8 files changed, 227 insertions, 14 deletions
diff --git a/cer/.DS_Store b/cer/.DS_Store
index 6ce9082..24c39b5 100644
--- a/cer/.DS_Store
+++ b/cer/.DS_Store
diff --git a/cer/sem2/cer.tex b/cer/sem2/cer.tex
index c8331ef..a55a489 100644
--- a/cer/sem2/cer.tex
+++ b/cer/sem2/cer.tex
@@ -1,4 +1,5 @@
-\font\tinyx=cmr8 at 6pt \font\tinym=cmmi8 at 6pt \font\tinyms=cmsy8 at 6pt \def\tiny{\tinym \tinyms \tinyx \baselineskip=6pt}
+\font\tinyx=cmr8 at 6pt \font\tinym=cmmi8 at 6pt \font\tinyxs=cmmi8 at 5pt \font\tinyms=cmmi8 at 5pt
+\def\tiny{\textfont0=\tinyx \textfont1=\tinym \scriptfont0=\tinyxs \scriptfont1=\tinyms \tinyx \baselineskip=6pt}
 \font\headx=cmb8 at 9pt \font\headm=cmmi8 at 9pt \font\headms=cmsy8 at 9pt \def\head{\headm \headms \headx \baselineskip=9pt}
 %\font\bigx=cmb14        \font\bigm=cmmi14        \font\bigms=\def\big{\bigx \bigm \baselineskip=14pt}
 
@@ -60,18 +61,6 @@
   \def\preamb{#1}
 }
 
-%\def\cp#1{#1}
-%\def\preambloop{
-%  & \vrule \text{##}
-%  \advance \@columns by -1
-%  \ifnum \@columns>0 \span\preambloop \fi
-%}
-%\def\makepreamb{
-%  \vrule width 1pt \text{##}
-%  \advance \@columns by -1
-%  \ifnum \@columns>0 \span\preambloop \fi
-%  \vrule width 1pt \cr
-%}
 
 \def\makedata{
   \def\text##1{\tiny \tolerance=10000 \hbadness=10000 \hbox to \@tablewidth{\hskip1em \vbox{\vskip1ex \noindent \advance \@tablewidth by -2.25em \hsize\@tablewidth ##1 \smallskip} \hskip1em} }
@@ -88,8 +77,9 @@
   \column{7.4in}{
     \item{\pad{{\bf Question:} \@question}}
     \item{\pad{{\bf Claim:} \@claim}}
+    \item{\smallskip{\bf Data:}\par\noindent\makedata}
     \item{\row{3.6015in}{2}{
-      \item{ {\bf Evidence:}\par \noindent\makedata\par\smallskip \@evidence \smallskip}
+      \item{ {\bf Evidence:}\par \@evidence \smallskip}
       \item{ {\bf Justification (Reasoning) of the Evidence:} \par\@justification \smallskip }
     }}
   }
diff --git a/cer/sem2/enthalpy/cer.log b/cer/sem2/enthalpy/cer.log
new file mode 100644
index 0000000..5fa8d3a
--- /dev/null
+++ b/cer/sem2/enthalpy/cer.log
@@ -0,0 +1,17 @@
+This is pdfTeX, Version 3.14159265-2.6-1.40.20 (TeX Live 2019) (preloaded format=pdftex 2019.5.8)  3 JUL 2019 23:03
+entering extended mode
+ restricted \write18 enabled.
+ %&-line parsing enabled.
+**../cer.tex
+(../cer.tex)
+*\see
+! Undefined control sequence.
+<*> \see
+        
+? x
+No pages of output.
+PDF statistics:
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diff --git a/cer/sem2/enthalpy/enthalpy.log b/cer/sem2/enthalpy/enthalpy.log
new file mode 100644
index 0000000..64810f7
--- /dev/null
+++ b/cer/sem2/enthalpy/enthalpy.log
@@ -0,0 +1,24 @@
+This is pdfTeX, Version 3.14159265-2.6-1.40.20 (TeX Live 2019) (preloaded format=pdftex 2019.5.8)  4 JUL 2019 00:01
+entering extended mode
+ restricted \write18 enabled.
+ %&-line parsing enabled.
+**enthalpy.tex
+(./enthalpy.tex (../cer.tex) [1{/usr/local/texlive/2019/texmf-var/fonts/map/pdf
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diff --git a/cer/sem2/enthalpy/enthalpy.pdf b/cer/sem2/enthalpy/enthalpy.pdf
new file mode 100644
index 0000000..205098b
--- /dev/null
+++ b/cer/sem2/enthalpy/enthalpy.pdf
diff --git a/cer/sem2/enthalpy/enthalpy.tex b/cer/sem2/enthalpy/enthalpy.tex
new file mode 100644
index 0000000..a98e9fa
--- /dev/null
+++ b/cer/sem2/enthalpy/enthalpy.tex
@@ -0,0 +1,61 @@
+\input ../cer.tex
+
+\name{Holden Rohrer}
+\course{FVS Chemistry AB 19.3}
+\teacher{Kerr} 
+
+\long\def\makedata{
+  \dimendef\tablewidth=1
+  \def\text##1{\tiny \hbadness=10000 \tolerance=10000 \hbox to \tablewidth{\hskip1em \vbox{\smallskip \noindent \advance \tablewidth by -2.25em \hsize\@tablewidth ##1 \smallskip} \hskip1em}}
+  \def\htext##1{\head \hbadness=10000 \tolerance=10000 \hbox to \tablewidth{\hskip0.5em \vbox{\smallskip \noindent \advance \tablewidth by -1.25em \hsize\@tablewidth ##1 \smallskip} \hskip0.5em}}
+  \def\thickline{\noalign{\hrule height 1pt}} 
+  \def\endliner{\noalign{\hrule height 0.05pt}}
+  \smallskip
+  \tablewidth=1in
+  \def\salign##1{\hbox{\quad \vbox{\smallskip \halign{##1} \smallskip}}}
+  \salign{\tablewidth 5in \vrule width 1pt \text{##} & \vrule width 0.05pt \text{##} & \vrule width 0.05pt \text{##} \vrule width 1pt \cr
+    \thickline
+    \omit \tablewidth 5in \vrule width 1pt \htext{Measurement} & \omit \vrule width 0.05pt \htext{Reaction 1 ($Mg + HCl$)} & \omit \vrule width 0.05pt \htext{Reaction 2 ($MgO + HCl$)} \vrule width 1pt \cr \endliner
+    Mass of HCl (g) & 100.50 & 100.57 \cr \endliner
+    Mass of solid (g) & 0.20 & 1.57 \cr \endliner
+    Total mass of reactants, m (g) & 100.70 & 102.14 \cr \thickline
+    Initial temperature ({\tinym $^{\circ}C$}) & 22.5 & 22.7 \cr \endliner
+    Temperature furthest from initial temperature ({\tinym $^{\circ}C$}) & 31.3 & 30.2 \cr \endliner
+    {\tinym $\Delta T$} ({\tinym $^{\circ}C$}) & 8.8 & 7.5 \cr \thickline
+    Heat released, {\tinym $q=cm\Delta T$} (J) & 3700 & 3200 \cr \endliner
+    Moles of solid reactant (mol) & 0.00823 & 0.0397 \cr \endliner
+    Enthalpy of reaction, {\tinym $\Delta H = {{-q}\over{mol}}*{{kJ}\over{1000J}}$} (kJ/mol) & -450 & -81 \cr \thickline
+  }
+  Combination of $\Delta H_1$, $\Delta H_2$, and $\Delta H_3$ needed (using Hess's Law): $\Delta H_{rxn} = -655{{kJ}\over{mol}}$
+
+  \smallskip
+  \tablewidth 3.5in
+  \salign{\vrule width 1pt \text{##} & \vrule width 0.05pt \text{##} \vrule width 1pt \cr \thickline
+    \omit \vrule width 1pt \htext{Reaction Equation} & \omit \vrule width 0.05pt \htext{$\Delta H$ (kJ/mol)} \vrule width 1pt \cr \endliner
+    Reaction 1: $Mg(s) + 2HCl(aq) \longrightarrow MgCl_2(aq)+H_2(g)$ & -450 \cr \endliner
+    Reaction 2: $MgO(s) + 2HCl(aq) \longrightarrow MgCl_2(aq)+H_2O(l)$ & -81 \cr \endliner
+    Reaction 3: $H_2(g) + {{1}\over{2}}O_2(g) \longrightarrow H_2O(l)$ & -286 (published value) \cr \thickline
+    Magnesium Combustion: $Mg(s)+{{1}\over{2}}O_2(g) \longrightarrow MgO(s)$ & -655 \cr \thickline
+  }
+  The published value for this reaction is -603 kJ/mol. Percent error of experimental value: 8.62\%
+  \smallskip
+}
+
+\question{How can you use Hess's Law to determine a reaction's enthalpy when you can't do so using a calorimeter?}
+\claim{Hess's Law accurately describes the enthalpy of an unknown reaction where other reactions' enthalpies (such that those reactions can be summed to the unknown reaction) are known.}
+
+\evidence{
+\def\enthunit{{{kJ}\over{mol}}}
+This experiment measured two chemical reactions, but three were used based off of commonly known values. The first of these is $Mg(s)+2HCl(aq) \longrightarrow MgCl_2(aq) + H_2(g)$. The enthalpy of this reaction was measured and recorded accurately because HCl was in excess, so the only limiting factor was the metal ($Mg$), of which the weight and molar mass was known. $\Delta T$ was also determined, and because a coffee cup calorimeter was used, that was able to be converted into released energy ($q$). The same process was completed for the dissolution of $MgO$ in $HCl$ by the reaction $MgO + 2HCl(aq) \longrightarrow MgCl_2(aq) + H_2O(l)$. Their respective enthalpies of reaction were $-450\enthunit$ and $-81\enthunit$. The third previously known reaction was the combination of gaseous diatomic hydrogen and oxygen into water. This reaction is impossible to conduct with our tools because it requires a bomb calorimeter to contain the gaseous components. Its enthalpy of reaction was $-286\enthunit$.
+
+From these, the reversal of the second reaction provided $H_2O(l)+MgCl_2(aq) \longrightarrow MgO(s)+2HCl(aq)$ at an enthalpy of $81\enthunit$. This allowed the ``cancelling'' of hydrochloric acid, magnesium chloride, hydrogen, and water when summing the chemical equations. After summation, we are left with $Mg(s) + {{1}\over{2}}O_2(g) \longrightarrow MgO(s)$ at a total enthalpy of $-450-(-81)-286=-450+81-286=-655\enthunit$ by Hess's Law. Because this is so similar to the previously published value of $-603 \enthunit$ (an experimental error of $8.62\%$, well within bounds of statistical significance), it can be considered to be an accurate measure of the enthalpy of the reaction (specifically magnesium combustion; however, this generalises).
+}
+
+\justification{
+Hess's Law's accuracy makes sense because it's logically consistent with enthalpy's property of being a state function. Starting with the assumed 0 enthalpy of elements in their standard form, most bonds release energy, and by conservation of energy, they must have proportionately lower potential energy. This is the energy of formation and is constant between all samples of a substance. In order for energy to be conserved during non-formation or -decomposition chemical reactions, the potential energy difference between products and reactants (enthalpy) must be equal to energy released. This is the foundation of Hess's Law: the known enthalpy amounts allow given compounds' enthalpies of formation to be treated like variables (because enthalpy of formation is an intrinsic property determined only by the identity of the substance and is identical between all samples of the substance) and thus the chemical equations summed like linear equations. This demonstrates the generality of the specific magnesium combustion example shown here.
+}
+
+\makeheader
+\vfil
+\makedoc
+\bye
diff --git a/cer/sem2/enthalpy/rate.tex b/cer/sem2/enthalpy/rate.tex
new file mode 100644
index 0000000..00e010d
--- /dev/null
+++ b/cer/sem2/enthalpy/rate.tex
@@ -0,0 +1,60 @@
+\input ../cer.tex
+
+\name{Holden Rohrer}
+\course{FVS Chemistry AB 19.3}
+\teacher{Kerr (sorry about the stretched out formatting; the writing just wouldn't fit)}
+
+\question{What are the effects of temperature and a reactant's particle size on reaction rate?}
+\claim{If you increase the temperature of a reaction, then the reaction rate will increase because particles experience more collisions at higher temperatures.
+
+{\bf Claim 2:} If you decrease the particle size of a reactant, then the reaction rate will increase because more of the reactants' surface area is exposed allowing more particles to make contact with eachother.}
+
+\long\def\makedata{
+  \dimendef\tablewidth=1
+  \def\text##1{\tiny \hbadness=10000 \tolerance=10000 \hbox to \tablewidth{\hskip1em \vbox{\smallskip \noindent \advance \tablewidth by -2.25em \hsize\@tablewidth ##1 \smallskip} \hskip1em}}
+  \def\htext##1{\head \hbadness=10000 \tolerance=10000 \vrule width 0.05pt \hbox to \tablewidth{\hskip0.5em \vbox{\smallskip \noindent \advance \tablewidth by -1.25em \hsize\@tablewidth ##1 \smallskip} \hskip0.5em}}
+  \def\style{\vrule width 1pt \htext{####} & \vrule width 0.05pt \text{####} & \vrule width 0.05pt \text{####} & \vrule width 0.05pt \text{####} & \vrule width 0.05pt \text{####} & \vrule width 0.05pt \text{####} \vrule width 1pt \crcr
+    \noalign{\hrule height 1pt}
+    \omit \vrule width 1pt \text{Trials} & \omit \htext{\fsub} & \omit \htext{Mass of Tablet ($mg$)} & \omit \htext{Volu\-me of Water ($L$)} & \omit \htext{Reac\-tion Time ($s$)} & \omit \htext{Reac\-tion Rate (${{mg}\over{L}}\over{s}$)} \vrule width 1pt \cr \endliner
+  }
+  \def\endliner{\noalign{\hrule height 0.05pt}}
+  \qquad Table A: Variation of Temperature
+  \smallskip
+  \tablewidth=0.55in
+  \def\fsub{Meas\-ured Reaction Temperature ($^{\circ}C$)}
+  \halign{\span\style
+    $\approx 20^{\circ}C$ & $24^{\circ}C$ & 1,000 & 0.2 & 34.2 & 146 \cr \endliner
+    $\approx 40^{\circ}C$ & $40^{\circ}C$ & 1,000 & 0.2 & 26.3 & 190 \cr \endliner
+    $\approx 65^{\circ}C$ & $65^{\circ}C$ & 1,000 & 0.2 & 14.2 & 352 \cr \endliner
+    $\approx 5^{\circ}C$ & $3^{\circ}C$ & 1,000 & 0.2 & 138.5 & 36 \cr
+    \noalign{\hrule height 1pt}
+  }
+  \bigskip
+  \qquad Table B: Variation of Particle Size (All at Room Temperature)
+  \smallskip
+  \def\fsub{Relat\-ive Particle Size (Small , Med\-ium, Large)}
+  \halign{\span\style
+    Full Tablet & Large & 1000 & 0.2 & 34.5 & 145 \cr \endliner
+    Eight Pieces & Medium & 1000 & 0.2 & 28.9 & 173 \cr \endliner
+    Tiny Pieces & Small & 1000 & 0.2 & 23.1 & 216 \cr
+    \noalign{\hrule height 1pt}
+  }
+  
+}
+
+\evidence{
+  The first claim is proven true by the evidence because the reaction rates do, in fact, increase with increasing temperatures. This is seen in the similar ordering of trials across both features. The temperatures ($T_k$ is the temperature of the $k$th trial) are, in order: $T_4 < T_1 < T_2 < T_3$. The reaction rates are, by a similar convention, $R_4 < R_1 < R_2 < R_3$. This clearly shows correlation of the two variables. In each of the trials, potentially confounding variables were controlled for---the mass of the tablet, volume of water in the beaker, and particle size (none of the tablets were crushed).
+  
+  The second claim follows a similar line of reasoning. Table B follows the changes in tablet size by crushing them into larger numbers of pieces (consequently, each piece has a smaller volume and disproportionately larger surface area and total surface area by the square-cube law). The reaction rate is highest for tiny pieces, slightly lower for eight pieces, and lowest for full, uncrushed tablets (as calculated from the increasing reaction times of the trials in the same order). This part of the experiment also controlled for similar variables: temperature was always room temperature as no heating or cooling occurred, and the masses/volumes of reactants were unchanged between trials. 
+}
+
+\justification{
+  Both of these results make sense within the context of collision theory. The temperature increases both the number of collisions because it is the same as kinetic energy and thus average velocities of the particles (faster particles cause more collisions) and it increases the number of effective collisions. The number of effective collisions increase because as the velocity of the particles go up, it becomes far more likely that the kinetic energy of a given collision is above the activation energy of the reaction.
+  
+  The second result can be justified similarly. The more pieces into which the tablet is crushed, the more of the ``internal'' surface area is exposed to the water. The more surface area is exposed to the water, the more chances there are for the water molecules to collide with the antacid's molecules, and the more reactions occur.
+}
+
+\makeheader
+\vfil
+\makedoc
+\bye
diff --git a/cer/sem2/solubility/solubility.tex b/cer/sem2/solubility/solubility.tex
new file mode 100644
index 0000000..a98e9fa
--- /dev/null
+++ b/cer/sem2/solubility/solubility.tex
@@ -0,0 +1,61 @@
+\input ../cer.tex
+
+\name{Holden Rohrer}
+\course{FVS Chemistry AB 19.3}
+\teacher{Kerr} 
+
+\long\def\makedata{
+  \dimendef\tablewidth=1
+  \def\text##1{\tiny \hbadness=10000 \tolerance=10000 \hbox to \tablewidth{\hskip1em \vbox{\smallskip \noindent \advance \tablewidth by -2.25em \hsize\@tablewidth ##1 \smallskip} \hskip1em}}
+  \def\htext##1{\head \hbadness=10000 \tolerance=10000 \hbox to \tablewidth{\hskip0.5em \vbox{\smallskip \noindent \advance \tablewidth by -1.25em \hsize\@tablewidth ##1 \smallskip} \hskip0.5em}}
+  \def\thickline{\noalign{\hrule height 1pt}} 
+  \def\endliner{\noalign{\hrule height 0.05pt}}
+  \smallskip
+  \tablewidth=1in
+  \def\salign##1{\hbox{\quad \vbox{\smallskip \halign{##1} \smallskip}}}
+  \salign{\tablewidth 5in \vrule width 1pt \text{##} & \vrule width 0.05pt \text{##} & \vrule width 0.05pt \text{##} \vrule width 1pt \cr
+    \thickline
+    \omit \tablewidth 5in \vrule width 1pt \htext{Measurement} & \omit \vrule width 0.05pt \htext{Reaction 1 ($Mg + HCl$)} & \omit \vrule width 0.05pt \htext{Reaction 2 ($MgO + HCl$)} \vrule width 1pt \cr \endliner
+    Mass of HCl (g) & 100.50 & 100.57 \cr \endliner
+    Mass of solid (g) & 0.20 & 1.57 \cr \endliner
+    Total mass of reactants, m (g) & 100.70 & 102.14 \cr \thickline
+    Initial temperature ({\tinym $^{\circ}C$}) & 22.5 & 22.7 \cr \endliner
+    Temperature furthest from initial temperature ({\tinym $^{\circ}C$}) & 31.3 & 30.2 \cr \endliner
+    {\tinym $\Delta T$} ({\tinym $^{\circ}C$}) & 8.8 & 7.5 \cr \thickline
+    Heat released, {\tinym $q=cm\Delta T$} (J) & 3700 & 3200 \cr \endliner
+    Moles of solid reactant (mol) & 0.00823 & 0.0397 \cr \endliner
+    Enthalpy of reaction, {\tinym $\Delta H = {{-q}\over{mol}}*{{kJ}\over{1000J}}$} (kJ/mol) & -450 & -81 \cr \thickline
+  }
+  Combination of $\Delta H_1$, $\Delta H_2$, and $\Delta H_3$ needed (using Hess's Law): $\Delta H_{rxn} = -655{{kJ}\over{mol}}$
+
+  \smallskip
+  \tablewidth 3.5in
+  \salign{\vrule width 1pt \text{##} & \vrule width 0.05pt \text{##} \vrule width 1pt \cr \thickline
+    \omit \vrule width 1pt \htext{Reaction Equation} & \omit \vrule width 0.05pt \htext{$\Delta H$ (kJ/mol)} \vrule width 1pt \cr \endliner
+    Reaction 1: $Mg(s) + 2HCl(aq) \longrightarrow MgCl_2(aq)+H_2(g)$ & -450 \cr \endliner
+    Reaction 2: $MgO(s) + 2HCl(aq) \longrightarrow MgCl_2(aq)+H_2O(l)$ & -81 \cr \endliner
+    Reaction 3: $H_2(g) + {{1}\over{2}}O_2(g) \longrightarrow H_2O(l)$ & -286 (published value) \cr \thickline
+    Magnesium Combustion: $Mg(s)+{{1}\over{2}}O_2(g) \longrightarrow MgO(s)$ & -655 \cr \thickline
+  }
+  The published value for this reaction is -603 kJ/mol. Percent error of experimental value: 8.62\%
+  \smallskip
+}
+
+\question{How can you use Hess's Law to determine a reaction's enthalpy when you can't do so using a calorimeter?}
+\claim{Hess's Law accurately describes the enthalpy of an unknown reaction where other reactions' enthalpies (such that those reactions can be summed to the unknown reaction) are known.}
+
+\evidence{
+\def\enthunit{{{kJ}\over{mol}}}
+This experiment measured two chemical reactions, but three were used based off of commonly known values. The first of these is $Mg(s)+2HCl(aq) \longrightarrow MgCl_2(aq) + H_2(g)$. The enthalpy of this reaction was measured and recorded accurately because HCl was in excess, so the only limiting factor was the metal ($Mg$), of which the weight and molar mass was known. $\Delta T$ was also determined, and because a coffee cup calorimeter was used, that was able to be converted into released energy ($q$). The same process was completed for the dissolution of $MgO$ in $HCl$ by the reaction $MgO + 2HCl(aq) \longrightarrow MgCl_2(aq) + H_2O(l)$. Their respective enthalpies of reaction were $-450\enthunit$ and $-81\enthunit$. The third previously known reaction was the combination of gaseous diatomic hydrogen and oxygen into water. This reaction is impossible to conduct with our tools because it requires a bomb calorimeter to contain the gaseous components. Its enthalpy of reaction was $-286\enthunit$.
+
+From these, the reversal of the second reaction provided $H_2O(l)+MgCl_2(aq) \longrightarrow MgO(s)+2HCl(aq)$ at an enthalpy of $81\enthunit$. This allowed the ``cancelling'' of hydrochloric acid, magnesium chloride, hydrogen, and water when summing the chemical equations. After summation, we are left with $Mg(s) + {{1}\over{2}}O_2(g) \longrightarrow MgO(s)$ at a total enthalpy of $-450-(-81)-286=-450+81-286=-655\enthunit$ by Hess's Law. Because this is so similar to the previously published value of $-603 \enthunit$ (an experimental error of $8.62\%$, well within bounds of statistical significance), it can be considered to be an accurate measure of the enthalpy of the reaction (specifically magnesium combustion; however, this generalises).
+}
+
+\justification{
+Hess's Law's accuracy makes sense because it's logically consistent with enthalpy's property of being a state function. Starting with the assumed 0 enthalpy of elements in their standard form, most bonds release energy, and by conservation of energy, they must have proportionately lower potential energy. This is the energy of formation and is constant between all samples of a substance. In order for energy to be conserved during non-formation or -decomposition chemical reactions, the potential energy difference between products and reactants (enthalpy) must be equal to energy released. This is the foundation of Hess's Law: the known enthalpy amounts allow given compounds' enthalpies of formation to be treated like variables (because enthalpy of formation is an intrinsic property determined only by the identity of the substance and is identical between all samples of the substance) and thus the chemical equations summed like linear equations. This demonstrates the generality of the specific magnesium combustion example shown here.
+}
+
+\makeheader
+\vfil
+\makedoc
+\bye