Chemistry Study Guide for the HESI Exam

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Stoichiometry

Stoichiometry studies the quantitative relationships between reactants and products in a chemical reaction. It is based on the law of conservation of mass, which states that matter cannot be created or destroyed in a chemical reaction. This law allows us to predict how much of a reactant is needed or how much of a product will be formed using balanced chemical equations.

Consider the reaction of hydrogen gas and oxygen to form water:

\[2 \text{H}_2 + \text{O}_2 \rightarrow 2 \text{H}_2 \text{O}\]

This balanced equation tells us that two moles of hydrogen (\(\text{H}_2\)) react with one mole of oxygen (\(\text{O}_2\)) to produce two moles of water (\(\text{H}_2 \text{O}\)).

Dimensional analysis is used to convert between different units in chemistry, particularly in stoichiometry. The key to dimensional analysis is to set up conversion factors correctly so that units cancel out, leaving the desired unit. The molar mass of the substances is often needed to set up these factors.

For example, to determine how many grams of water can be produced from four moles of hydrogen gas, follow these steps:

  1. Use the balanced equation to find the mole ratio between water and hydrogen. Since the coefficient in the balanced chemical equation of water is \(2\), and the coefficient of hydrogen is \(2\), this means that the mole ratio between water and hydrogen is:
\[\frac{2 \text{ mol } \text{H}_2 \text{O}}{2 \text{ mol } \text{H}_2}\]
  1. Set up the conversion. To convert from moles of water to grams of water you need the molar mass of water, which is \(18\) grams per mole. This is found from the atomic weights of two hydrogen atoms and one oxygen atom:
\[4 \text{ mol } \text{H}_2 \times \frac{2 \text{ mol } \text{H}_2 \text{O}}{2 \text{ mol } \text{H}_2} \times \frac{18 \text{ g } \text{H}_2 \text{O}}{1 \text{ mol } \text{H}_2 \text{O}}\]
  1. Solve:

\(\require{cancel}\) \(4 \cancel{\text{ mol } \text{H}_2} \times \frac{2 \cancel{\text{ mol } \text{H}_2 \text{O}}}{2 \cancel{\text{ mol } \text{H}_2}} \times \frac{18 \text{ g } \text{H}_2 \text{O}}{1 \cancel{\text{ mol } \text{H}_2 \text{O}}} = 72.08 \text{ g } \text{H}_2 \text{O}\)

Thus, four moles of hydrogen gas produce \(72.08\) grams of water.

Oxidation and Reduction (Redox) Reactions

Oxidation is the loss of electrons. Reduction is a gain of electrons (which reduces the oxidation state). A great way to remember it is using the acronym OIL RIG (“Oxidation Is Losing, Reduction Is Gaining”). In redox reactions, oxidation and reduction occur simultaneously.

Oxidation Numbers and Rules

Oxidation numbers help track the movement of electrons in redox reactions. Here are the five key rules for assigning oxidation numbers:

  • Elements in their natural state have an oxidation number of \(0\).

  • Monoatomic ions have oxidation numbers equal to their charge. Group I metals always have oxidation numbers of +1. Group II metals always have oxidation numbers of +2. Fluorine always has an oxidation number of –1. Chlorine, bromine, and iodine usually have oxidation numbers of –1 except in compounds with oxygen or fluorine.

  • Oxygen usually has an oxidation number of \(–2\), except in peroxides (where it is \(–1\)).

  • Hydrogen is usually \(+1\), except when bonded to metals in hydrides (where it is \(–1\)).

  • The sum of oxidation numbers in a neutral compound is \(0\), and in a polyatomic ion it equals the ion’s charge. A polyatomic ion is made up of more than one atom.

Let’s try an example.

For the compound sulfuric acid, \(\text{H}_2\text{SO}_4\), what is the oxidation number of sulfur (\(\text{S}\))?

Solution

According to the rules we just discussed, hydrogen (\(\text{H}\)) has an oxidation number of \(+1\), oxygen (\(\text{O}\)) has an oxidation number of \(–2\) , and sulfur has an oxidation number that will make the sum of the oxidation numbers of all elements in the compound equal to \(0\) (a neutral molecule).

If we call the oxidation number of sulfur \(x\), this means:

\[2(\text{H})+x+4(\text{O})=0\]

So, plugging in the oxidation numbers of hydrogen and oxygen and solving for \(x\) gives an oxidation number of \(+6\) for sulfur.

Redox Equations

Redox (oxidation-reduction) equations represent chemical reactions where electrons are transferred between substances. These reactions are essential in biological processes like cellular respiration. A redox equation consists of two half-reactions:

  • oxidation half-reaction—shows the loss of electrons
  • reduction half-reaction—shows the gain of electrons

Consider this reaction as an example:

\[2\text{Na} + \text{Cl}_2 \rightarrow 2\text{Na}\text{Cl}\]

In this reaction sodium (\(\text{\bf{Na}}\)) is oxidized because it loses an electron to become \(\text{Na}^+\) and chlorine (\(\text{\bf{Cl}}\)) is reduced because it gains an electron to become \(\text{Cl}^–\). Sodium is the electron donor and chlorine is the electron acceptor.

  • The oxidation half-reaction shows the loss of one electron in sodium:
\[\text{Na} \rightarrow \text{Na}^+ + \text{e}^-\]
  • The reduction half-reaction shows the gain of one electron in chlorine:
\[\text{Cl}_2 + 2 \text{e}^- \rightarrow 2 \text{Cl}^-\]

Acids and Bases

An acid is a substance that donates protons (\(\text{H}^+\) ions), while a base (also called an alkaline compound or alkali) is a substance that accepts protons or produces hydroxide ions (\(\text{OH}^-\)) in solution.

The pH Scale

The acidity or basicity of a chemical or solution can be measured using the pH scale.

  • The pH scale ranges from \(\mathbf{0}\) to \(\mathbf{14}\), with a pH of \(\mathbf{7}\) being neutral.

  • A solution with a pH lower than \(\mathbf{7}\) is classed as acidic.

  • A solution with a pH higher than \(\mathbf{7}\) is classed as basic.

ph-scale_1760283025.png

Hydrochloric acid (\(\text{HCl}\)), which is part of the acid that is found in the stomach, is a very strong acid with a pH of \(1\). Water is neutral, with a pH of \(7\). Sodium bicarbonate (\(\text{NaHCO}_3\)) has a pH of \(9\), making it a weak alkali. Sodium hydroxide (\(\text{NaOH}\)) is a very strong alkali that, when concentrated, can have a pH of \(14\).

The pH scale is logarithmic with a base of \(\mathbf{10}\). This means that each unit difference corresponds to a change of a factor of \(10\). For example, pH \(3\) is \(10\) times more acidic than pH \(4\) and \(1\text{,}000\) times more acidic than pH \(6\).

Examples of Acids

Acids release \(\text{H}^+\) ions in water and have a pH less than \(\bf{7}\). These are a few examples:

  • Hydrochloric acid (\(\text{HCl}\)) is found in stomach acid.

  • Sulfuric acid (\(\text{H}_2 \text{SO}_4\)) is used in car batteries.

  • Acetic acid (\(\text{CH}_3 \text{COOH}\)) is found in vinegar.

Examples of Bases

Bases release \(\text{OH}^-\) ions in water and have a pH greater than \(\bf{7}\). These are a few examples:

  • Sodium hydroxide (\(\text{NaOH}\)) is used in soap making.

  • Ammonia (\(\text{NH}_3\)) is used in household cleaners.

  • Calcium hydroxide (\(\text{Ca}[\text{OH}]_2\)) is used in cement.

Litmus Paper Testing

Litmus paper is a simple and widely used pH indicator that helps determine whether a substance is acidic or basic. It is made by treating paper with natural dyes extracted from lichens, which change color depending on the pH of the solution.

There are two main types of litmus paper:

  • red litmus paper—Used to test for bases, it stays red in acidic solutions, but turns blue in basic (alkaline) solutions.

  • blue litmus paper—Used to test for acids, it stays blue in basic (alkaline) solutions, but turns red in acidic solutions.

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