Chemistry Study Guide for the HESI Exam

General Information

Chemistry is more prevalent in nursing than you may initially assume. Nurses must understand the use of the medication they are administering, the conversions, and how the medications work. Furthermore, advancing your career into more specialized fields requires you to possess not only a basic understanding of chemistry, but perhaps organic chemistry and physical chemistry as well. Having a solid foundation in chemistry is key to understanding how certain drugs interact, which directly affects patient outcomes.

On this section of the HESI exam, you will find 25 questions and have 25 minutes to answer them. Here are some basic chemistry concepts you should understand to do well on the Chemistry section of the HESI exam:

Scientific Notation

Scientific notation is a way to express very small or very large numbers in a simpler format. It’s commonly used in chemistry to work with small values like the mass of a proton (\(1.67\times 10^{-27} \, \text{kg}\)) or large values like Avogadro’s constant (\(6.022\times 10^{23}\)).

Notation Parts

Scientific notation consists of three main parts: the mathematical sign, significand, and exponent. These are shown in the number below:

\[-2.3 \times 10^3\]

• mathematical sign: negative (\(-\))
• significand: \(2.3\)
• exponent: \(10^3\)

Mathematical Sign

The mathematical sign tells you whether the number is positive or negative. A positive number, indicated by the plus sign (\(+\)), is one in which the value is greater than \(0\), while a negative number, indicated by the minus sign (\(-\)), has a value less than \(0\). In the above example, the sign is negative.

Significand

The significand may be referred to as the coefficient and is always a number between \(\mathbf{1}\) and \(\mathbf{9}\). It represents the base value of the number in scientific notation. In the above example, the significand is \(2.3\). The significand always has only one place to the left of the decimal point.

Exponent

The exponent tells you how many places you need to move the decimal point to obtain the complete number. The exponent can either be positive or negative:

• With a positive exponent, you move the decimal point to the right (making a larger number). For example, for \(2.5 \times 10^4\), you move the decimal point four places to the right to get \(25\text{,}000\).

• With a negative exponent, you move the decimal point to the left (making a smaller number). For example, for \(7.1 \times 10^{-3}\), you move the decimal point three places to the left to get \(0.0071\).

Using Scientific Notation

One key skill you’ll need to master is converting from standard to scientific notation, and vice versa. These are the steps for converting from standard to scientific notation:

1. Identify where the decimal point currently is. If you are given the number \(5\text{,}600\text{,}000\), you should know this is the same as \(5\text{,}600\text{,}000.0\) (notice the decimal point at the end).

2. Move the decimal point so that only one nonzero digit is to the left. In this case, we will move it six places left and end up with \(5.6\).

3. Count how many places the decimal point moved and use that as the exponent. Since we moved six places left, the exponent is \(6\). Note: When we move the decimal point to the right, the exponent becomes negative.

4. Write the number in scientific notation:

\[5.6 \times 10^6\]

Converting from scientific to standard notation just reverses this process. Let’s try an example problem.

Convert \(3.2 \times 10^{-5}\) to the standard form.

Solution

To solve this problem, we simply reverse the process we just learned:

1. Identify the exponent. The exponent is \(–5\), which means we need to move the decimal point five places to the left. Note: We’re moving to the left with a negative exponent because this is the reverse operation of what we discussed above.

2. Determine the base value (significand), which in this case is \(3.2\).

3. Move the decimal point five places left. After moving it, we get the answer:

\[0.000032\]

The Metric System

Most measurements in nursing use the metric system, but you may need to approximate the US customary and/or imperial measurements as a comparison.

Metric Prefixes

Metric prefixes are unit prefixes that are placed before a basic unit of measurement and are used to represent different magnitudes of measurement. Understanding these prefixes is necessary for converting accurately in nursing calculations.

The following chart shows the most common metric prefixes:

Prefix	Symbol	Factor	Meaning
kilo-	k	\(10^{3}\)	\(1\text{,}000\) times larger
hecto-	h	\(10^{2}\)	\(100\) times larger
deca-	da	\(10^{1}\)	\(10\) times larger
base unit	-	\(10^{0}\)	standard unit
deci-	d	\(10^{-1}\)	\(10\) times smaller
centi-	c	\(10^{-2}\)	\(100\) times smaller
milli-	m	\(10^{-3}\)	\(1\text{,}000\) times smaller
micro-	\(\mu\)	\(10^{-6}\)	\(1\text{,}000\text{,}000\) times smaller
nano-	n	\(10^{-9}\)	\(1\text{,}000\text{,}000\text{,}000\) times smaller

Weight or Mass

• Mass (or weight) is usually measured in grams (g) or kilograms (kg), where \(1 \text{ kg} = 1\text{,}000 \text{ g}\).

• \[1 \text{ kg} = 2.2 \text{ lb}\]
• \[1 \text{ lb} = 16 \text{ oz}\]
• \[1 \text{ st} = 14 \text{ lb}\]

Note: You should be familiar with the basic abbreviations for common units like pounds (lb) and ounces (oz). The abbreviation “st” stands for stone, an imperial unit common in the UK.

Volume

Volume is usually measured in milliliters (mL) or cubic centimeters (\(\text{\bf{cm}}^3\)).

• \[1 \text{ mL} = 0.001 \text{ L}\]
• \[1 \text{ mL} = 1 \text{ cm}^3\]

For example:

\[5 \text{ L} = 5\text{,}000 \text{ mL} = 5\text{,}000 \text{ cm}^3\]

For larger quantities, US gallons (gal) are used in the United States, and imperial gallons (imp gal) are used in the United Kingdom. You should also be familiar with cups (c), fluid ounces (fl oz or oz), and US pints (pt). Some helpful conversions are:

• \[1 \text{ c} = 8 \text{ fl oz} = 240 \text{ mL}\]
• \[1 \text{ imp gal} \approx 1.2 \text{ gal}\]
• \[1 \text{ gal} = 231 \text{ in}^3 \approx 3.79 \text{ L}\]
• \[1 \text{ imp gal} \approx 4.55 \text{ L}\]
• \[1 \text{ imp gal} \approx 277.42 \text{ in}^3\]
• \[1 \text{ L} = 2.11 \text{ pt}\]

Distance

• Length or distance is commonly measured in meters (m), centimeters (cm), or millimeters (mm). You may need to be able to convert to these distances from US units like inches (in), feet (ft), yards (yd), and miles (mi).

• \[1 \text{ in} = 2.54 \text{ cm}\]
• \[1 \text{ ft} = 30.5 \text{ cm}\]
• \[1 \text{ yd} = 91.44 \text{ cm}\]
• \[1 \text{ km} = 0.621 \text{ mi}\]

Measuring Temperature

Temperature is a crucial measurement in science, particularly in healthcare. It is proportional to the average kinetic energy of a body and is used to assess body temperature, monitor fevers, and ensure the safe storage of medications. There are three main temperature scales: Fahrenheit, Celsius, and Kelvin.

Fahrenheit

Fahrenheit (\(\text{}^{\circ} \text{F}\)) is primarily used in the United States, mainly in thermometers and weather reports. It is based on a scale where water freezes at \(\mathbf{32^{\circ}} \textbf{F}\) and boils at \(\mathbf{212^{\circ}} \textbf{F}\). In this scale, the normal human body temperature is approximately \(\mathbf{98.6^{\circ}} \textbf{F}\).

Celsius

Celsius (\(\text{}^{\circ} \text{C}\)) is the standard temperature scale in most of the world and is widely used in science and medicine. It is based on the metric system, making conversions easier. In this scale, water freezes at \(\mathbf{0^{\circ}} \textbf{C}\) and boils at \(\mathbf{100^{\circ}} \textbf{C}\), while the normal human body temperature is approximately \(\mathbf{37^{\circ}} \textbf{C}\).

This is the formula to convert from Fahrenheit to Celsius:

\[\text{}^{\circ}\text{C} = (\text{}^{\circ}\text{F} - 32) \times \frac{5}{9}\]

To convert from Celsius to Fahrenheit, use this formula:

\[\text{}^{\circ}\text{F} = \left( \text{}^{\circ}\text{C} \times \frac{9}{5} \right) + 32\]

Kelvin

Kelvin (K) is an absolute scale used primarily in scientific research and physics. It is based on absolute zero, the point at which all molecular motion stops, which corresponds to \(\mathbf{0} \textbf{ K}\). On this scale, water freezes at \(\mathbf{273.15} \textbf{ K}\) and boils at \(\mathbf{373.15} \textbf{ K}\), while the normal human body temperature is approximately \(\mathbf{310.15} \textbf{ K}\).

To convert from Celsius to Kelvin, use this formula:

\[\text{K} = \text{}^{\circ}\text{C} + 273.15\]

Atoms

Everything around you, from the air you breathe to the water you drink, is made of matter. And at the core of all matter are atoms. Atoms are the smallest units of an element that retain the element’s properties. They are the building blocks of everything in the universe.

Atomic Structure

Elements are made of atoms, and atoms are made of three types of particles: negatively charged electrons, positively charged protons, and neutral neutrons.

The Nucleus

Protons and neutrons make up the nucleus of the atom. The nucleus has a very small diameter compared to the overall size of the atom, but it is where most of the mass is concentrated.

Orbits

The electrons orbit the nucleus in shells (also known as electron clouds), and most of the volume of the atom is taken up by the free space between the nucleus and electrons.

Electric Charge of Atoms

Atoms are usually neutral, meaning the number of protons equals the number of electrons. However, when an atom gains or loses electrons, it becomes an ion. Ions are important in biological processes such as nerve signaling and muscle contractions.

• Cations are positively charged ions that form when an atom loses electrons (\(\text{Na}^+\), \(\text{Ca}^{2+}\)).

• Anions are negatively charged ions that form when an atom gains electrons (\(\text{Cl}^-\), \(\text{O}^{2-}\)).

Note: The \(+\) sign indicates a positive charge and the \(-\) sign indicates a negative charge.

The Periodic Table

The periodic table organizes all known elements based on their atomic number and properties. It provides valuable insights into how elements interact with one another.

Retrieved from: https://commons.wikimedia.org/wiki/File:Simple_Periodic_Table_Chart-en.svg

There are many elements in the periodic table, but try to memorize the properties of these common ones:

Element	Symbol	Atomic Number	Atomic Weight (amu)
Hydrogen	H	1	1.01
Carbon	C	6	12.01
Nitrogen	N	7	14.01
Oxygen	O	8	16.00
Sodium	Na	11	22.99
Magnesium	Mg	12	24.31
Phosphorus	P	15	30.97
Sulfur	S	16	32.07
Chlorine	Cl	17	35.45
Calcium	Ca	20	40.08

Periods

The rows of the periodic table are called periods. Each period represents a new energy level (shell) being added to the atom’s structure. As you move down the table, elements increase in size because they have more electron shells.

Electron Configurations

Electron configurations show how electrons are arranged around the nucleus of an atom in energy levels (or shells) and sublevels (\(s, p, d, f\)). This arrangement determines an element’s chemical behavior.

Electrons fill orbitals in a specific order, following three main rules:

1. Aufbau principle: Electrons occupy the lowest energy orbitals available first, following the order shown in the image below:

Retrieved from: https://openstax.org/books/chemistry-2e/pages/6-4-electronic-structure-of-atoms-electron-configurations

1. Pauli exclusion principle: Each orbital holds a maximum of two electrons, and they must have opposite spins.

2. Hund’s rule: When electrons fill a set of orbitals of equal energy, one electron goes into each orbital first, all with the same spin, before pairing up.

For example, consider the electron configuration of nitrogen, which has an atomic number of 7, meaning that a neutral atom of nitrogen contains seven protons and seven electrons. These electrons would be arranged with the following electron configuration:

\[1s^22s^22p^3\]

Groups

The columns of the periodic table are called groups. Elements in the same group share chemical properties because they have the same number of valence (outermost) electrons. By looking at an element’s group, you can predict how it will behave in chemical reactions.

Electrical Charges

Elements in the same group tend to form ions with the same charge. For example:

• Group \(\bf{1}\) (alkali metals)—Elements in this group tend to lose one electron to form a \(+1\) charge.

• Group \(\bf{2}\) (alkaline earth metals)—Elements in this group tend to lose two electrons to form a \(+2\) charge.

• Group \(\bf{17}\) (halogens)—Elements in this group tend to gain one electron to form a \(–1\) charge.

Electrons in the Outer Shell

The electrons in the outer shell are called valence electrons. The number of valence electrons determines the reactivity of an element. Elements with very few valence electrons or close to a full shell are very reactive, while elements with full outer shells are nonreactive. To know which are the valence electrons, we write the electron configuration in increasing order of energy according to the Aufbau principle (shown in the figure below), and the valence electrons will be those with the greatest energy level.

Retrieved from: https://openstax.org/books/chemistry-2e/pages/6-4-electronic-structure-of-atoms-electron-configurations

Atomic Number

The atomic number is the number of protons in the nucleus of an atom. This number uniquely identifies each element.

For example:

• Hydrogen (\(\text{H}\)) has one proton, so its atomic number is \(1\).
• Carbon (\(\text{C}\)) has six protons, so its atomic number is \(6\).

Mass Number

The mass number is the total number of protons and neutrons in an atom’s nucleus. Electrons are incredibly light, so they do not significantly contribute to the overall mass of an atom. The mass number is always a whole number and specific to a particular isotope. For instance, the mass number of oxygen-\(16\) is \(16\) and the mass number of helium-\(4\) is \(4\).

To find the number of neutrons (\(n\)) in an atom, subtract the atomic number (\(a\)) from the mass number (\(m\)):

\[m - a = n\]

For example, Carbon-\(13\) has a mass number of \(13\) and an atomic number of \(6\). So, the number of neutrons it has is:

\[n = 13 - 6 = 7\]

Isotopes

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons.

For example, carbon-12 (\(^{12}_{6}\text{C}\)) and carbon-13 (\(^{13}_{6}\text{C}\)) have the same atomic number (\(6\)), meaning that they are the same element, but they have different mass numbers (\(12\) and \(13\)) due to the varying number of neutrons.

Atomic Weight

The atomic weight (or atomic mass) is the average mass of an element’s isotopes, usually measured in atomic mass units (amu or u). It represents the weighted average of the mass numbers of all the natural isotopes of that element. For instance, the atomic weight of oxygen is \(15.999 \text{ amu}\) and the atomic weight of helium is \(4.002602 \text{ amu}\).