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diff --git a/notes/notes.tex b/notes/notes.tex index 61bc1a8..ea8ad9d 100644 --- a/notes/notes.tex +++ b/notes/notes.tex @@ -662,7 +662,7 @@ colorlinks=true \subsubsection{Activation Energy} Activation energy ($E_A$) is the minimum amount of energy to initiate a chemical reaction (this is brought on by an increase in free energy of compounds during the reaction---meaning that energy is absorbed by the reaction before the ``hump'' of the activation energy is reached). This is a barrier to the reaction and must be overcome before the reaction will proceed. For all stable compounds, the reaction energy is positive because otherwise they are unstable and decompose readily (the higher the activation energy, the less likely the reaction is to occur and thus more slowly it occurs). \subsubsection{Chemical Bonds} - Chemical reactions rearrange atoms in molecules by making and/or breaking bonds. Bonds, as mentioned in previous units, are stable electron configurations---implying that energy must be input to start decomposing or making bonds. Note that energy still must be input to start most spontaneous reactions, but they will continue on their own because energy from previous reactions supplies activation energy for other reactant molecules. + Chemical reactions rearrange atoms in molecules by making and/or breaking bonds. Bonds, as mentioned in previous units, are stable electron configur\-ations---implying that energy must be input to start decomposing or making bonds. Note that energy still must be input to start most spontaneous reactions, but they will continue on their own because energy from previous reactions supplies activation energy for other reactant molecules. \subsubsection{Reaction Rate} Reaction rate is how quickly reactants convert to products. This depends inversely on activation energy. High activation energies (like in graphite to diamond reaction) mean low reaction rates, and low activation energies (like baking soda and vinegar) mean very high reaction rates. This is completely unrelated to $\Delta G_{rxn}$. \subsubsection{Collision Theory} @@ -670,7 +670,7 @@ colorlinks=true \subsubsection{Effective Collisions} Effective Collisions, as mentioned in the previous topic, are necessary (as compared to simply collisions where all possible touching of molecules results in a reaction) for chemical reactions to occur. What contrasts these effective collisions is a high enough kinetic energy and the reactants' facing the correct directions to exchange atoms. \subsubsection{Effect of Concentration on Reaction Rate} - Similar to temperature increasing collision rate by increasing pressure/temperature, concentration causes more frequent collisions---increasing the rate of reaction (concentration meaning number of molecules per volume in an aqueous solution, for example). + Similarly to temperature increasing collision rate by increasing molecule movement and pressure, increased concentration causes more frequent collis\-ions---increasing the rate of reaction (concentration meaning number of molecules per volume in an aqueous solution, for example). \subsubsection{Effect of Pressure on Reaction Rate} There are 2 other ways to increase pressure other than temperature: increasing the amount of reactants and decreasing volume (which makes sense with the ideal gas law). However, note that increasing the pressure is only useful with reactions exclusively between gases because liquids and solids don't become more dense with higher pressures. \subsubsection{Effects of Surface Area on Reaction Rate} @@ -750,6 +750,74 @@ colorlinks=true A bomb calorimeter is a specific type of calorimeter which doesn't use a bath of water (solution) for the reaction. Instead, it has a separated vessel inside of a bath of water to measure more reactions. This are often used to measure combustion heat release. \subsubsection{The Calorie Content of Foods} Nutrition information labels give calorie counts (note that a food Calorie is a kilocalorie or $4.184kJ$). These are determined with bomb calorimeters by burning the food and determining heat released during the combustion reaction. + \subsection{Thermochemical Equations} + \subsubsection{Enthalpy} + Bonds contain potential energy, and breaking or forming bonds can release or absorb energy. All reactants and products contain chemical energy. Enthalpy ($H$) measures the heat and internal energy in a system. Enthalpy in a chemical reaction is a state function, a type of function which depends only on the beginning and end of the process (not the middle). This makes sense because enthalpy of reaction is the same if a catalyst is used or not. Enthalpy is also a linear function, so the only important part is the potential energy of the involved reactants and products. This potential energy is called enthalpy of formation ($\Delta H_f$). It is defined relative to the zero-point of an element in a ``standard state'' at STP. Chlorine gas, for example has a $\Delta H_f$ of $0 \frac{kJ}{mol}$. + \subsubsection{Enthalpy of Reaction} + Enthalpy of a reaction ($\Delta H_{rxn}$) is the energy absorbed (positive) or released (negative) during a chemical reaction. Hess's Law of heat formation states that this property is linear with enthalpy of formation, i.e. the enthalpy of a reaction is the enthalpy of the reactants subtracted from the enthalpy of the products. A thermochemical equation is a chemical equation which includes this information and states of involved substances. + \subsubsection{Manipulating Equations} + If the direction of a chemical equation is reversed, $\Delta H_{rxn}$ becomes $-\Delta H_{rxn}$. If a chemical equation is multiplied by a coefficient $k$, $\Delta H_{rxn}$ becomes $k\Delta H_{rxn}$. Note that $\Delta H_{rxn}$ typically refers to integer molar amounts of compounds in a balanced chemical equation. + \subsubsection{Enthalpy of Combustion} + Combustion, another chemical reaction, also has a certain amount of enthalpy. Combustion reactions are typically very formulaic, so the energy of a given combustion reaction (such as burning propane) is known by the scientific community. Note, however, combustion enthalpy can still be determined with standard techniques. + \subsection{Enthalpy and Phase Changes} + \subsubsection{Heat and Phase Changes} + Previously, heating and cooling curves have been described. These are graphs of a substance across temperature and added heat, showing its phase changes. While in different states of matter, substances have different specific heats. While changing between substances, temperature is constant, but energy is not. For melting, the energy to transform a mole of a substance at the melting point, but completely solid, is called the molar heat of fusion ($\Delta H_{fus}$). The heat of fusion can also be given in terms of $\frac{J}{g}$ but is typically provided as $\frac{J}{mol}$. The energy released during the opposite process, freezing, is equal to $-\Delta H_{fus}$. Boiling/vaporizing follows a similar pattern: the molar heat of vaporization is $\Delta H_{vap}$, and the molar heat of condensation (not a real measurement) is $-\Delta H_{vap}$. These can also be given in terms of mass. + \subsubsection{Effects of Heat on Phase Changes} + During phase changes, temperature doesn't change, so kinetic energy doesn't change. However, heat input has to increase some form of energy, so it increases potential energy---enough so that the movement of the particles overcomes the intermolecular forces. Similar to chemical reactions, potential energy increases during endothermic reactions and decreases during exothermic reactions. + \subsubsection{Relationship between Enthalpy} + $\Delta H_{vap} > \Delta H_{fus} > 0$ + \subsubsection{Sweat and Temperature Regulation} + Because evaporation is endothermic, the human body uses sweat to cool itself by making sweat evaporate (temperature doesn't increase because no temperature changes occur during phase changes). + \subsection{Enthalpy of Reaction} + \subsubsection{Equation Combination} + By Hess's Law of Heat Formation, chemical equations can be added like linear equations. This is most useful for combined processes (like writing photosynthesis from its component reactions). The goal of combining equations is to write a balanced equation of the products respective to reactants as chosen. + \subsubsection{Diagramming Hess's Law} + Because enthalpy is a state function, and the potential energy of a given compound or group of compounds is constant, a set of chemical equations can be diagrammed as a set of transformations from one set of compounds (i.e. $2C(s)+H_2(g)$ to $C_2H_2(g)$ is one transformation), and from that, the difference in the final product's potential energy and the reactants is the change in enthalpy. Note also that upwards arrows on the diagram are positive, and downward arrows are negative. + \subsubsection{Reaction Energy Diagrams and Enthalpy Diagrams} + Reaction energy diagrams and enthalpy diagrams are both ways to describe the change in potential energy of a chemical reaction. Enthalpy diagrams are more detailed in the process, but reaction energy diagrams describe reaction progress with more detail. Also, reaction energy diagrams only display one reaction at a time unlike enthalpy diagrams displaying the total and individual reactions. Both put energy on the vertical axis. +\section{Mixtures, Solutions, and Solubility} + \subsection{Mixtures and Solutions} + \subsubsection{Heterogenous and Homogenous Mixtures} + As described extensively in previous lessons, homogenous mixtures are approximately congruent between any two arbitrarily chosen volumes, but heterogenous mixtures are not. + \subsubsection{Solution} + A solution is a type of homogenous mixture which appear as one phase. It has two components: a solvent (the ``dissolver''), and a solute (the ``dissolved''). + \subsubsection{Suspensions and Colloids} + Suspensions are similar to solutions, but instead of being homogenous they are heterogenous mixtures which have large enough particles to settle out or be filtered out. This includes muddy rivers, milk, and blood. Colloids are a subset of suspensions with significantly smaller particles dispersed in a way that makes it difficult to easily filter or quickly settle. + A few features of suspensions, colloids, and suspensions can be used to differentiate them. The first of these is Brownian motion: the constant random motion of particles is observed in colloids and solutions but not in suspensions. Quick settling is observed in suspensions by definition but not in most colloids or any solutions. + \subsubsection{The Tyndall Effect} + The Tyndall Effect differentiates colloids from solutions because they can be easily confused. In colloids, light passing through is scattered by the particles like dust in air or flour in water. This scattering, the Tyndall Effect, notably does not occur in solutions. The scattering of light looks like a coherent beam of light. + \subsubsection{Techniques to Separate Solutions} + There are a few common techniques to separate solutions. These have been listed before, and include distillation (boiling liquids and collecting the vapor of liquids as they boil at different temperatures), crystallization (vaporizing the solvent to leave the solute(s)), and chromatography (separating solutes by density or particle size). + + Colloids are slightly easier to separate. Centrifuges (spinning the colloid to separate solutes by density with centrifugal force), long standing (leaving the solute to settle out of the solvent), and boiling/heating (heating or running electricity through the colloid to coagulate---thicken to a solid or semisolid---it). + \subsection{Reactions in Aqueous Solutions} + \subsubsection{Polarity of Water} + Water is a highly polar covalent compound. The Oxygen has a partial negative charge, and the Hydrogens/hydrogen end have a partial positive. This makes it better as a dissociative solvent for ionic compounds. Dissociation is when ionic compounds separate into their respective ions with the same charges they have while in the compound. These individual atoms (like $Na^+$ or $Br^-$ from $NaBr$) and their parent compounds (such as $NaBr$) are called electrolytes. + + Dissociation occurs in significant amounts in aqueous solutions because of a process called hydration. Water molecules' partially negative oxygen atoms surround positive ions (cations) like $Na^+$ to help them stay in solution and prevent reassociation. In order to cancel the electric charge, more than one water molecules surround the ions. Similarly, the partially negative hydrogen side surrounds negative ions (anions) like $Cl^-$. + \subsubsection{Ionization} + Water can also dissolve some covalent compounds by ionizing them---a chemical reaction. For example, sulfur dioxide ($SO_2$, a covalent compound) takes $OH^-$ from the water molecule to make itself a negative ion. This also occurs with ammonia and hydrochloric acid. Note that this, like dissociation, only occurs to water-soluble molecules. + \subsubsection{State of Matter in Chemical Reactions} + As described in the lesson on chemical equation symbols, symbols like $(g)$ and $(l)$ describe the state of matter of reactants and products (gas and liquid, respectively), and $(aq)$ describes reactants or products dissolved in an aqueous solution. Where insoluble products form (remaining as $(s)$ instead of becoming $(aq)$ like with $AgCl$), gases either bubble out of the solution and solids form precipitates separate from but within the solution. + \subsubsection{Double-Displacement Reactions} + Double-displacement reactions occur when ions switch places in ionic compounds. These often occur in aqueous solutions, and when they do, ionic compounds in aqueous solution can be represented as their individual ions in aqueous solution (e.g. $NaCl(aq)=Na^+(aq) + Cl^-(aq)$ but $AgCl(s) \neq Ag^+(s) + Cl^-(s)$). + + After these dissociated ions are rewritten, any ion which is on both sides in the same phase can be eliminated (e.g. $H^+(aq) + Cl^-(aq) + Na^+(aq) + OH^-(aq) \longrightarrow Na^+(aq) + Cl^- + H_2O(l)$ becomes $H^+(aq) + OH^- \longrightarrow H_2O(l)$). This is called a net ionic equation. Notably, if all reactants and products are aqueous in a double-displacement reaction, the net ionic equation is empty and thus no real reaction has occurred. This leads us to show that double-displacement reactions in solution occur only when a gas, precipitate, or pure liquid are formed. + \subsubsection{Types of Electrolytes} + Strong electrolytes and weak electrolytes are the two types of electrolytes. Strong electrolytes dissociate or ionize completely like $HCl(g)$ or $KOH(s)$. This means that they conduct electricity very well when in solution. Weak electrolytes, on the other hand partially dissociate or ionize because their reactions are easily reversible---meaning that a significant amount of reactant is left in the solution. Nonelectrolytes aren't really electrolytes beccause they don't dissociate or ionize in solution. These solutions do not conduct electricity. + \subsection{Solutions and Solubility} + \subsubsection{The Dissolving Process} + Dissolution does not occur instantaneously. It's a process which requires that solute molecules or ions are attracted to solvent molecules, causing the solvent to try and surround the solute. This pulls apart the solute from the other solute molecules, and completes the mixing process. The necessary attraction between solute and solvent is why the ``like dissolves like'' rule tends to accurately describe dissolution of solutions (e.g. water, a highly polar covalent molecule easily dissolves highly polar ionic compounds). + \subsubsection{The Rate of Dissolution} + The rate of dissolution is how quickly a solid solute dissolves in a liquid solvent. This increases proportionally with stirring, surface area, and temperature. + \subsubsection{Solubility} + Solubility is the ``amount of solute which will dissolve in a volume of solvent at a given temperature and pressure.'' This changes only with the solute's and solvent's identity. This is different from the dissolution rate and follows the like dissolves like rule significantly. + + The concentration of a solute compared to the maximum concentration based on solubility defines three types of solutions: unsaturated solutions have less solute than the ``maximum,'' saturated solutions have exactly an equal amount, and supersaturated solutions have higher than the maximum. Supersaturated solutions are highly unstable and are typically created by slowly cooling a saturated solution at a higher temperature. + \subsubsection{Environmental Conditions' Effects on Solubility} + Increasing temperature increases the solubility of solids and liquids but decreases the solubility of gases. Increasing pressure, however, increases the solubility of gases and does not affect the solubility of solids and liquids. + + Solubility graphs show the change in solubility with temperature. These are determined experimentally because they are different for different solutes. These can be used to determine if a given solution is saturated, unsaturated, or supersaturated at a specific temperature. \end{document} |