Comprehensive Chemistry Study Guide: Salts, Stoichiometry, ) Redox, and Organic Compounds

Salts, Ions, and Precipitation Reactions

Salts are chemical compounds consisting of positive and negative ions, held together by electrical forces rather than sharing electrons like molecules. Positive ions are known as cations, while negative ions are anions. Because opposite charges attract, a strong bond called an ionic bond is formed. Iones are created when atoms gain or lose electrons. For example, a sodium atom ($Na$) loses its outermost electron to become a sodium ion ($Na \rightarrow Na^+ + e^-$), while a chlorine atom gains one to become a chloride ion ($Cl + e^- \rightarrow Cl^-$). When these react, sodium chloride ($NaCl$) is formed.

Salts are organized in a regular pattern called a crystal lattice, where positive and negative ions alternate. This gives salts their characteristic structure; common table salt consists of millions of tiny crystals built from $Na^+$ and $Cl^-$ ions. Ions can be classified as simple or polyatomic. Simple ions consist of only one element, such as $Na^+$, $Mg^{2+}$, or $Cl^-$. Polyatomic (compound) ions consist of multiple atoms acting as a single unit, such as nitrate ($NO_3^-$), sulfate ($SO_4^{2-}$), and carbonate ($CO_3^{2-}$).

When a salt is introduced to water, it may be easily soluble (letopløselig) or poorly soluble (tungtopløselig). Solubility occurs because water is a polar molecule with positive and negative ends. The positive end of water attracts negative ions, and the negative end attracts positive ions, pulling them away from the crystal lattice. In precipitation reactions, different salt solutions are mixed. If the combination of ions forms a poorly soluble salt, they attract each other to form a solid ionic lattice, which appears as a precipitate (bundfald). For instance, mixing silver nitrate and sodium chloride results in a reaction where silver ions and chloride ions form silver chloride: $AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)$. The ionic steps are $Ag^+ + Cl^- \rightarrow AgCl(s)$. While $AgCl$ forms a white precipitate, the sodium and nitrate ions remain dissolved as $NaNO_3(aq)$.

Solubility can often be predicted: salts containing sodium ($Na^+$), potassium ($K^+$), or nitrates ($NO_3^-$) are almost always soluble. Conversely, silver chloride ($AgCl$) and barium sulfate ($BaSO_4$) are typical poorly soluble salts. The Periodic Table helps predict ion charges: Group 1 elements typically form $+1$ ions, Group 2 elements form $+2$ ions, and Group 17 elements form $-1$ ions. Potential errors in these experiments include contaminated test tubes, inconsistent solution amounts, or difficulty seeing very small precipitates.

Stoichiometry and the Decomposition of Sodium Hydrogen Carbonate

Stoichiometry involves calculating the amounts of substances in chemical reactions. The amount of substance is measured in moles ($n$), where one mole equals Avogadro's constant ($6.02 \times 10^{23}$) particles. Molar mass ($M$) is the mass of one mole of a substance ($g/mol$). For sodium hydrogen carbonate ($NaHCO_3$), also known as baking soda, the molar mass is calculated by adding the atomic masses: $23 + 1 + 12 + (16 \times 3) \approx 84\,g/mol$. This salt consists of $Na^+$ and $HCO_3^-$ ions and is used in baking because it releases carbon dioxide when heated.

In an experimental setting, heating $NaHCO_3$ leads to chemical decomposition. By weighing a crucible and lid ($21.470\,g$) and then adding the salt ($28.445\,g$), the mass of the reactant was determined to be $6.975\,g$. Upon heating, the substance decomposes into sodium carbonate, carbon dioxide, and water. The balanced equation is: $2NaHCO_3(s) \rightarrow Na_2CO_3(s) + CO_2(g) + H_2O(g)$. The mass of the solid remainder decreases because $CO_2$ and $H_2O$ escape as gases, though the total mass of the system is conserved.

At the particle level, heating increases the movement of ions until the bonds in the hydrogen carbonate ions become unstable and rearrange. To calculate the amount of moles in the experiment: $n = \frac{m}{M} = \frac{6.975\,g}{84\,g/mol} \approx 0.083\,mol$. According to the 2:1 ratio in the reaction schema, this produces $0.0415\,mol$ of $Na_2CO_3$. With a molar mass of $106\,g/mol$ for $Na_2CO_3$, the theoretical yield is approximately $4.4\,g$. Equivalent amounts of substance refer to the ratios defined by coefficients in a balanced equation. Potential errors include insufficient heating time, weighing the crucible while hot, or the substance absorbing moisture during cooling.

Gases and the Molar Mass of Magnesium

The experiment to determine the molar mass of magnesium utilizes the reaction between magnesium ribbon and hydrochloric acid ($HCl$): $Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_{2}(g)$. Specifically, the stoichiometry shows a 1:1 ratio between $Mg$ and $H_2$, meaning finding the moles of hydrogen gas reveals the moles of magnesium used. In the experiment, $0.309\,g$ of magnesium reacted with $65\,mL$ of $1\,M$ hydrochloric acid. The hydrogen gas was collected in a Hayduck apparatus to measure its volume, while temperature and atmospheric pressure were recorded.

The ideal gas law is used to find the substance amount ($n$) of the gas: $P \times V = n \times R \times T$. Pressure ($p$), volume ($V$), the gas constant ($R$), and temperature ($T$) in Kelvin must be known. In this trial, the amount of hydrogen produced was $0.01259\,mol$. Using the formula $M = \frac{m}{n}$, the experimental molar mass for magnesium was calculated as $24.54\,g/mol$. This is very close to the table value of $24.31\,g/mol$, representing an error of only $0.95\%$.

The molar volume of a gas is the volume one mole occupies at standard conditions ($22.4\,L$ at STP). In this lab, it was more accurate to use the ideal gas law to account for specific lab pressure and temperature. It was also determined that $HCl$ was in excess; while $0.065\,mol$ of $HCl$ was provided, only $0.0252\,mol$ was required, leaving $0.0398\,mol$ remaining. Potential errors include leaks in the apparatus, improper leveling of the water in the measuring tube, air trapped in the tubing, or the presence of a magnesium oxide ($MgO$) layer on the metal surface.

Molecular Structures, Polarity, and Alkanes

Chemical solubility and miscibility depend heavily on molecular structure and polarity. A fundamental rule is "like dissolves like"; polar substances dissolve in polar solvents (e.g., glucose and $NaCl$ in water), while non-polar substances dissolve in non-polar solvents (e.g., stearin in rensebenzin). Water is polar because of electronegativity—oxygen pulls electrons more strongly than hydrogen, creating an uneven charge distribution. Covalent bonds (electron pair bonds) occur when atoms share electrons, typically between non-metals.

Molecules are described using molecular formulas and structural formulas. For example, heptane has the molecular formula $C_7H_{16}$. Its structural formula, $CH_3–CH_2–CH_2–CH_2–CH_2–CH_2–CH_3$, shows it is a straight chain of carbon atoms. Alkanes are saturated hydrocarbons containing only single bonds. They are named based on the number of carbons: methane (1), ethane (2), propan (3), butan (4), pentan (5), hexan (6), and heptan (7). All alkanes end with the suffix "-an."

In contrast to covalent bonding, ionic bonding occurs between a metal and a non-metal where electrons are transferred ($Na \rightarrow Na^+ + e^-$; $Cl + e^- \rightarrow Cl^-$). The resulting ions attract to form $NaCl$. In experiments involving solubility, errors can arise from insufficient shaking, unclean test tubes, or difficulty seeing layers when densities are similar.

Metals and the Reactivity Series (Redox)

The reactivity series (spændingsrækken) ranks metals based on their tendency to lose electrons. Metals that lose electrons easily are "base" (uædle) metals (e.g., Magnesium), while those that do not are "noble" (ædle) metals (e.g., Copper). When a metal loses electrons, it undergoes oxidation ($Mg \rightarrow Mg^{2+} + 2e^-$). When an ion gains electrons, it undergoes reduction ($Cu^{2+} + 2e^- \rightarrow Cu$). These processes always occur simultaneously in a redox reaction.

A metal further to the left on the reactivity series can displace (skubbe ud) ions of a metal to its right. For example, magnesium can reduce copper ions because magnesium is more reactive: $Mg(s) + Cu^{2+}(aq) \rightarrow Mg^{2+}(aq) + Cu(s)$. Experimental results ranked reactivity as follows: Magnesium, Zinc, Iron, Aluminum, Hydrogen, and Copper. Note that Aluminum often appears less reactive than expected due to a protective oxide layer ($Al_2O_3$).

Metals also react with acids to produce hydrogen gas ($H_2$) if they are more reactive than hydrogen. For example: $Zn + 2HCl \rightarrow ZnCl_2 + H_2$. Here, zinc is oxidized to $Zn^{2+}$ while hydrogen ions are reduced to $H_2$ gas, visible as bubbles. Oxidation numbers are used to track electron movement; if the oxidation number increases, the substance is oxidized. Potential errors include contaminated pipettes, dirty metal surfaces, or slow reactions that are difficult to observe.

Experimental Redox and Titration Analysis

Redox reactions can be observed through color changes, such as the disappearance of the purple color of potassium permanganate ($KMnO_4$) when it is reduced. In experiments with tin(II) chloride and sodium sulfite, reactions occurred because the substances could be oxidized. However, sodium sulfate ($Na_2SO_4$) showed no reaction because the sulfur in sulfate already has its highest oxidation state ($+6$) and cannot be oxidized further. Redox balancing requires that both the number of atoms and the total charge are equal on both sides of the equation.

Quantitative analysis can be performed via titration, such as determining the acetic acid ($CH_3COOH$) content in vinegar. Acetic acid is a weak acid that releases hydrogen ions ($H^+$) in water. It is neutralized by a base like sodium hydroxide ($NaOH$) in a 1:1 ratio: $CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O$. During titration, a phenolphthalein indicator is used, which turns pink at the equivalence point—where the moles of acid equal the moles of base.

In a trial using $2.0\,mL$ of vinegar and $0.10\,M \, NaOH$, an average of $16.25\,mL$ of base was needed. The calculation performed was: $n(NaOH) = 0.10\,mol/L \times 0.01625\,L = 0.001625\,mol$. Thus, $n(CH_3COOH) = 0.001625\,mol$. The concentration was $c = \frac{0.001625\,mol}{0.0020\,L} = 0.8125\,mol/L$. Using the molar mass of acetic acid ($\approx 60.05\,g/mol$), the mass percentage was found to be $4.88\%$, matching the label's $5\%$ specification. The pH started at $3.76$ and rose to $8.44$ at the equivalence point. Error sources include misinterpreting the indicator color, incorrect burette readings, or uncleaned equipment.

Hydrocarbons and Combustion Reactions

Hydrocarbons consist solely of carbon and hydrogen. Alkanes follow the general formula $C_nH_{2n+2}$. Octane ($C_8H_{18}$), used as a model for gasoline, has a structural formula of $CH_3–(CH_2)_6–CH_3$. Combustion reactions occur when hydrocarbons react with oxygen. In complete combustion (plenty of oxygen), products are carbon dioxide and water: $C_5H_{12} + 8O_2 \rightarrow 5CO_2 + 6H_2O$. In incomplete combustion (limited oxygen), products can include carbon monoxide ($CO$) or pure carbon ($C$), which is seen as soot.

Smaller molecules like pentane have lower boiling points, evaporate more easily, and ignite more readily with cleaner flames. Larger molecules like dodecane ($C_{12}H_{26}$) have higher boiling points and produce more soot because it is harder for oxygen to reach all carbon atoms. Isomers are molecules with the same molecular formula but different structural arrangements, leading to different properties like boiling points. For example, pentane ($C_5H_{12}$) has three isomers: n-pentane, 2-methylbutane, and 2,2-dimethylpropane.

Carbon's ability to form four bonds and long chains stems from its four valence electrons. Potential experimental errors in combustion include varying liquid amounts, changing oxygen levels, or the crucible being pre-heated from previous trials, affecting evaporation and burning rates.