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         Some possible indicators of a chemical change (meaning that these indicators do not conclusively show that a given change is chemical) include color change, solid disappearance, gas formation, precipitate formation, and light and heat production.
 
         Physical changes are of the sort of melting, boiling, cutting, and bending. Chemical changes are of the sort of burning, cooking, water reactions, and air reactions.
+\section{Chemical Bonds and Compounds}
     \subsection{Ionic Bonding}
         \subsubsection{Ionic Bonding}
 	Ionic bonding is a relatively simple type of bond where the majority of electrons stay in the valence of one atom. This definition defines the restriction of difference in electron negativity: any two elements can form an ionic bond if the difference of their electron negativities $>1.7$. The charge of the ions of a given element also make a difference. In addition, only elements with different ion charges (for example, $Al^{+3}$ can form a bond with $O^{-2}$, but not $Ca^{+2}$).
@@ -506,6 +507,118 @@ colorlinks=true
 	\subsubsection{Alloys}
 	Alloys are a type of mixture. Alloys are mixtures made by melting together metals and are thus homogenous. Properties of alloys are different from the properties of its composite metals, which is what makes them useful.
 
-	Alloys include brass (copper \& zinc), rose gold (gold \& copper), bronze (copper \& tin or aluminum), and steel (mostly iron, chromium, \& nickel). 
+	Alloys include brass (copper \& zinc), rose gold (gold \& copper), bronze (copper \& tin or aluminum), and steel (mostly iron, chromium, \& nickel).
+    \subsection{Intermolecular Forces}
+        \subsubsection{Intramolecular vs Intermolecular Forces}
+	Intramolecular forces are, by definition, acting within a molecule. These are chemical bonds (covalent and ionic). Intermolecular forces act between molecules in a substance, attractively or repulsively. These can affect interactions between molecules and have 3 classes: hydrogen bonds, London dispersion forces, and Van der Waals forces. Intermolecular forces are significantly weaker than intramolecular forces.
+	\subsubsection{Hydrogen Bonds}
+	A hydrogen bond is "the attraction of a single hydrogen between a highly electronegative atom and another highly electronegative atom in a different molecule or a different region of a large molecule. These are commonly seen in water because of the strongly electronegative nature of the oxygen. It pulls electrons away from the hydrogen, creating a charge differential which acts like a hydrogen bond between hydrogens and oxygens of different molecules. Note that hydrogen bonding only occurs when hydrogens are directly bonded to highly electronegative atoms.
+
+	Hydrogen bonding affects the properties of substances which affect significantly because it is only roughly $\frac{1}{10}$ the strength of a real bond. The properties which it creates (when present) include higher melting and boiling points, high surface tension, high viscosity, highly structured solid forms (less dense than liquid form), and in some cases (such as water) polarity. This polarity can lead to macroscopic solubility properties. Ethane, a nonpolar, non hydrogen bonding chemical can only dissolve nonpolar covalent substances in contrast to water being able to dissolve ionic and polar covalent substances readily.
+
+	These are very common in organic molecules because oxygen and hydrogen are often present as major building blocks.
+	\subsubsection{Van der Waals Forces}
+	Van der Waals forces, unlike hydrogen bonds, occur in all atoms. These are the attractive or repulsive forces created by uneven distribution of electrical charge. This uneven distribution can occur from random variation, and it can occur from differences in atoms' electronegativities. Between polar molecules, a Van der Waals force is present, attracting the negative pole of one atom to the positive pole of the other. This is called a dipole-dipole interaction. Another common Van der Waals force is dipole-induced dipole interaction between polar and nonpolar molecules. The positive or negative end of the polar molecule pushes the same force away (negative repels electrons, positive attracts electrons), pulling them together. This induced force is weaker than a dipole-dipole interaction. The third Van der Waals interaction is called a London dispersion force. This is a induced dipole-induced dipole interaction. While it occurs in all atoms, it is only significant in nonpolar-nonpolar interactions because of how weak it is. London dispersion forces cause molecules to attract.
+	\subsubsection{Applications of Intermolecular Forces}
+	As these intermolecular forces fundamentally change the properties of materials, they can be used to stick to surfaces (such as in geckos' feet), or design soap which disrupts water's intermolecular forces giving it a high surface tension, melting point, and density.
+    \subsection{Nomenclature of Ionic Compounds}
+        \subsubsection{Charge of Ions}
+	When forming ionic compounds, one of the atoms will necessarily lose electron(s) to the other, and because this usually aims for a full valence shell the number is typically the same between different bonds. The number which it loses or gains is called its oxidation state, and varies between atoms. Metals' oxidations are +1 for group 1, +2 for group 2, +3, and multiple, variant oxidation states for transition metals. For nonmetals in group 16 and 17, the oxidation states are -2 and -1 respectively.
+
+	Ionic compounds are neutral, allowing for algebraic determination of ion ratios/quantities. For example, if element A has an oxidation state of A and element B has an oxidation state of B, Ax+By=0 where x is the number of A ions and B is the number of B ions. This works consistently.
+	\subsubsection{Nomenclature of Ionic Compounds}
+	Nomenclature is the rules for naming compounds, as develoepd by the International Union of Pure and Applied Chemistry (IUPAC). For ionic compounds, it is always [cation metal] + [anion nonmetal modified with -ide]. This can also be performed in reverse: if one takes the name of an ionic compound, the component ions can be determined and from that the elements, ratio, and formula in that order.
+	\subsubsection{Polyatomic Ions}
+	This is the same as monoatomic ionic compounds, but without the -ide modification for anions. For example, $(NH_4)_2CO_3$ is ammonium carbonate.
+	\subsubsection{Transition Metals}
+	As mentioned earlier, transition metals are an exception to the singular oxidation states, and can each form multiple ions with different charges because the d and f orbitals allow different stable electron configurations. These have no easily discernible pattern, so it is easiest to memorize or look up oxidation states of a given transition metal. When converting chemical formulas with transition metals to the name of that ion, one uses roman numerals to signify the oxidation state. For example, in $MnO$, the $Mn$ has a charge of $+2$ because oxygen always has an oxidation state of $-2$ and ionic compounds have 0 net charge, this meaning that the name is ``magnesium (II) oxide.'' Note that roman numerals ought be used only when the metal ions have multiple common charges.
+    \subsection{Nomenclature of Covalent Compounds}
+        \subsubsection{Nomenclature for Covalent Compounds}
+	The element further left on the periodic table is listed first (e.g. $CO_2$ is carbon dioxide, not oxygen carbide). The second element ends with -ide (e.g. carbon dioxide). Numerical prefixes are used to indicate the number of atoms of a given element (1 of the second element=mono, 2=di, 3=tri, 4=tetra, 5=penta, 6=hex) such as in carbon tetrachloride or dinitrogen tetroxide.
+	\subsubsection{Exceptions}
+	The most major exception to these rules are traditional names such as water instead of dihydrogen monoxide, where the traditional name would be preferred. The other exception are compounds with hydrogen. In hydrogen compounds, a prefix for the amount of hydrogen isn't used such as in hydrogen sulfide ($H_2S$) or hydrogen sulfate ($H_2SO_4$). However, these chemicals are often acidic, which means they follow unique naming rules.
+	\subsubsection{Acids and Bases}
+	Acids increase the concentration of $H^+$ ions in a water (aqueous) solution. In the same solution, bases increase the concentration of $OH^-$ ions. Both can be either ionic or covalent, but they have significant structural differences: acids generally contain H atoms bonded to other atoms or polyatomic ions and ionic bases contain $OH^-$ in their formula. Note that covalent bases do not contain $OH^-$. Acids turn blue litmus paper red while bases turn red litmus paper blue. Tastewise, acids taste sour and bases taste bitter. Bases also feel slippery or soapy.
+
+	Acids have two different methods by which they are named, depending on what type of acid they are. Binary acids (acids with exactly two atoms) are \textit{hydro-[nonmetal]-ic}. Oxyacids (acids with polyatomic ions, in turn with one or more oxygen atoms) change dependent on the type of polyatomic ion. \textit{-ate} ions form \textit{-ic} acids and \textit{-ite} anions form \textit{-ous} acids (e.g. nitrate -> nitric acid and nitrite -> nitrous acid).
+
+	Bases' nomenclature is slightly less formulaic. Most covalent bases are amines, meaning that they are similar in structure to ammonia, but instead of a hydrogen, an R-group (a carbon or string of carbons) is present. The number of R-groups changes the name, such as primary amines (1 r-group) or tertiary amines (3 r-group). A hydrogen ion can attach to the Nitrogen's extra pair of electrons. Ionic bases follow the same rules as other ionic compounds.
+    \subsection{Polymers}
+        \subsubsection{Formation of Polymers}
+	Monomers are the basic units of polymers---simple molecules like ethylene which can combine with eachother by covalent bonding. Monomers share electrons (such as deconstructing a double bond to form a single bond to another atom) in order to form these bonds. Because they can each form covalent bonds with more than one other monomer, these can occur as repeating subunits of a large molecule (polymer). The process of forming this large molecule (also known as a macromolecule) is called polymerization. The properties of a given polymer are, notably, dissimilar from its constituent monomers. Ethylene, for example, is a colorless, flammable gas with a sweet taste and odor, but as part of a polymer, this forms the plastic of water bottles.
+	\subsubsection{Polymer Classification}
+	There are two types of polymers: natural and synthetic. These are precisely as they sound, made in nature and manmade respectively. Both are immensely useful in technological applications. For natural polymers, these make gloves rubber bands (from natural rubber), gloves, scarves, carpets (from wool), ties, and silk fabrics (silkworm cocoons). DNA (carries genetic information), starch (food and paper-making), and cellulose (plant stems, and used as thread such as cotton jeans) are three more incredibly common natural polymers.
+
+	Synthetic polymers uses include: polysterene foam (makes take-out containers, packaging materials, egg cartons, and insulation), nylon (used to make ropes, nets, backpacks, cooking utensils, and stockings), polyethylene (plastic bags, bottles, toys), and vulcanized rubber (tires, shoe soles). The reason that these are so useful in contrast to natural polymers is because they are flexible, hard, lightweight, strong, resistant, and often extremely cheap to produce. However, these are often not biodegradable, causing the synthetic polymers to build up in landfills as recycling is expensive. Also, the materials used to make synthetic polymres aren't always available.
+    \subsection{Properties and Uses of Saturated Hydrocarbons}
+        \subsubsection{Isomers and Hydrocarbons}
+	Isomers are chemicals which are similar in one of two properties (if they are similar in both, they are the same chemical). These are structural isomers (same chemical formula) and geometric isomers (the structure looks the same, but the numbers are different). Hydrocarbons are chemicals with exclusively hydrogens and carbons, and saturated hydrocarbons are hydrocarbons with only single bonds. These are also known as alkanes. In contrast, unsaturated hydrocarbons contain at least one double or triple bond.
+	\subsubsection{Hydrocarbon Notation}
+	Hydrocarbons and organic compounds in general are often represented with standard Lewis structures, but there are some notational abbreviations which can be taken when drawing the chemicals. The first of these is abbreviation of methyl groups ($CH_3$, $CH_2$ and $CH$; note that $CH_4$ is never present in hydrocarbons with more than a single $CH_4$ because all of its atoms have full valence shells, leaving no more room for bonding). These methyl groups can also be written in opposite chemical order such as $H_3C$ for convenience. The other notational abbreviation is a methyl group being written as a "joint" (no $C$, just connecting two lines or being the end of a line). Note that when this last notational trick is used in conjunction with double or triple bonds, the extra line(s) are written adjacent to where they would normally connect, not touching anything.
+	\subsubsection{Straight Chains, Branched Chains, and Cycloalkanes}
+	All compounds listed in this lesson are alkanes. Straight chains are the simplest type of alkane, as a singular straight line of carbons (with attached hydrogens). Branched chains are slightly more complex, with more than 2 "ends" (i.e. carbons attached to only one other carbon). Cycloalkanes have no "ends," forming rings of carbon.
+	\subsubsection{Properties}
+	These are generally:
+	\begin{itemize}
+            \item Fairly unreactive
+	    \item Combustible
+	    \item Insoluble in water but soluble in ether and other organic solvents
+	    \item having a higher boiling point with increased molecular mass.
+	\end{itemize}
+	\subsubsection{Alkane Uses}
+	Alkanes have several uses, often in mechanical applications. These include gasoline (pentane and octane), natural gas (methane and ethane), lubricating oil (carbon lengths of 17-35 carbons), and asphalt (carbon lengths $>35$ carbons)
+	\subsubsection{Nomenclature}
+	For the 10 smallest straight chains, the naming system is a simple $[prefix]-ane$. For straight chains with 5 or greater carbons, the prefix follows Greek numerical prefixes, but the first 4 are methane, ethane, propane, and butane in that order.
+	Branch chains have a multi-step naming system: one starts by naming the longest chain in the molecule and naming that (i.e. nonane or methane). Then, the substituents (chains coming off of the main chain) are identified and labeled (such as methyl groups).Next, all carbons on the long chain are numbered sequentially---starting at the end with the closest branch. After that, substituent groups are listed alphabetically (e.g. methyl before nitrogen before oxygen), and included in the name based on the location of the substituent groups. This becomes a molecule name such as 3,6-dimethylnonane if two methyls are attached to carbons 3 and 6.
+	Cycloalkanes have a similar naming system, but with a couple of exceptions. Instead of starting the counting at the end, the counting starts at the longest substituent (because there is no end) and continues in the direction of the closest substituent. Numbering continues otherwise as normal, using standard cycloalkane names such as cyclopentane or cyclobutane, unless there is only one substituent. In the case of one substituent, no numbering is used. The name of the substituent is simply prefixed.
+    \subsection{Properties and Uses of Unsaturated Hydrocarbons}
+        \subsubsection{Alkenes, Alkynes}
+	Alkenes are hydrocarbons with at least one double bond between carbon atoms. The chemical formula has the form $C_nH_{2n}$ assuming only one double bond. Their naming convention is the same as alkanes except with the suffix "-ene." Alkynes have at least one triple bond between carbon atoms. They use the form $C_nH_{2n-2}$, and their naming convention is the same as alkanes except with "-yne." Both have a number prefixed to indicate where the double/triple bond is if the chain is long. These can be cyclic.
+	\subsubsection{Cis Isomers and Trans Isomers}
+	Because double and triple bonds use pi bonds, the molecules don't have freedom of movement about these bonds, meaning that there are two possible positions in which the adjacent carbons can occur relative to eachother. These are cis isomers (the next carbon is in the same direction on both sides) and trans isomers (the next carbons are in opposite directions). Trans isomers are relatively uncommon. These are prefixed to compounds for identification
+	\subsubsection{Aromatic Hydrocarbons}
+	These are the double-bonded analog of cyclohexanes. The aromatic hydrocarbons always have 6 bonds in their basic unit (benzene), which alternates between double and single bonds. These are flat molecules, and can form into structures with multiple benzene units (such as naphthalene).
+	\subsubsection{Boiling Points and Vapor Pressure}
+	The boiling points of alkenes are lower than alkynes, and for both, as the boiling point increases with chain length. As mentioned in earlier lessons, vapor pressure is the pressure exerted by a vapor when the liquid and vapor are in equilibrium at a given temperature. For alkenes, vapor pressure decreases with the size of the molecule increasing, and like most substances increases with temperature.
+	\subsubsection{Properties of Unsaturated Hydrocarbons}
+	Almost all are combustible and insoluble in water. Alkynes have typically higher boiling points than alkenes or aromatics. These are all less dense than water, and alkenes tend to be pH-neutral while alkynes are typically acidic.
+	\subsubsection{Applications of Unsaturated Hydrocarbons}
+	Unsaturated hydrocarbons are sometimes present in food such as in polyunsaturated margarine or aspirin and other analgesic medicines (they contain chemicals derived from aromatic hydrocarbons). Rubber is made up of alkene monomers, polysterene is made of hydrocarbons, and acetylene is an alkyne used as fuel in welding torches, gas lamps, and electric generators.
+\section{Chemical Reactions and Limiting Reactants}
+    \subsection{Evidence of Chemical Reactions}
+	\subsubsection{Chemical Changes and Chemical Reactions}
+	Chemical reactions accompany all chemical changes, and are effectively equivalent. Chemical reactions, by definition, are a rearrangement of the atoms in reactants which result in a new substance (product). Oxidation as in rust, combustion, and rising bread are all chemical reactions. An arrow ($\longrightarrow$) is often used to represent the change from listed reactants to listed products.
+	\subsubsection{Rate of a Chemical Reaction}
+	Chemical reactions can occur over a short period of time (fast) such as a firework, or they can take a long period of time (slow), such as limestone decaying. The speed at which this occurs is called the rate.
+	\subsubsection{Indicators/Evidence of a Chemical Reaction}
+	As noted in previous lessons, when a chemical reaction occurs, several things can happen which are common between chemical reactions. These include gas formation, solid (precipitate) formation, a color change, and always some form of energy change. The enrgy change can change the temperature, a decrease if the reaction is endothermic and an increase if the reaction is exothermic. Light is also emitted as another signal of energy change.
+
+	These indicators do not prove a chemical reaction as a hard and fast rule. Physical changes can be accompanied by gas formation (mentos and coke), solid formation (freezing), temperature change (boiling water), and color change (mixing dyes). Generally, the more indicators of a chemical reaction which are present, the more likely a chemical reaction has occurred.
+    \subsection{Writing and Balancing Chemical Equations}
+        \subsubsection{Equations}
+	Equations represent a given chemical reaction (changing listed reactants on the left to listed products on the right, separated by pluses in both cases). These equations can represent each substance with words (such as water or sodium bicarbonate) or chemical formulae (such as $H_2O (l)$ or $12NaHCO_3 (g)$). In special cases, a "model" (pictographic representation) of the formula can be used.
+
+	The formulae are the most common and accepted. The letters next to the chemicals' formulae represent the state (s is solid, l is liquid, g is gas, aq is aqueous). The numbers before the chemicals represent "counts" (ratios) of moles/molecules and allow balancing by using algebra and the principle of conservation of matter (if $O_2$ is at left and $H_2O$ is at the right, with no other oxygens, there will be the same number of oxygens on each side (double at the right)). The smallest possible numbers are preferrable.
+
+	Often, reaction conditions are added above the arrow or the arrow becomes bidirectional to symbolize a reversible reaction. The reaction conditions are mostly self explanatory and there are 4 common symbols: a $\Delta$ for a heated reaction, $25^\circ$ for a reaction at 25$^\circ$, $2atm$ for a reaction at 2 atmospheres, and $Pt$ for platinum as a catalyst.
+    \subsection{Percent Composition and Molecular Formula}
+        \subsubsection{Mole Ratios}
+	A mole is a specific number of a given particle (ion, atom, molecule, etc.). The number is Avogadro's number and is chosen so that the mass of a mol of $^{12}C$ is 12g. Because this isn't reliant on mass, just the amount of particles every chemical has a molar ratio between its composite elements easily determinable from the empirical formula. $CO_2$, for example, has a mole ratio of $C:O=1:2$. Note that this isn't always just the numbers of elements listed in the formulae. $N_2O_4$ has a mole ratio of $N:O=2:4=1:2$, with $1:2$ preferred as the simplest ratio and $2:4$ as the true ratio. These ratios also need not always include every element. For ammonium nitrate ($NH_4NO_3$), the $H:O$ ratio is $4:3$.
+	\subsubsection{Empirical and Molecular Formulae}
+	The true ratios and mole ratios mentioned earlier generate molecular and empirical formulae respectively (e.g. a given ratio of elements such as $C:H=1:2$ becomes a formula such as $CH_2$). This implies that all empirical formulae have the simplest ratio of their elements, but molecular formulae can have any multiple of a simple ratio.
+	\subsubsection{Percent Composition}
+	In order to accurately analyse compounds, it can be useful to know the proportions of masses of an element. This is not computationally difficult as the percent composition of a given element in a given compound is the mass of the element in the compound divided by the mass of the compound. Note that for determining this ratio molecular and empirical ratios are equally well-suited. These ratios can also be determined for more complex chemicals than just elements. In hydrates (crystals which integrate water in their lattice), one can determine the percent composition of water by an identical calculus. Note that in a hydrate, the salt forming the crystal is called anhydrous when the crystal doesn't have water.
+    \subsection{Types of Reactions}
+        \subsubsection{Types of Chemical Reactions}
+	There are five main types of chemical reactions: synthesis (2+ reactants combine to form a single product), decomposition (a single compound breaks down to form 2+ new substances), combustion (reaction of a substance with oxygen---includes oxydation), single replacement (reaction in which one ion displaces another to form a new compound), and double replacement (two ionic compounds exchange ions to form new products)
+	\subsubsection{Differentiating Single and Double Replacement Reactions}
+	In single replacement, both the reactants and products have one element and one compound, where the reactant element becomes part of the product compound. In double replacement, reactants and products are ionic compounds and the same ions are present---just reordered---between the reactants and products.
+	\subsubsection{Activity Series}
+	The activity series describes how reactive elements are compared to eachother. For example, lithium is much more reactive than gold. This can allow us to make predictions about if a reaction will take place or not. For example, we know that $AgNO_3+Cu$ are more likely to react than $Cu(NO_3)_2+Ag$ because copper in its pure form is more reactive than silver is in its pure form. This generalises. Note that halogens' and metals' activity series are different. Metals' activity series is long, but the halogens' activity series is short ($F_2 > Cl_2 > Br_2 > I_2$).
+    \subsection{Limiting Reactant and Percent Yield}
+        \subsubsection{Limiting Reactants and Stoichiometry}
+	Stoichiometry is the relationship between the relative quantities of substances taking part in a reaction or forming a compound, typically a ratio of whole integers. Stoichiometry is typically used in chemistry to determine the ratios between substances in which they react. For example, $H_2$ and $N_2$ react to form $NH_3$ in a ratio of $H:N=3:1$. However, if there are 5mol of $H_2$ and 1mol of $N_2$, then there is left 2mol of $H_2$, called the excess reactant. Because more nitrogen would allow continued reaction, nitrogen is called the limiting reactant. In general, the limiting and excess reactants in a reaction with uneven proportions can be determined with a similar calculus. Stoichiometry also extends to comparisons between reactants and products. If $2O_2 + H_2O \rightarrow H_2O_5$, 10 fully reacted moles of oxygen form 5 moles of the substance on the right (note that the substance makes no sense and isn't real).
+	\subsubsection{Theoretical Yield and Percent Yield}
+	Theoretical yield is ``the ideal maximum amount of a product that can be produced during a reaction, calculated from stoichiometric relationships.'' All examples above assumed that reactions occurr perfectly, consuming all of the reactants and creating the correct amount of product and nothing else. The amount of product created was the theoretical yield. However, real life doesn't usually have the same yield (actual yield) as predicted. Actual yield is almost always less than theoretical yield due to suboptimal reactant conditions and random factors. Percent yield is the ratio between actual yield and theoretical yield.
 \end{document}