atom has no overall electric charge. The electrons are packed around the nucleus in a series of layers or shells. \item The electrons fill up the shells progressively with: \begin{itemize} \item 2e$^{-}$ in the first shell \item 8e$^{-}$ in the second \item 8e$^{-}$ in the third \item 18e$^{-}$ in the fourth. \end{itemize} \end{itemize} \subsection{Sub-atomic Particles and Mass} The neutron is the heaviest of the non-fundamental sub-atomic particles. It has a mass of about $1.675\times10^{27}kg$, slightly more than the proton. It has no electronic charge, is located in the nucleus of the atom, and is written as $n^{0}$. The next heaviest is the proton, with a mass of about $1.673\times10^{27}kg$. It is located in the nucleus of the atom, and has a positive electric charge of $1.602\times10^{19}C$ (Coulombs). It is written as $P^{+}$. The smallest sub-atomic particle (that is part of the atom itself) is the electron, which is also a fundamental particle. It has a mass of $9.109\times10^{31}kg$ (i.e. several orders of magnitude less than both the neutron and the proton), and a negative electronic charge of $1.602\times10^{19}C$. It is written as $e^{-}$. \begin{sidewaystable} %\begin{table}[htb!] \begin{tabular}{|l||r|r|r|r|r|r|} \hline Element & Atomic Number (Z) & Mass Number (A) & No. of p$^{+}$ & No. of n$^{0}$ & No. of e$^{-}$ & Electron Configuration \\ \hline \hline H & 1 & 1 & 1 & 0 & 1 & 1 \\ \hline He & 2 & 4 & 2 & 2 & 2 & 2 \\ \hline Li & 3 & 7 & 3 & 4 & 3 & 2.1 \\ \hline Be & 4 & 9 & 4 & 5 & 4 & 2.2 \\ \hline B & 5 & 11 & 5 & 6 & 5 & 2.3 \\ \hline C & 6 & 12 & 6 & 6 & 6 & 2.4 \\ \hline N & 7 & 14 & 7 & 7 & 7 & 2.5 \\ \hline O & 8 & 16 & 8 & 8 & 8 & 2.6 \\ \hline F & 9 & 19 & 9 & 10 & 9 & 2.7 \\ \hline Ne & 10 & 20 & 10 & 10 & 10 & 2.8 \\ \hline \end{tabular} \caption{The electron configuration of the first 10 elements.} %\end{table} \end{sidewaystable} \begin{figure}[htb!] \includegraphics[scale=0.75]{diagram-of-atom} \caption{A diagram of a typical atom, with \ce{20n^0}, \ce{20p+} and \ce{20e-}.} \end{figure} $Z=20$ $A=40$ $e^{-}\approx1.602\times10^{-19}C$ $20\times e^{-}\approx-32.04\times10^{-19}C$ (Coulombs - measurement of charge). \subsection{Isotopes} An isotope is an element that has a different number neutrons in its nucleus. e.g. \begin{itemize} \item Hydrogen $_{1}^{1}$H \item Deuterium $_{1}^{2}$H \item Tritium $_{1}^{3}$H \end{itemize} The more common nomenclature (naming system) for isotopes gives the element name followed by a number that is the atomic mas number. e.g. \begin{itemize} \item C12 = $_{6}^{12}$C \item C13 = $_{6}^{13}$C \item C14 = $_{6}^{14}$C \end{itemize} Isotopes can be radioactive or non-radioactive, The factor that determines whether they are radioactive is if the neutron to proton ratio is higher than 1.5:1. \subsection{Uses of Isotopes} \paragraph{Medical} \begin{itemize} \item Medical isotopes are used for treatment using radiation (radiotherapy). \item Medical isotopes are used for imaging purposes (x-rays). \item Medical isotopes are used to highlight soft tissues that cannot be imaged by x-rays. \end{itemize} \paragraph{Industrial} \begin{itemize} \item Isotopes can be used for ``on-line'' control of processes such as mining, metallurgy, mineral processing and welding checks. \end{itemize} \paragraph{Scientific} \begin{itemize} \item The decay of isotopes can be used to work out the age of rocks, fossils and earth processes such as plat collisions and fold mountain formation. \end{itemize} \subsection{Electronic Structure} The number of electrons around an atom is the same as the number of protons, The electrons `pack' around the nucleus in a series of levels or shells. The strength of the electric charge on an electron means that they pack around the nucleus in a pattern. \begin{table}[htb!] \centering \begin{tabular}{lr} \hline Shell & No. of \ce{e-} \hfill \\ \hline 1 & 2 \\ 2 & 8 \\ 3 & 8 \\ 4 & 18 \\ 5 & 18 \\ 6 & 32 \\ \hline \end{tabular} \caption{Electronic structure for the first 6 shells.} \end{table} All elements with the same number of electrons in their outer shell (the valence electrons) will behave in a similar manner chemically. \subsection{Valencies} \begin{itemize} \item Each of the groups in the periodic table have a distinctive valency. \item The valency of an element (or radical) is the number of electrons available or electrons spaces that need to be filled. \end{itemize} \begin{table}[htb!] \centering \begin{tabular}{lrc} \hline Group & No. of e$^{-}$ in outer shell & Valency \\ \hline I & 1 & $1+$ \\ II & 2 & $2+$ \\ III & 3 & $3+$ \\ IV & 4 & $4\pm$ \\ V & 5 & $3-$ \\ VI & 6 & $2-$ \\ VII & 7 & $1-$ \\ 0 & 8 & $0$ \\ \hline \end{tabular} \caption{Group valencies.} \end{table} Valencies 1+ to 3+, and 3- to 1-, form ionic bonds. 4+/- forms covalent bonds, and 3- can form covalent forms sometimes, as well as ionic. \begin{itemize} \item For the transition metals, assume they have a valency of +2 with the notable exceptions being Fe(III) which has a valency of +3, and Hg(I) which has a valency of +1. \item Radicals behave as if they were single atoms. \end{itemize} \begin{table} \centering \begin{tabular}{lll} \hline Radical Name & Formula & Valency \\ \hline Sulphate & \ce{SO4} & $-2$ \\ Nitrate & \ce{No3} & $-1$ \\ Carbonate & \ce{CO3} & $-2$ \\ Phosphate & \ce{PO4} & $-3$ \\ Hydrogen Carbonate & \ce{HCO3} & $-1$ \\ Hydroxide & \ce{OH} & $-1$ \\ Ammonium & \ce{NH4} & $+1$ \\ \hline \end{tabular} \caption{Valencies of common radicals.} \end{table} \section{Periodic Table \& Properties of Elements} \subsection{What is the Periodic Table?} It is a summary of all the known elements. The elements are arranged in ascending order according to atomic number (Z). The vertical columns are called groups and contain elements that have similar chemical properties (dictated by the electron arrangement of the outer shell). The horizontal rows, or periods, show changing chemical properties across a series of elements as their outer electron shell fills up. \subsection{History of the Periodic Table} In 1817 Johann D\"{o}bereiner noticed that the atomic mass of Strontium fell between that of Calcium and Barium, and all three had similar chemical properties. He called this a ``triad'' of elements and found three triads altogether: \begin{itemize} \item Calcium, Strontium, Barium \item Chlorine, Bromine, Iodine \item Lithium, Sodium, Potassium \end{itemize} In 1865, John Newlands listed all of the known elements by atomic mass and noted that every eighth element had similar chemical properties. He called this the ``Law of Octaves''. \subsection{Group I -- The Alkali Metals} ($\downarrow$progressively more vigorous) \[2Li_{(s)}+2H_{2}O_{(l)}\longrightarrow2LiOH_{(aq)}+H_{2(g)}\] \[2Na_{(s)}+2H_{2}O_{(l)}\longrightarrow2NaOH_{(aq)}+H_{2(g)}\] \[2K_{(s)}+2H_{2}O_{(l)}\longrightarrow2KOH_{(aq)}+H_{2(g)}\] Each of these elements (Li, Na, K) have the same outer shell of electrons (only 1e$^{-}$). This means that they combine with water in the same way; two atoms of the elements in question will combine with two molecules of water to form two molecules of the alkali salt of that element and hydrogen gas. Each reaction is progressively more vigorous (the reaction takes place in water). This group is known as the alkali metals. All of these metals react in a similar way because they have a similar electron configuration i.e. they have a single electron in their outer shell. The reason for the increasing activity of the Group I metals down the group is due to the position of the outer electron, As you go further down the group, this single outer electron has a smaller attractive force to the protons in the nucleus. e.g. \begin{figure}[htb!] \centering \includegraphics[scale=0.8]{diagram-of-lithium-atom} \caption{A diagram of a Lithium atom (Li).} \end{figure} \begin{figure}[htb!] \centering \includegraphics[scale=0.8]{diagram-of-sodium-atom} \caption{A diagram of a Sodium atom (Na).} \end{figure} \begin{figure}[htb!] \centering \includegraphics[scale=0.8]{diagram-of-potassium-atom} \caption{A diagram of a Potassium atom (K).} \end{figure} \subsection{Group VII -- The Halogens} \begin{itemize} \item They are all reactive non-metals. \item Fluorine is the most reactive while Iodine is the most stable i.e. stability increases down the group. \item They are all coloured, with intensity of colour increasing down the group. \item Their state at R.T.P. changes down the group with both Fluorine and Chlorine being gases, Bromine is a liquid and Iodine is a solid. \item All of the Halogens are highly reactive, particularly with Groups I and II metals. They all combine with water to form strong acids. \begin{itemize} \item HF, Hydrofluoric acid \item HCl, Hydrochloric acid \item HBr, Hydrobromic acid \item HI, Hydroiodic acid \end{itemize} \item Fluorine, a pale yellow gas, is used in the form of fluorides in drinking water and toothpaste because it reduces tooth decay by hardening the enamel on teeth. \item Chlorine, a pale green gas, is used to make PVC plastic as well as household bleaches. It is also used to kill bacteria and viruses in drinking water. \item Bromine, a brown liquid, is used to make disinfectants, medicines and fire retardants. \item Iodine, a purple/black solid, is used in medicines and disinfectants and also as a photographic chemical. \item Astatine is a radioactive black solid. \end{itemize} \subsection{Group 0 -- The Noble Gases} These are Helium, Neon, Argon, Krypton, Xenon and Radon. They are all gases, exist as single atoms, colourless, and strongly unreactive. The only element in the group that forms compounds is Xenon, and this can combine with fluorine to form XeF$_{6}$. \begin{description} \item[Helium] an inert lifting gas \item[Argon] used in lights \item[Neon] used in LASERs \item[Xenon] % ... ??? \end{description} \subsection{The Transition Elements} \begin{itemize} \item The transition elements are a group of metals that occur in the central panel of the periodic table and include elements: \begin{itemize} \item 21$\rightarrow$30 \item 39$\rightarrow$48 \item 57$\rightarrow$80 \item 89$\rightarrow$112 \end{itemize} \item They have a number of features in common: \begin{itemize} \item High densities \item High melting and boiling points \begin{itemize} \item except Mercury (Hg) \end{itemize} \item They form strongly coloured compounds \end{itemize} \item They often act as catalysts in reactions e.g. \begin{itemize} \item Pt \item Fe \item Co \end{itemize} \item \end{itemize} \subsection{Bits and Pieces} \begin{itemize} \item Metal oxides will form basic (alkali) solutions when mixed with water. \item Non-metal oxides will form acidic solutions when mixed with water. \item An oxide that forms a neutral solution when mixed with water is called amphoteric. \item The position of an element in the group is determined by the number of electrons in its outer shell. \end{itemize} \section{The Mole Concept} \subsection{Size of Atoms} Atoms have only recently been actually observed. They are often smaller than 10$^{-30}$m. \subsection{Avagadro's Constant} A convenient way of calculating the number of moles is to use Avagadro's constant ($N_{A}$): \[N_{A}\approx6.023\times10^{23}mol^{-1}\] Or more simply:\[ 6\times10^{23}\] \subsection{Calculating Formulae} \paragraph{Example 1} If we were to combust some Magnesium ribbon in air, with careful calculations and measurements we can calculate the formula of the compound formed. If we started with 0.24g of Mg and ended up with 0.40g of Magnesium Oxide, we can work out the formula of the product. \begin{table}[htb!] \centering \begin{tabular}{|l|ll|} \hline Element & Mg & O \\ Masses & 0.24 & 0.16 \\ No. of Moles & $\frac{0.24}{24g}=0.01$ & $\frac{0.16}{16g}=0.01$ \\ \hline \end{tabular} \caption{Results of combustion of Magnesium.} \end{table} $\therefore$ formula $=$ \ce{MgO}. \paragraph{Example 2} An unknown chemical contains 0.12g of Carbon and 0.02g of Hydrogen. \begin{table}[htb!] \centering \begin{tabular}{|l|ll|} \hline Element & C & H \\ Masses & 0.12 & 0.02 \\ No. of Moles & $\frac{0.12}{12g}=0.01$ & $\frac{0.02}{1g}=0.02$\\ \hline \end{tabular} \caption{Results of combustion of an unknown chemical.} \end{table} $\therefore$ formula $=$ \ce{CH2}. This is an empirical formula. A mass spectrometer is used and the relative molecular mass is found to be 56g. If the empirical formula is calculated as a mass then \ce{CH2} would have a mass of 14g. Calculation: $\frac{56g}{14g}=4$. $\therefore$ actual formula $=$ \ce{C4H8}. \section{Formulae and Equations} \subsection{Stoichiometry} \begin{itemize} \item Stoichiometry is the measurement of equations to make sure that they balance (law of conservation of mass). E.g.: \item \ce{CaCO3} will break down into \ce{CaO} and \ce{CO2}. \[\ce{CaCO3_{(s)}} \longrightarrow \ce{CaCO_{(s)}} + \ce{CO2_{(g)}}\] \item $100g\longrightarrow56g+44g$ \item $\ce{2H2} + \ce{O2} \rightarrow \ce{2H2O}$ \item $\frac{2.00{\mr{g}\ce{NaCl}}}{58.44{\mr{g}\ce{NaCl}\mr{mol}}^{-1}}=0.034\ mol$ \end{itemize} \subsection{Group I} \begin{itemize} \item They form highly stable compounds. \begin{itemize} \item They have a valency of +1 because they form ions by losing an electron. \item All compounds of Group I elements are soluble in water. (Most Group VII are.) \item Group I metals are all light grey metals. \item The formulae of Group I compounds depends on the valency of the anion (negatively charged ion) they are attached to. e.g. \begin{itemize} \item Anions with $-1$ valency \begin{itemize} \item \ce{NaCl} \item \ce{NaI} \item \ce{NaNO3} \end{itemize} \item Anions with $-2$ valency \begin{itemize} \item \ce{Na2CO3} \item \ce{NA2So4} \item \ce{Na2O} \end{itemize} \item Anions with $-3$ valency \begin{itemize} \item \ce{Na3PO4} \end{itemize} \item Anions with $-4$ valency \begin{itemize} \item \ce{Na4C} (Sodium Carbide) \end{itemize} \end{itemize} \end{itemize} \end{itemize} \subsection{Group VII} \begin{itemize} \item Formulae of Group VII element compounds are determined by the valency of the cation (positively charged ion) they are attached to. e.g. \begin{itemize} \item Cations with a +1 valency \begin{itemize} \item NaCl \begin{itemize} \item NH$_{4}$I \end{itemize} \item Cations with a +2 valency \begin{itemize} \item CaCl$_{2}$ \item PbF$_{2}$ \end{itemize} \item Cations with a +3 valency \begin{itemize} \item AlI$_{3}$ \item GaCl$_{3}$ \end{itemize} \end{itemize} \end{itemize} \end{itemize} \subsection{Hydrocarbons} \beqn \ce{2C2H6} + \ce{7O2} & \longrightarrow & \ce{4CO2} + \ce{6H2O} \\ \ce{2C8H18} + \ce{25O2} & \longrightarrow & \ce{16CO2} + \ce{16H2O} \\ \ce{C6H12O6} + \ce{6O2} & \longrightarrow & \ce{6CO2} + \ce{6H2O} \\ \ce{C12H22O11} + \ce{12O2} & \longrightarrow & \ce{12CO2} + \ce{11H2O} \\ \ce{CH4O6} + \ce{2O2} & \longrightarrow & \ce{CO2} + \ce{2H2O} \\ \ce{2C2H6O2} & \longrightarrow & \ce{4CO2} + \ce{6H20} \\ \ce{C3H8O6} + \ce{5O2} & \longrightarrow & \ce{3CO2} + \ce{4H2O} \\ \ce{2C4H10} + \ce{36O2} & \longrightarrow & \ce{8CO2} + \ce{10H2O} \\ \ce{C5H12} + \ce{8O2} & \longrightarrow & \ce{5CO2} + \ce{6H2O} \eeqn \section{Ionic Bonding} \section{Covalent Bonding} \section{Limestone} \subsection{Introduction} \begin{itemize} \item The main mineral in Limestone is $CaCO_{3}$ or Calcite. \item Lime-rich rocks might be called limestone, chalk or marble. \end{itemize} \subsection{Uses of Carbonate Rocks} \begin{itemize} \item Building stone \begin{itemize} \item as a structural unit i.e. columns or blocks in 'classical' buildings \item as a decorative stone i.e. 'dressing stone' \end{itemize} \end{itemize} \section{Metals} \subsection{Properties of metals} \begin{itemize} \item Physical properties \begin{itemize} \item Usually solid at RTP. \item When struck they will produce a ring tone. \item Ductile: they can be drawn out into a wire. \item Malleable: they take on the shape of an object that hits them. \item Good conductor of heat. \item Good conductor of electricity. \item Most metals have a shiny surface when clean (or cut) (metallic lustre). \item Most have a high density ($=\frac{mass}{volume}$). \item High melting/boiling point. \end{itemize} \item Chemical properties \begin{itemize} \item Metals + dilute acids: any metal above copper in the reactivity series will react with an acid to produce a soluble salt of hydrogen gas e.g. \[ \ce{2HCl_{(aq)}}+\ce{Mg_{(s)}} \longrightarrow \ce{MgCl2_{(aq)}} + \ce{H2_{(g)}} \] \item Metals react with oxygen to form metal oxides e.g. \[ \ce{2Fe_{(s)}}+\ce{O2_{(g)}} \longrightarrow \ce{2FeO_{(s)}} \] Note: With the exception of Group I metal oxides, all other oxides are insoluble. \item Metals high in the reactivity series react readily with water (down to Mg) to produce hydrogen gas + a metal hydroxide. \end{itemize} \end{itemize} \subsection{Metallic Bonding} \begin{itemize} \item Metals have their own type of bonding that is quite different from ionic or covalent bonding. \item Metallic bonds are fairly strong bonds, shown by the fact that with one exception, all metals are solids at RTP. \item The arrangement of the metal is as a lattice with the individual atoms surrounded by large numbers of delocated electrons. \item These delocalised electrons mean that electric current can be conducted through the structure. \end{itemize} \begin{figure}[htb!] \centering \includegraphics[scale=1.2]{dc-current} \caption{A representation of DC (direct current).} \end{figure} \begin{figure}[htb!] \centering \includegraphics[scale=1.2]{ac-current} \caption{A representation of AC {alternating current}.} \end{figure} \begin{sidewaystable} \begin{tabular}{|l|p{4cm}p{4cm}p{4cm}p{4cm}|} \hline Reactivity Series & Reaction with Dilute Acid & Reaction with Air/Oxygen & Reaction with Water & Ease of Extraction \\ \hline Potassium (K) & Produce \ce{H2} with decreasing vigour & Burn very brightly and vigourously & Produce \ce{H2} with decreasing vigour in cold water & Difficult to extract \\ Sodium (Na) & $\downarrow$ & $\updownarrow$ & $\downarrow$ & $\updownarrow$ \\ Calcium (Ca) & $\downarrow$ & Burn to form an oxide with decreasing vigour & React with steam with decreasing vigour & Easier to extract \\ Magnesium (Mg) & $\downarrow$ & $\downarrow$ & $\downarrow$ & $\downarrow$ \\ Aluminium (Al) & $\downarrow$ & $\downarrow$ & $\downarrow$ & $\downarrow$ \\ Carbon (C) & $\downarrow$ & $\downarrow$ & $\downarrow$ & $\downarrow$ \\ Zinc (Zn) & $\downarrow$ & $\downarrow$ & $\downarrow$ & $\downarrow$ \\ Iron (Fe) & $\downarrow$ & $\downarrow$ & $\downarrow$ & $\downarrow$ \\ Lead (Pb) & $\downarrow$ & React slowly to form the oxide & Do not react with water or steam & $\downarrow$ \\ Copper (Cu) & Do not react with dilute acids & $\downarrow$ & $\updownarrow$ & $\downarrow$ \\ Silver (Ag) & $\updownarrow$ & Do not react & $\updownarrow$ & Found as element in nature \\ Gold (Au) & $\updownarrow$ & $\updownarrow$ & $\updownarrow$ & $\updownarrow$ \\ Platinum (Pt) & $\updownarrow$ & $\updownarrow$ & $\updownarrow$ & $\updownarrow$ \\ \hline \end{tabular} \caption{The Reactivity Series} \end{sidewaystable} \subsection{Alloys} The reason for using alloys is that we can tailor the properties of the metal to specific purposes. It is a crystal lattice like an elemental metal, but the lattice is irregular because the atoms are of different sizes. \begin{figure} \centering \includegraphics[scale=2.2]{alloy-crystal-lattice} \caption{The crystal lattice of an alloy.} \end{figure} \subsection{Competition Reactions in Aqueous Solutions} \begin{itemize} \item If a metal from high in the reactivity series is placed in an aqueous solution of a metal from lower in the series then the first metal will displace the second. e.g.: \begin{itemize} \item \[\ce{Mg_{(s)}} + \ce{CuSO4_{(aq)}} \longrightarrow \ce{Cu_{(s)}} + \ce{MgSO4_{(aq)}}\] \item \[\ce{MG_{(s)}} + \ce{Cu{(aq)}} \longrightarrow \ce{Mg_{(aq)}} + \ce{Cu_{(s)}} \] \item \[\ce{Zn_{(s)}} + \ce{Pb(NO3)2_{(aq)}} \longrightarrow \ce{Zn(NO3)2_{(aq)}} + \ce{Pb_{(s)}} \] \end{itemize} \item Most metal hydroxides are insoluble (the exceptions being the Group I metals) so an easy way to test for the presence of metal ions in solution is to add a few drops of $NaOH$. A coloured precipitate often forms. \end{itemize} \subsection{Aluminium -- The Problem Metal} \section{Organic Fuels} \subsection{Sources of fuels} \begin{itemize} \item The main source of combustible fuels is either petroleum or gas. \item The main sources of petroleum are from OPEC countries, the North Sea, Russia and Bass Strait. \item The method of recovering usable portions involves a process known as catalytic cracking where the petroleum is heated in a furnace then pumped into a tower structure and allowed to cool. \item The various fractions obtained have different numbers of carbon atoms in a chain. \begin{itemize} \item A chain of $C_{1{\rightarrow}5}$ is a gas. \item $C_{6{\rightarrow}20}$ are liquids. \item $C_{>20}$ is a solid. (All at R.T.P.) \end{itemize} \item The carbon atoms are mostly joined to hydrogen atoms so they are called hydrocarbons. \item Fuels are generally burned. As such the chemical reaction is a combustion reaction. \item The products of a combustion reaction, where complete reaction occurs, are $\ce{CO2} + \ce{H2O}$. \end{itemize} \section{Acids, Bases \& Salt Preparations} \subsection{Oxygen \& Oxides} Oxygen is a colourless odourless gas which is essential to life as we know it on this planet. Oxygen is the second most abundant gas in the air (21\%). It is also found in many different compounds in the earth's crust, making it the most abundant elemsnt there also at 34\% (silicon is next at 26\%). When an element or a compound reacts with oxygen it is said to have undergone an oxidisation (or combustion) reaction. If the reaction involves a flame then it is called burning. The products of these reactions of elements or compounds with oxygen are \textsc{Oxides}. Both elements and compounds burn in pure oxygen more fiercely than in air, because pure oxygen contains no nitrogen which is a suppressor of reactions. Examples of oxidations: \[ \ce{4K + O2 {\longrightarrow} 2K2O} \] \[ \ce{S + O2 {\longrightarrow} SO2} \] \[ \ce{4Fe + 3O2 {\longrightarrow} 2Fe2O3} \] \[ \ce{4NH3 + 7O2 {\longrightarrow} 4NO2 + 6H2O} \] \[ \ce{2H2S + 3O2 {\longrightarrow} 2SO2 + 2H2O} \] \subsection{Properties of the Oxides of the Elements} Oxides of the elements can be made by burning the elements in oxygen. Oxides which react with acids are called \textsc{Basic} oxides. Some basic oxides are soluble in water. When they dissolve they form a solution whose pH is greater than 7 (full range indicator turns from green to blue). These soluble basic oxides are called \textsc{Alkaline} oxides. Oxides which dissolve in water to give a solution whose pH is less than 7 (full range indicator turns from green to orange or red) are \textsc{Acidic} oxides. Some oxides are insoluble in water. If they are soluble in acids the yare basic oxides. If they are soluble in alkalis they are acidic oxides. A few oxides are insoluble in water but soouble in both acids and alkalis. They show both acidic and basic properties and are called \textsc{Amphoteric} oxides. Metallic elements on the left-hand-side of the Periodic Table form \textsc{Basic} oxides; non-metallic elemnts on the right-hand-side of the Periodic Table form \textsc{Acidic} oxides. \subsection{Acids, Bases \& Salts} Acids and alkalis (soluble bases) are a constant part of everyday life. The most well-known acid-containg substance is probably vinegar, but there are many others that we come across every day. \begin{table} \centering \scalebox{0.9}{\begin{tabular}{ll|ll}\hline \multicolumn{2}{|c||}{\textbf{Acids}} & \multicolumn{2}{c|}{\textbf{Alkalis}} \\ \hline Vinegar & Acetic & Garden lime & Calcium (Hydr)Oxide \\ Fizzy drinks & Ascorbic & Oven cleaner & Sodium/Potassium Hydroxide \\ Rust remover & Phosphoric & Soap & Potassium Hydroxide \\ Fruit juice & Citric & Household cleaner & Sodium Hydroxide/Ammonia \\ Kettle descaler & Hydrochloric & Baking soda & Sodium Hydrogen Carbonate \\ (Car) Battery acid & Sulphuric & Washing soda & Sodium Carbonate \\ Fertilisers \& Explosives & Nitric & Milk of Magnesia & Magnesium/Sodium Carbonate \\ TCP & Phosphoric & & \\ Fabric softeners & Sorbic & & \\ \hline \end{tabular}} \caption{Some common uses of Acids and Alkalis.} \end{table} The word `acid' means sharp or sour --- acids are known for their characteristic sour taste. Lemon juice, uncorked wine, and milk that has gone off all contain acids. Acids and alkalis were found to change the colour of some plant pigments. The colour extracted from red cabbage and from lichens have different colours in acidic and alkaline solution. The indicator derived from lichen is called \textsc{Litmus}, and goes red in acid and blue in alkalis. \textsc{Full-range Indicator} (similar to \textsc{Universal Indicator}) is a mixture of indicators, which can show how strongly acidic or alkaline a substance is. These substances which change colour when placed in acids and alkalis are called \textsc{Indicators}. The effects of acids can be removed by alkalis --- this kind of reaction is called a \textsc{Neutralisation} reaction. The properties of acidic and alkaline solutions are best explained in terms of the ions which they contain. An \textsc{Acid} is a compound which in aqueous solution produces hydrogen ions \ce{H+}. An \textsc{Alkali} is a compound which in aqueous solution produces hydroxyl ions \ce{OH-}. Acids, when pure, are composed of molecules. It is only in the presence of water that they release hydrogen ions. (In fact, each \ce{H+} ion which is set free bonds to a water molecule forming the ion \ce{H3O+}.) \textsc{Strong Acids} are almost totally broken down into ions in aqueous solution (fully dissociated) e.g.\ \ce{Hcl}, \ce{HNO3}, \ce{H2SO4}. \textsc{Weak Acids} are only partly broken down into ions when in aqueous solution e.g.\ \ce{CH3CO2H} (ethanoic acid). \[ \ce{CH3CO2H + H2O {\rightleftharpoons} CH3CO2- + H3O+} \] \begin{table} \begin{tabular}{|l|c|c|c|} \hline Name & Formula & Ionic Formula & Strong or Weak? \\ \hline Hydrochloric acid & \ce{HCl} & \ce{H+ / Cl-} & S \\ Sulphuric acid & \ce{H2SO4} & \ce{2H+ / SO_4^2-} & S \\ Nitric acid & \ce{HNO3} & \ce{H+ / NO3-} & S \\ Acetic or Ethanoic acid & \ce{CH3CO2H} & \ce{H+ / CH3CO2-} & W \\ Citric acid & \ce{C6H8O7} & \ce{3H+ / C6H5O7^{3-}} & W \\ \hline \end{tabular} \caption{Important Acids} \end{table} \begin{table} \begin{tabular}{|l|c|c|c|} \hline Name & Formula & Ionic Formula & Strong or Weak? \\ \hline Sodium hydroxide & \ce{NaOH} & \ce{Na+ / OH-} & S \\ Calcium hydroxide & \ce{Ca(OH)2} & \ce{Ca^2+ / 2OH-} & S \\ Magnesium hydroxide & \ce{Mg(OH)2} & \ce{Mg^2+ / 2OH-} & S \\ Sodium Hydrogen carbonate & \ce{NaHCO3} & \ce{Na+ / HCO3-} & W \\ Ammonia (hydroxide) & \ce{NH4OH} & \ce{NH4+ / OH-} & W \\ \hline \end{tabular} \caption{Important Bases} \end{table} \subsubsection{The Role of Water} Acids are covalently-bonded molecules. E.g.\ Hydrogen Chloride (\ce{HCl}). \begin{table}[htb!] \centering \begin{tabular}{r|ll} \hline & \ce{HCl} in \ce{H2O} & \ce{HCl} in organic solvent \\ \hline Temperature Change & Marked rise & Little change \\ Reaction with \ce{Mg} & Fizzes and \ce{Mg} disappears & No reaction \\ Reaction with marble & Fizzes and marble disappears & No reaction \\ Electrical conductivity & Yes & No \\ Universal indicator paper & Dark red & No change of colour \\ \hline \end{tabular} \caption{Reactions of \ce{HCl}.} \end{table} \[ \ce{HCl_{(g)} + H2O_{(l)} {\longrightarrow} H3O+_{(aq)} + Cl-_{(aq)}} \] The molecules dissociate into ions. Atoms form ions in water. One of these ions is \ce{H+_{(aq)}}. The pH is a measure of the strength of an acid or alkali, i.e.\ how much it dissociates into ions when put in water. \subsection{Salt Preparations} To obtain pure samples of soluble salts it is necessary to ensure that all of the acid has been used up in the reaction and that no excess of the other reagent is left to contaminate the product. \begin{enumerate} \item An excess of solid is added to the acid to make sure all the acid is used up. \item The excess solid is filtered off. \item Salt srystals are then produced by evaporation and crystallisation: \begin{enumerate} \item The dilute solution of the salt is evaporated first directly then slowly over a steam bath. \item The solution is tested for saturaion by removing a droplet of the solution on a glass rod. \item If crystals form quickly when the droplet is cooled, then the solution is left to crystallise. \item If crystals do not form quickly evaporation is continued until the solution is saturated \end{enumerate} \end{enumerate} \subsection{Ionic Equations} Many chemical reactions involve ionic substances. Ionic formulae can be shown in the equation for the reaction. In these equations the numbers of atoms and the charges on both sides must balance. \subsubsection{Acids Reacting With Bases} Many insoluble metal oxides react with dilute acids to form soluble salts. \[ \ce{General\:equation{:\;\;} Base + Acid {\longrightarrow} Salt + Water} \] \[ \ce{Word\: equation{:\;\;} Copper(II)\: oxide{(s)} + Hydrochloric\: acid{(aq)} {\longrightarrow} Copper(II)\: chloride{(aq)} + water{(l)} } \] \[ \ce{Ordinary\: equation{:\;\;} CuO + 2HCl {\longrightarrow} CuCl2 + H2O } \] \[ \ce{Ionic\: equation{:\;\;} CuO + 2H+ + 2Cl- {\longrightarrow} Cu^{2+} + 2Cl- + H2O } \] Here O and H ions are spectator ions, so the ionic equation could be written as: \[ \ce{O^{2-} + 2H+ {\longrightarrow} H2O } \] However, the Cu ions change from being solid to aqueous, so it is usual to include them in the final ionic equation. It is also usual to write the base as a molecular formula. \[ \ce{Ionic\: equation{:\;\;} CuO + 2H+ {\longrightarrow} Cu^{2+} + H2O } \] \subsubsection{Acids Reacting With Carbonates} Carbon dioxide is given off when any dilute acid reacts with any carbonate at room temperature. \beqn \mr{General \: equation:} \; \; \mr{Carbonate} + \mr{Acid} & \longrightarrow & \mr{Salt} + \mr{Carbon \: dioxide} + \mr{Water} \\ \mr{Word \: equation:} \; \; \mr{Calcium \: carbonate(s)} + \mr{Hydrochloric \: acid(aq)} & \longrightarrow & \mr{Calcium \: chloride(aq)} + \\ & & \mr{Carbon \: dioxide(g)} + \mr{water(l)} \\ \mr{Ordinary \: equation:} \; \; \ce{CaCO3} + \ce{2HCl} & \longrightarrow & \ce{CaCl2} + \ce{CO2} + \ce{H2O} \\ \mr{Ionic \: equation:} \; \; \mr{CaCO3} + \mr{2H+} + \mr{2Cl-} & \longrightarrow & \mr{2Cl}- + \mr{CO2} + \mr{H2O} \\ \eeqn Here \ce{Cl} ions are spectator ions. Since the \ce{Ca} ions in Calcium carbonate change from being in the solid state to being in aqueous solution the formula of calcium carbonate is shown as a molecular formula. Ionic equation omitting spectator ions: \[ \ce{ CaCO3 + 2H+ {\longrightarrow} CO2 + H2O + Ca_{(aq)}^{2+} } \] \subsubsection{Acids Reacting With Metals} Hydrogen is given off when most dilute acids (Nitric acid is an exception) react with metals. Heat is needed for the less reactive metals and lead and copper are too unreactive to give this reaction at all. \beqn \mr{General \:equation:} \; \; \mr{Metal} + \mr{Acid} & \longrightarrow & \mr{Salt} + \mr{Hydrogen} \\ \mr{Word \: equation:} \; \; \mr{Zinc(s)} + \mr{Sulphuric \: acid(aq)} & \longrightarrow & \mr{Zinc\: sulphate(aq)} + \mr{Hydrogen(g)} \\ \mr{Ordinary\: equation:} \; \; \ce{Zn} + \ce{H2SO4} &\longrightarrow& \ce{ZnSO4} + \ce{H2} \\ \mr{Ionic\: equation:} \; \; \ce{Zn} + \ce{2H+} + \ce{SO4^{2-}} &\longrightarrow& \ce{Zn^{2+}} + \ce{SO4^{2-}} + \ce{H2} \\ \eeqn Here sulphate ions are spectator ions, so they can be omitted from the final ionic equation. Ionic equation omitting spectator ions: \[ \ce{ Zn + 2H- {\longrightarrow} Zn_{(aq)}^{2+} + H2{(g)} } \] \section{Ammonia, Fertilisers, Equilibrium \& Sulphur} \subsection{Nitrogen} Nitrogen is a colourless, odourless gas which makes up about 78\% of the composition of the air. It is a diatomic gas covalently bonded by a triple covalent bond. It is an extremely unreactive and many of its uses (e.g.\ Liquid nitrogen for fast freezing food, Nitrates for fertiliser) depend on its inertness. However, Nitrogen is essential for healthy life as it is an important component of protein molecules. Nitrogen is found in a variety of inorganic compounds either as Ammonium compounds \ce{NH4+} or as Ammonia \ce{NH3}. \subsection{Ammonia} Ammonia is a weak base which produces an alkaline solution. It is sometimes used in cleaning agents as a degreasing agent. It convert grease into soluble soaps. This reaction is called Saponification. Ammonia gas dissolves in water to form an aqueous solution called Ammonium hydroxide: \[ \ce{NH3} + \ce{H2O} \rightleftharpoons \ce{NH4+} + \ce{OH-} \] Ammonia is in great demand throughout the world. It has a variety of uses, as shown by the pie chart below. \bigskip \begin{center} \begin{tikzpicture} \draw (0,0) circle (4cm); \draw (0,0) -- (288:4cm); \draw (0,0) -- (324:4cm); \draw (0,0) -- (342:4cm); \draw (0,0) -- (360:4cm); \draw (0,0) (144:2cm) node {\parbox{4cm}{\center Fertilisers \\ 80\%}}; \draw (0,0) (306:5cm) node {\parbox{3cm}{\center Wood pulp production 10\%}}; \draw (0,0) (333:5cm) node {\parbox{3cm}{\center Nitric acid 5\%}}; \draw (0,0) (351:5cm) node {\parbox{3cm}{\center Nylon 5\%}}; \end{tikzpicture} \end{center} \bigskip The manufacture of ammonia from nitrogen and hydrogen in the laboratory is almost impossible. \[ \ce{N2} + \ce{3H2} \rightleftharpoons \ce{2NH3} \] In industry, as the conditions can be carefully controlled, it becomes possible (this process is called the Haber Process after the German chemist Fritz Haber). \subsection{Chemical Equilibria} How far do chemical reactions go? When dry blue powdered copper (II) sulphate crystals are heated a vapour is given off and a white residue remains. If the vapour is condensed and collected it can be shown to be water. The white residue is anhydrous Copper(II) sulphate. \[ \ce{ CuSO4 + 5H2O {\xrightarrow{forwards\: reaction}} CuSO4 + 5H2O \;\;\;\; endothermic\:({{\Delta}\mr{H}+}) } \] When a few drops of water are added to the cold anhydrous Copper(II) sulphate the crystals turn blue and heat energy is recovered. \[ \ce{ CuSO4 + 5H2O {\xrightarrow{backwards\: reaction}} CuSO4 + 5H2O \;\;\;\; exothermic\:({{\Delta}\mr{H}-}) } \] This type of reaction is an example of a reversible reaction. A reversible change is one in which the products formed are capable of reacting with each other to reform the original substance or substances. Reactions that can proceed in both directions are given a $\rightleftharpoons$ sign between the reactants and products. The special kind of balance which keeps still by moving is called an \textbf{Equilibrium}. Chemical equilibria are very common in industrial processes where gases are involved because the reactions take place within a closed system. Factors which affect the rate of a chemical reaction might effect the position of an equilibrium reaction as there are two opposing rates to consider. Reaction conditions must be carefully chosen so that the process is economically viable. Some chemical reactions are reversible, in that not only can reactants form products, but also products can form reactants, e.g.\ heated iron reacts with steam to form iron oxide and hydrogen, while heated iron oxide reacts with hydrogen to form iron and steam. For a particular reaction, it is possible for reactants and products to be present in a \textbf{dynamic equilibrium} with one another. This means that the amounts of reactants and products remain constant, due to the rate at which reactants form products exactly equalling the rate at which products form reactants. The position of an equilibrium (i.e.\ the relative amounts of reactants and products present at equilibrium --- effectively the yield of products) can be affected by changes in the reaction conditions. Changes can be predicted using \textbf{Le Chatelier's Principle}. This states that: \begin{quote} \textit{Whenever a system in dynamic equilibrium is disturbed, it tends to respond in such a way as to oppose the disturbance and so restore equilibrium.} \end{quote} \subsubsection{Le Chatelier's Principle} \paragraph{Concentration} if the concentration of a reactant or product is altered when a chemical system is in equilibrium the equilibrium will try to restore it. So if a product is removed as the equilibrium proceeds then the equilibrium will try to increase the products' concentration and so increase the yield. \paragraph{Temperature} Most reactions give out heat when the reactants react to form the products, i.e.\ the formation of products is exothermic. This means that the formation of reactants from products is endothermic --- it absorbs heat. If a reaction of this sort is in equilibrium, and an increase in termperature is applied, the position of equilibrium will move in an attempt to lower the temperature. This means that more reactants will be formed, as this will absorb heat. Therefore the yield of products is reduced by an increase in temperature. \paragraph{Pressure} (\emph{\small For reactions involving gases}) If, for example, there are less gas molecules on the products side of the equation than on the reactants side, then a movement to form more products will produce a decrease in pressure. This, if the pressure is increased on a reaction of this sort in equilibium, then the posiion of equilibrium will move in an attempt to lower the pressure. Therefore more products will be formed and the yield of products is increased by an increase in pressure. % Two important industrial processes which involve reversible % reactions which can reach dynamic equilibrium are now % described. ---- Where should this go? \subsection{Manufacture of Ammonia --- The Haber Process} The Haber Process is the syntheses of Ammonia (\ce{NH3}) from Nitrogen and Hydrogen. \[ \ce{Nitrogen + Hydrogen {\rightleftharpoons} Ammonia} \] \[ \ce{N2{(g)} + 3H2{(g)} {\rightleftharpoons} 2NH3{(g)}} \] The production of Ammonia is exothermic and there are less gas molecules on the products side of the equation than on the reactants side. Therefore, a high yield of Ammonia is favoured by a low temperature and a high pressure. Conditions used to obtain a satisfactory yield at a satisfactory rate: \begin{description} \item[Moderate temperature (450$\ensuremath{^\circ}$C)] Higher temperature favours the decomposition of Ammonia, thereby reducing the yield. Lower temperature reduces the rate of reaction. \item[Moderate pressure (200 atmospheres)] Higher pressure increases cost of industrial plant, lower pressure decreases both yield and rate. \item[Catalyst (iron)] Helps to produce a satisfactory rate, but does not change the equilibrium yield, as it speeds up the production of ammonia and the decomposition of Ammonia equally. \end{description} \input{haber.pic} \subsection{Manufacture of Nitric Acid} One of the main uses of Ammonia is to convert it into Nitric acid. The Nitrogen dioxide is formed in a two step process. First the Ammonia is oxidised by air in the presence of a Platinum catalyst to produce Nitrogen monoxide. \[ \ce{4NH3 + 5O2 {\rightleftharpoons} 4NO + 6H2O} \] This reaction is exothermic. As the mixture cools the Nitrogen monoxide reacts with more air to give Nitrogen dioxide. \[ \ce{6NO + 3O2 + 2H2O {\rightleftharpoons} 4HNO3 + 2NO} \] \[ \ce{2NO + O2 {\rightleftharpoons} 2NO2} \] \[ \ce{2MO2 + H2O {\rightleftharpoons} 2HNO3} \] On reaction with water Nitric acid is formed. Apart from fertilisers Nitric acis is also used to make explosives. An Italian chemist called Professr Sobrero converted Nitric acid into Nitroglycerine, a very unstable oily liquid. It was Alfred Nobel who lost his borther in a Nitroglycerine explosion who was responsible for stabilising it by mixing it with clay to form Dynamite. \subsection{Manufacture of Sulphuric Acid --- The Contact Process} Sulphuric acid is manufactured from sulphur in the following stages: \begin{enumerate} \item Sulphur is burned in air to make Sulphur dioxide (\ce{SO2}) \item \beqn \ce{Sulphur\: dioxide + Oxygen} & \rightleftharpoons & \ce{Sulphur\: trioxide} \\ \ce{2SO2{(g)} + O2{(g)}} & \rightleftharpoons & \ce{2SO3{(g)}} \eeqn The production of Sulphur trioxide in this stage is exothermic and there are less gas molecules on the products side of the equation than on the reactants side. Therefore a high yield of Sulphur trioxide is favoured by a low temperature and a high pressure. In practice, to obtain a satisfactory yield at a satisfactory rate, the following conditions are used: \begin{description} \item[Temperature of 450$^\circ$C] This is a compromise termperature, as a lower temperature would decrease the rate of reaction and a higher temperature would decrease the yield. \item[Pressure only just above atmospheric] Rate and yield are satisfactory wwithout the need for high pressure. \item[Catalyst of Vanadium pentoxide (\ce{V2O5})] This increases the rate of reaction (but has no effect on yield) \end{description} \item The Sulphur trioxide is absorbed in concentrated Sulphuric acid and the resulting liquid is added to water. \end{enumerate} The Sulphur dioxide needed is obtained from: \begin{itemize} \item The combustion of Sulphur \item The removal of unpleasant-smelling Sulphur compounds from petroleum oil \item The removal of Hydrogen sulphide from natural gas (not North Sea gas) \item Roasting metal sulphides in order to extract metals from their ores \end{itemize} \input{contact.pic} Sulphur dioxide and air are purified by electrostatic precipitation. A Sulphuric acid plant must be sited in a place where supplies of Sulphur dioxide can reach it, that is, near a port where caroes of Sulphur can arrive by ship, or near a gas refinery or a metal smelter. \bigskip \begin{center} \begin{tikzpicture} \coordinate (O) at (0,0); \draw (O) circle (6cm); \coordinate (A) at (360 * 2 / 7:6cm); \coordinate (B) at (360 * 2 / 7 + 360 * 1 / 7:6cm); \coordinate (C) at (360 * 2 / 7 + 360 * 1 / 7 + 360 * 1 / 7:6cm); \coordinate (D) at (360 * 2 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7:6cm); \coordinate (E) at (360 * 2 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7:6cm); \coordinate (F) at (360 * 2 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 14:6cm); \coordinate (G) at (360 * 2 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 14 + 360 * 1 / 42:6cm); \coordinate (H) at (360 * 2 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 14 + 360 * 1 / 42 + 360 * 1 / 42:6cm); \coordinate (I) at (360 * 2 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 7 + 360 * 1 / 14 + 360 * 1 / 42 + 360 * 1 / 42 + 360 * 1 / 42:6cm); \draw (O)--(A) node [shift={(4cm,-3cm)}] {\parbox{3cm}{\center Fertilisers: \\ phosphates and ammonium sulphate}} ; \draw (O)--(B) node [shift={(2.5cm,1cm)}] {\parbox{3cm}{\center Soaps and detergents}} ; \draw (O)--(C) node [shift={(1.75cm,2.75cm)}] {\parbox{3cm}{\center Soaps Paints and pigments, e.g.\ \ce{TiO2}}} ; \draw (O)--(D) node [shift={(-1.25cm,3.25cm)}] {\parbox{4cm}{\center Batteries \\ Pharmaceuticals \\ Insecticides}} ; \draw (O)--(E) node [shift={(-2.75cm,0.75cm)}] {\parbox{4cm}{\center Plastics and many chemicals, including metal sulphates}} ; \draw (O)--(F) node [shift={(-2cm,0cm)}] {\parbox{2cm}{\center Fibres, e.g.\ nylon}} ; \draw (O)--(G); \coordinate (Gpoint) at ($(G) - (0.5,0.25)$); \node at ($(Gpoint) + (1.5,-0.5)$) (Glabel) {Dyes}; \draw[->] (Glabel) -- (Gpoint); \draw (O)--(H); \coordinate (Hpoint) at ($(H) - (0.25,0.25)$); \node at ($(Hpoint) + (2.5,-0.5)$) (Hlabel) {Oil and petrol}; \draw[->] (Hlabel) -- (Hpoint); \draw (O)--(I); \coordinate (Ipoint) at ($(I) - (0.25,0.5)$); \node at ($(Ipoint) + (2.5,0.5)$) (Ilabel) {\parbox{2cm}{\center Pickling iron and steel}}; \draw[->] (Ilabel) -- (Ipoint); \end{tikzpicture} \end{center} \bigskip \subsection{How The Chemical Industry Is Organised} All manufactueres -- whether of chemical products of anything else -- are engaged in a process which may be summed up as: \[ \mr{raw \: materials} \rightarrow \mr{processing} \rightarrow \mr{products} \] which means taht they are concerned with: \[ \mr{supplies} \rightarrow \mr{manufacture} \rightarrow \mr{sales} \] The chemical industry is a business which has grown up to provide a livelihood for those who work in it, and a profit for shareholders who have invested money in chemical companies. At the same time it is there to meet the need for an ever-growing range of materials and products. \subsubsection{Processing \& Manufacture} The major divisions of chemical industry are: \begin{tabular}{ll} \hline agricultural chemicals & heavy chemicals \\ dyestuffs & metals \\ explosives & paints \\ fibres & plastics \& petrochemicals \\ fien chemicals & pharmaceuticals \\ \hline \end{tabular} \subsubsection{Locating a Chemical Plant} The decision on where to build a particular chemcial plant is an extremely complicated one. These are some of the more important things whcih have to be thought about: \begin{description} \item[Communications] road, rail, sea and air access are neede for bringing in raw materials and distributing products. \item[Environment] this must not me affected too much while building the plant, or disposing of its waste. \item[Power supply] this must be available, or it must be possible to generate power at the site. \item[Utilities] examples are water and fuel; these must be available at reasonable cost. \item[Workers] these must be recruited, and perhaps moved to the area and housed. \end{description} \subsubsection{The Supply of Raw Materials} The main raw metrials for the chemical industry are: oil and natural gas, coal, metal ores, minerals like salt and limstone, and the gases of the air. All of these must be round, extracted and transported to where they are needed. There are important issues here: \begin{itemize} \item Some raw materials are running out. Since new mineral deposits are beign found all the time, and enw methods for exploiting them are beign developed, it is not possible to say precisely how long any particular materials will last. Our own supplies of fossil fuels are a good example. Coal will last longer than the others, if wecontinue extraction at the present rate. The chemical industry is already adapting to likely changes in its supply of raw materials. \item Our supply of some materials, notably tin, mercury, uranium and gold, depends on the fact that miners in some parts of the world work in dangerous conditions for low rates of pay. So long as the world wants these materials, someone has to get them out. Social and political changes in different parts of the world may have a considerable effect on the prices and supply of some metal ores. \end{itemize} \subsection{The Manufacture \& Use of Fertilisers} Ammonia and sulphuric acid are both used in the manufacture of fertiliesrs. Aqueous ammonia can be used directly as a nitrogenous fertiliser. Ammonia can also be used to manufacture nitric acid and ammonium salts, e.g.\ ammonium nitrate \ce{PNH4NO3}, is manufactured by the reaction of ammonia and nitric acid. Ammonium salts (e.g.\ ammonium nitrate), minerals and some phosphates are soluble in water, and if these are washed into lakes and rivers they cause excessive plant growth. The decay of this plant material deoxygenates the water. Nitrates are very difficult to remvoe from domestic water sources, and are potentially harmful to health. The main elements required for healthy plant growth are: \begin{description} \item[Nitrogen] taken into plant roots in the form of nitrate ions; a lack causes stunted growth, harsh and fibrous leaves \item[Phosphorus] a lack causes stunted growth, and stunted grey leaves \item[Potassium] a lack causes stunted growth, and premature death of leaves \end{description} NPK fertiliers must contain the appropriate compositions of tehse elements. In addition, physical state, solubility and pH are other factors which must be considered when choosing a fertiliser. \begin{table} \begin{tabular}{lll} \hline \it Raw material & Where it comes from & Used to make \\ \hline \rm nitrogen & the air & ammonia for all fertilisers \\ water & reservoirs & ammonia \\ natural gas & under the sea & ammonia \\ potassium chloride & mines & compound fertilisers \\ phosphate rock & mined & compound fertilisers \\ sulphur & mined & compound fertilisers \\ \end{tabular} \caption{Raw materials for the manufacture of fertilisers.} \end{table} \begin{figure} \begin{tikzpicture} \node (water) {water}; \end{tikzpicture} \end{figure} \subsubsection{Problems with Fertilisers} \section{Volumetric Analysis} Many chemical reactions take place in aqueous solution and so it is important to have a standard amount in solution. Concentration is measured in $\frac{\mr{mol}}{\mr{dm}^3}$. A 1 molar solution contains 1 mole of solute dissolved in exactly 1$dm^3$ of solvent, usually water. So a 1 molar solution of Sodium chloride contains 58.5g of Sodium chloride dissolved in 1000cm$^3$ of water. \subsection{The Volumes \& Concentrations of Solutions} The concentration can be stated in two ways: \begin{enumerate} \item In $\frac{\mr{mol}}{\mr{dm}^3}$, the amount, in moles, of solute contained in 1 dm$^3$ of solution. \item In $\frac{\mr{g}}{\mr{dm}^3}$, the amount, in g, of solute contained in 1 dm$^3$ of solution. \end{enumerate} There then exists two similar simple relationships between the concentration of a solution, amount of solute and the volume of the solution: \begin{enumerate} \item $\mr{Concentration}\: (\frac{\mr{mol}}{\mr{dm}^3}) = \frac{\mr{amount}}{\mr{volume}}$ \item $\mr{Concentration}\: (\frac{\mr{g}}{\mr{dm}^3}) = \frac{\mr{mass}}{\mr{volume}}$ \end{enumerate} These can be combined to form: \[ \mr{Concentration}\: (\frac{\mr{g}}{\mr{dm}^3}) = \mr{Concentration}\: (\frac{\mr{mol}}{\mr{dm}^3}) \times \mr{Molar\: mass} (\frac{\mr{g}}{\mr{mol}}) \] \subsection{Titrations} \subsubsection{Mathod} \subsubsection{Calculation} \subsection{Volumetric Analysis} \subsubsection{The Pipette} \subsubsection{The Burette} \subsubsection{Performing a Titration} \section{Organic I} Organic Chemistry is the study of the chemistry of the element Carbon. Carbon has the unique ability to bond to both itself and many other elements (mainly non-metallic). The main sources of Carbon are plants, animals, coal, crude oil and natural gas. Today fossil fuels provide 93\% of our fuels while nuclear and hydoelectric power only provide 7\%. \begin{tabular}{lclllll} & $\nearrow$ & Coal 34\% & $\rightarrow$ & (heat without \ce{O2}) & $\rightarrow$ & Coal gas \\ Fossil fuels are & $\rightarrow$ & Crude oil 34\% & & $\downarrow$ & $\searrow$ & \\ & $\searrow$ & Natural gas 24\% & & Coke & & Coal tar \end{tabular} \subsection{Fuels} Natural gas is 97\% Methane. Petrol is a mixture of Hydrocarbons. \subsubsection{Fractional Distillation} \begin{tabular}{crll} \multirow{6}{*}{$\downarrow$} & 1 & Petroleum gas & \\ & 2 & Petroleum fraction & -- petrol engines, cars (gasoline) \\ & 3 & Paraffin & -- oil stoves, aircraft fuel (kerosene) \\ & 4 & Diesel & -- diesel engines \\ & 5 & Lubrication fraction & -- lubricants, waxes \& polishes \\ & 6 & Bitumen fraction & -- roads \end{tabular} \medskip (\emph{higher boiling point/heavier/larger molecules towards bottom}) Carbon dioxide is produced by: \begin{itemize} \item complete combustion (e.g.\ hydrocarbons) \item respiraton \item reaction of acids \& alkalis \end{itemize} \subsection{Naming Organic Compounds} \begin{table}[!htbp] \centering \begin{tabular}{lr} \hline Name & No.\ of Carbon Atoms \\ \hline Methane & 1 \\ Ethane & 2 \\ Propane & 3 \\ Butane & 4 \\ Pentane & 5 \\ Hexane & 6 \\ \hline \end{tabular} \caption{Names \& numbers of carbon atoms in alkanes.} \end{table} \begin{table}[!htbp] \centering \begin{tabular}{ll} \hline Side Chains (\emph{allyl}) & Formula \\ \hline methyl & \ce{CH3}-- \\ ethyl & \ce{C2H5}-- \\ propyl & \ce{C3H7}-- \\ \hline \end{tabular} \caption{Formulae of allyls.} \end{table} \[ \begin{array}{ccccccccc} &&&& \ce{CH3} &&&& \\ 1 && 2 && \textSFxi 3 && 4 && 5 \\ \ce{CH3} & --- & \ce{CH} & --- & \ce{CH} & --- & \ce{CH2} & --- & \ce{CH3} \\ && \textSFxi &&&&&& \\ && \ce{CH3} && \multicolumn{4}{c}{2,3-diMethylpentane} \end{array} \] \subsection{Homologous Series} A homologous series is a group of similar molecules. For example: \begin{itemize} \item Alkanes \item Alkenes \item Carbolic acid (with \ce{CO2H} group) \item Aldehydes \item Alcohols (with \ce{OH} group) \end{itemize} Structural isomerism is where two molecules have the same molecular formula, but have different structual formulae. \subsection{Alkenes} General formula is \ce{C_nH_{2n}}. \begin{itemize} \item Contain a \ce{C\,=\,C} double covalent bond \item Are unsaturated hydrocarbons \item Need at least 2 Carbon atoms \end{itemize} \begin{itemize} \item Ethene -- \ce{C2H4} \\ \includegraphics{ethene} \item Propene -- \ce{C3H6} \\ \includegraphics{propene} \item Butene -- \ce{C3H8} \\ \begin{itemize} \item but-1-ene -- \ce{CH3CH2CH\,=\,CH2} \item but-2-ene (two isomers) -- \ce{CH3CH\,=\,CHCH3}\\ \end{itemize} \end{itemize} \includegraphics{but-2-ene-1} \hskip 2cm \emph{and} \hskip 2cm \includegraphics{but-2-ene-2} \begin{figure}[!h] \centering \begin{framed} \includegraphics{additionreaction} \end{framed} \caption{Alkenes undergo addition reactions (regardless of light conditions)} \end{figure} The test to determine whether is substance is an Alkane or Alkene (i.e.\ a test for unsaturation) is to add Bromine water. \subsection{Alkanes} Alkanes undergo a substitution reacction with Halogens in the presence of light. \subsection{Reactions of Alkanes \& Alkenes} \subsubsection{Addition Reactions} \subsection{Polymers} \subsubsection{Synthetic Organic Polymers} \begin{figure}[!h] \centering \begin{framed} \includegraphics{cyclohexanereaction} \end{framed} \caption{Addition reaction of Cyclohexane and Bromine} \end{figure} \subsubsection{Polymerisation} \section{Rates of Reaction} Different chemical reactions take place at different speeds. Some reactions seem instantaneous on mixing the reactants, e.g. \\ Silver nitrate solution \& \\ Sodium chloride. As soon as these two reactants are mixed a white precipitate is formed. Some reactions require energy ebfore they start, e.g.\ the burning of a piece of Magnesium. Some reactions take a long time, e.g.\ the rusting of Iron. \[ \ce{Reactants {\longrightarrow} Products} \] You can measure the reate of appearance of a product or the rate of disappearance of a reactant. So you need to measure a property of the product/reactant that changes with time: \begin{itemize} \item Mass \item Volume \item Colour \end{itemize} \subsubsection*{Methods of measuring rates of reactions} (\emph{for reactions which involve the production of a gas}) \begin{itemize} \item Measure the volume of the gas product at fixed time intervals. \item If \ce{CO2}: measure the loss in mass of the reaction mixture. \end{itemize} \subsubsection*{Factors which affect the rate of reaction} \begin{itemize} \item Concentration \item Temperature \item Surface Area \item Presence of a catalyst (can be organic or inorganic) \end{itemize} \begin{framed} \textbf{\large Note:} In order for a chemical reaction to take place, the particles must collide with sifficient force. Not all collisions produce a reaction --- the collision must have the right properties (force, type \&c.) \end{framed} \subsection{The Collision Theory} To understand why different reactions proceed at different rates, and why the rates are affected by concentration and other factors, we have to consider reactions at a mmolecular level. It is obvious that if two particles (atoms, molecules of ions) are to react they have to collide with each other, although, at ordinary room temperatures, only a very small proportion of molecular collisions result in reaction. The rate of a reaction will be increased by any change which makes the reactant particles collide more frequently, or which makes a greater proportion of collisions result in reaction. %Factors 1. - 4. will now be discussed ---- where does this go? \subsection{Concentration} The greater the concentration of any reactant (i.e.\ teh more particles per unit volume), the greater will be the number of collisions per second -- therefore the faster will be the reaction. For a gaseous reaction, an increase in pressure produces a decrease in volume and an increase in concentration. Therefore an increase in pressure makes a gaseous reaction faster. \subsection{`Lump' Size} For reactions involving solids, the smaller the `lump' size the faster the reaction --- this is because the smaller the `lump' size the greater the total surface area where the reactant particles can come into contact with each other. \subsection{Temperature} All chemical reactions are speeded up by an increase in temperature. A temperature rise means that the particles move faster. If they move faster then: \begin{itemize} \item They collide more often \item On average, the collisions are more energetic --- so that collisions are more likely to result in reaction (i.e.\ a greater proportion of all collisions result in reaction). \end{itemize} A temperature rise thus increases reaction rates in two ways, of which the second is the more effective. \subsection{Catalysts} A catalyst may be defined as a substance which alters (generally increases) the rate of a chemical reaction without itself being used up. \begin{itemize} \item \textbf{Chemical catalysts} --- These function in two ways: \begin{itemize} \item All catalysts change the `route' of reaction to one which requites less energy to `activate' it. The new route is therefore a faster one. \item Some catalysts also act by bringing together the reacting aprticles on to the surface of the catalyst, therefore increasing the reactant concentrations and the rate of reaction. \end{itemize} Examples of catalysts: \begin{itemize} \item Manganese(IV) oxide in the decomposition of Hydrogen peroxide into Oxygen and water. \item Iron in the manufacture (synthesis) of Ammonia (\ce{NH3}) from Nitrogen and Hydrogen --- the Haber Process \item Vanadium(V) oxide in the manufacture of Sulphuric acid, via Sulphur trioxide, from Sulphur dioxide and Oxygen --- the Contact Process \end{itemize} \item \textbf{Enzymes} \begin{itemize} \item Enzymes are biological catalysts and they speed up the chemical reactions which go on in living things \item They are always proteins \item They are specific in their action, in that each enzyme controls one particular reaction, or type of reaction \item They are destroyed by heat, this process being called \textbf{denaturation} \item Therefore they stop functioning above about 45$^\circ$C, working best over a limited range of temperature (35--40$^\circ$C approx.) \item When a reactant (\textbf{substrate}) molecule collides with a molecule of the correct enzyme, it can fit into a depression on the surface of the enzyme molecule, called the \textbf{active site}. The reaction then takes place and the molecules of product leave the active site, freeing it for another substrate molecule. \item The active site of a particular enzyme has a specific shape into which only one kind of substrate will fit. This is why enzymes are specific in their action. \end{itemize} Examples of enzymes: \begin{itemize} \item Catalase, present, e.g.\ in the liver, catalyses the decomposition of Hydrogen peroxide into Oxygen and water. This process is vital, as Hydrogen peroxide is poisonous (and is closely involved in the process of ageing) \item Zymase, present in yeast, catalyses the fermentation of glucose into alcohol (ethanol) \item Various protein-digesting enzymes (proteases), present in biological washing powders, remove protein stains, such as gravy, from clothes \end{itemize} \end{itemize} \section{Enthalpy} \subsection{Enthalpy Changes ($\Delta H$)} \subsection{Calculations Using Bond Energies} \section{Organic II} \subsection{Naming Organic Compounds} \subsection{Manufacture of Ethanol} \subsection{Esters} \subsection{Carboxylic Acids} \section{Electro-Chemistry} \subsection{Electrolysis} %\subsubsection{General Principle of Electrolysis} --- is this header really necessary? \begin{itemize} \item At the cathode (negative), metal or hydrogen is produced. \begin{itemize} \item Hydrogen is produced if the metal is \ce{Al} or higher. Metal is produced if it is \ce{Zn} or lower (all on the reactivity series). \end{itemize} \item At the anode (positive), non-metals (apart from Hydrogen) are formed. \begin{itemize} \item \ce{F2}, \ce{Cl2}, \ce{Br2}\ \ce{I2} are formed in preference to oxygen, unless the solution is dilute, in which case oxygen is formed from \ce{OH-} ions, according to the formula: \[ \ce{4OH-} \longrightarrow \ce{2H2O} + \ce{O2} + \ce{4e-} \] \item Nitrates and Sulphates do not decompose and Oxygen will be formed at the anode instead \end{itemize} \item Anything above Zinc in the reactivity series must be extracted from its natural ore using electrolysis. Zinc and below can be extracted by heating with Carbon. \item Electrolyte -- a substance that conducts electricity when dissolved in water or molten. \end{itemize} \subsubsection{Reactivity Series} \begin{table}[htb!] \centering \begin{tabular}{llc} Element & Mnemonic & Reactivity \\ \hline Potassium & Police & \textsc{high} \\ Sodium & Seargent & $\uparrow$ \\ Calcium & Charlie & \\ Lithium & & \\ Beryllium & & \\ Magnesium & M-- & \\ Aluminium & A-- & \\ $\uparrow$ \ce{H2} formed & & \\ (Carbon) & & \\ $\downarrow$ metal formed & & \\ Zinc & Z-- & \\ Iron & I-- & \\ (Hydrogen in Electrolysis) & & \\ Nickel & N-- & \\ Tin & T-- & \\ Lead & L-- & \\ Hydrogen & Has & \\ Copper & Caught & \\ Mercury & Me & \\ Silver & Stealing & \\ Gold & Gold & $\downarrow$ \\ Platinum & Plates & \textsc{low} \\ \hline \end{tabular} \linebreak \\[4pt] \parbox{0.75\textwidth}{\small More reactive elements prefer to \emph{stay} as ions when in solution, and are more likely to be formed at the cathode. In the reaction: \ce{X_{(s)} + Y+_{(aq)}} Where \ce{X} is \emph{more} reactive than \ce{Y}, the result will be: \ce{X+_{(aq)} + Y_{(s)}} (the charges could be $2+$, $3+$ \&c.)} \caption{Table of the Reactivity Series} \end{table} \begin{table}[htb!] \centering \begin{tabular}{lc} Element/Oxide ion & Reactivity \\ \hline \ce{I2} & \textsc{high} \\ \ce{Br2} & $\uparrow$ \\ \ce{Cl2} & \\ \ce{F2} & $\downarrow$ \\ \ce{O2} & \textsc{low} \\ \ce{NO3} & Never formed \\ \ce{SO4} & Never formed \\ \hline \end{tabular} \linebreak \\[4pt] \parbox{0.75\textwidth}{\small This assumes a concentrated solution; if dilute, it is more likely that simply Oxygen will be formed.} \caption{Non-Metal Reactivity Series} \end{table} \subsubsection{Metals} Metals are stable in a lattics structure where all their electrons are floating around freely (incidentally, this allows electrical conductivity). They are \emph{also} stable when they have a completely full/empty out electron shell. This makes them ions. So metals are found in nature either as ions or in in the lattice (sea of electrons) structure. (Carbon doesn't normally form ions --- it forms covalent bonds.) \begin{sidewaystable} \thispagestyle{empty} \centering \scalebox{0.9}{ \begin{tabular}{|rl||l|l|l|l|l|l|l|l|l|l|} \hline Name of Ion & $\longrightarrow$ & Fluoride & Chlorine & Hydroxide & Bromide & Oxide & Sulphide & Sulphate & Nitrate & Carbonate & Phosphate \\ $\downarrow$ & Symbol of ion & \ce{F-} & \ce{Cl-} & \ce{OH-} & \ce{Br-} & \ce{O^{2-}} & \ce{S^{2-}} & \ce{SO4^{2-}} & \ce{NO^{3-}} & \ce{CO3^{2-}} & \ce{PO4^{3-}} \\ \hline \hline Hydrogen & \ce{H+} & \ce{HF} & \ce{HCl} & \ce{H2O} & \ce{HBr} & \ce{H2O} & \ce{H2S} & \ce{H2SO4} & \ce{HNO3} & \ce{H2CO3} & \ce{H3PO4} \\ \hline Sodium & \ce{Na+} & \ce{NaF} & \ce{NaCl} & \ce{NaOH} & \ce{NaBr} & \ce{Na2O} & \ce{Na2S} & \ce{Na2SO4} & \ce{NaNO3} & \ce{Na2CO3} & \ce{Na3PO4} \\ \hline Potassium & \ce{K+} & \ce{KF} & \ce{Kcl} & \ce{KOH} & \ce{KBr} & \ce{K2O} & \ce{K2S} & \ce{K2SO4} & \ce{KNO3} & \ce{ K2CO3} & \ce{K3PO4} \\ \hline Lithium & \ce{Li+} & \ce{LiF} & \ce{LiCl} & \ce{LiOH} & \ce{LiBr} & \ce{Li2O} & \ce{Li2S} & \ce{Li2SO4} & \ce{LiNO3} & \ce{Li2CO3} & \ce{Li3PO4} \\ \hline Ammonium & \ce{NH4+} & \ce{NH4F} & \ce{NH4Cl} & \ce{NH4OH} & \ce{NH4Br} & \ce{(NH4)2O} & \ce{(NH4)2S} & \ce{(NH4)2SO4} & \ce{NH4NO3} & \ce{(NH4)2CO3} & \ce{(NH4)3PO4} \\ \hline Zinc & \ce{Zn^{2+}} & \ce{ZnF2} & \ce{ZnCl2} & \ce{Zn(OH)2} & \ce{ZnBr2} & \ce{ZnO} & \ce{ZnS} & \ce{ZnSO4} & \ce{Zn(NO3)2} & \ce{ZnCO3} & \ce{Zn3(PO4)2} \\ \hline Lead & \ce{Pb^{2+}} & \ce{PbF2} & \ce{PbCl2} & \ce{Pb(OH)2} & \ce{PbBr2} & \ce{PbO} & \ce{PbS} & \ce{PbSO4} & \ce{Pb(No3)2} & \ce{PbCO3} & \ce{Pb3(PO4)2} \\ \hline Magnesium & \ce{Mg^{2+}} & \ce{MgF2} & \ce{MgCl2} & \ce{Mg(OH)2} & \ce{MgBr2} & \ce{MgO} & \ce{MgS} & \ce{MgSO4} & \ce{Mg(No3)2} & \ce{MgCO3} & \ce{Mg3(PO4)2} \\ \hline Calcium & \ce{Ca^{2+}} & \ce{CaF2} & \ce{CaCl2} & \ce{Ca(OH)2} & \ce{CaBr2} & \ce{CaO} & \ce{CaS} & \ce{CaSO4} & \ce{Ca(No3)2} & \ce{CaCO3} & \ce{Ca3(PO4)2} \\ \hline Copper (II) & \ce{Cu^{2+}} & \ce{CuF2} & \ce{CuCl2} & \ce{Cu(OH)2} & \ce{CuBr2} & \ce{CuO} & \ce{CuS} & \ce{CuSO4} & \ce{Cu(No3)2} & \ce{CuCO3} & \ce{Cu3(PO4)2} \\ \hline Iron (III) & \ce{Fe^{3+}} & \ce{FeF3} & \ce{FeCl3} & \ce{Fe(OH)3} & \ce{FeBr3} & \ce{Fe2O3} & \ce{Fe2S3} & \ce{Fe2(SO4)3} & \ce{Fe(NO3)3} & \ce{Fe2(CO3)3} & \ce{FePO4} \\ \hline Aluminium & \ce{Al^{3+}} & \ce{AlF3} & \ce{AlCl3} & \ce{Al(OH)3} & \ce{AlBr3} & \ce{Al2O3} & \ce{Al2S3} & \ce{Al2(SO4)3} & \ce{Al(NO3)3} & \ce{Al2(CO3)3} & \ce{AlPO4} \\ \hline \end{tabular}} \caption{Table of Ions \& Ionic Compounds} \end{sidewaystable} \subsubsection{Molten Ionic Compounds} \subsubsection{Aqueous Solutions} \paragraph{Inert Electrodes} \paragraph{Metal Electrodes} \subsection{Producing Electricity From Chemical Reactions} \subsubsection{Diagram of Simple Cell} \begin{table} \begin{tabular}{lp{6cm}p{5cm}}\hline \it anion & \it test & \it test result \\ \hline carbonate (\ce{CO3^{2-}}) & add dilute acid & effervescence, carbon dioxide produced \\ aqueous chloride (\ce{Cl-}) & acidify with dilute nitric acid, then add aqueous silver nitrate & white precipitate \\ aqueous iodide (\ce{I-}) & acidify with dilute nitric acid, then add aqueous silver nitrate & yellow precipitate \\ aqueous nitrate (\ce{NO3^-}) & add aqueous sodium hydroxide, then aluminium foil; warm carefully & ammonia produced \\ aqueous sulphate (\ce{SO4^{2-}}) & acidify, then add aqueous barium nitrate & white precipitate \\ \hline \end{tabular} \caption{Tests for anions.} \end{table} \end{document}