Understanding Matter and Its Composition
All material things in the universe are forms of matter. Matter is defined as anything that has both mass and volume, meaning it occupies space. From the air we breathe to the cells in our bodies, everything around and inside us is composed of matter. Scientists categorize matter in two primary ways: by physical state (solid, liquid, gas) and by chemical composition (pure substances or mixtures).
For AP Biology and general chemistry understanding, the composition of matter—what it's made of—is especially important. The chemical composition of a substance dictates its properties, how it interacts with other substances, and how it behaves under different physical conditions.
Three States of Matter
Before diving into chemical composition, it's important to understand how matter exists physically in nature. The three most common physical states are solids, liquids, and gases.
Solids
Solids have a fixed shape and volume.
The atoms or molecules in a solid are tightly packed together in a rigid structure.
These particles can only vibrate slightly in place, and do not move freely.
This dense packing and low particle movement account for the rigidity and incompressibility of solids.
Liquids
Liquids have a fixed volume but no fixed shape.
The particles in a liquid are still closely packed but not as rigidly arranged as in solids.
They can move past one another, allowing the liquid to flow and take the shape of its container.
Liquids are slightly compressible and exhibit intermolecular forces weaker than those in solids.
Gases
Gases have neither fixed shape nor volume.
Gas particles are widely spaced and in constant, random motion.
Gases expand to fill the entire volume of a closed container and are highly compressible.
The particles in a gas experience very weak or negligible intermolecular forces.
These differences in the physical behavior of matter help us understand how particles interact, which is essential when analyzing pure substances.
Classification of Matter by Composition
Matter is also categorized based on what it's made of. The broadest division in this classification is between pure substances and mixtures.
Pure substances are materials made from only one type of particle (atom or molecule) and have a uniform and definite composition.
Mixtures, on the other hand, consist of two or more different substances physically combined. These will be discussed in a separate subtopic.
Pure Substances
A pure substance is a type of matter that has a constant composition and consistent properties throughout. Every sample of a pure substance has the same chemical makeup, regardless of its origin or quantity.
There are two main categories of pure substances:
Elements
Compounds
Elements
An element is a pure substance consisting of only one type of atom. Each element has unique physical and chemical properties and is represented on the periodic table by a symbol (e.g., O for oxygen, H for hydrogen, Au for gold).
Elements cannot be broken down into simpler substances using chemical reactions.
Examples include helium (He), iron (Fe), and nitrogen (N₂).
Some elements exist in diatomic form (e.g., O₂, N₂) in nature, but they are still made up of only one type of atom.
Compounds
A compound is a pure substance composed of two or more elements chemically bonded in a fixed ratio. These elements lose their individual properties and create a new substance with unique characteristics.
For instance, water (H₂O) is made up of two hydrogen atoms and one oxygen atom, always in that exact proportion.
Another example is sodium chloride (NaCl), which consists of equal numbers of sodium and chloride ions.
In compounds, the elements are held together by chemical bonds, and the composition is always uniform and definite. This fixed ratio is a key part of what defines a compound’s elemental composition.
Formula Units
In the context of compounds, especially ionic compounds, the term formula unit is used to describe the simplest whole-number ratio of ions that represent the compound’s composition.
A formula unit does not represent a single molecule, as ionic compounds do not form discrete molecules.
For sodium chloride (NaCl), the formula unit reflects a 1:1 ratio of sodium ions to chloride ions.
Understanding formula units is important because it allows chemists to express the composition of a compound and use it for stoichiometric calculations.
This concept is particularly useful when analyzing ionic compounds, where atoms are arranged in repeating crystal lattice structures rather than individual molecules.
The Law of Definite Proportions
Also known as the Law of Constant Composition, the Law of Definite Proportions is a foundational chemical law that states:
A pure compound always contains the same elements in the same fixed proportion by mass, regardless of the source or sample size.
Example:
In water (H₂O), the mass ratio of hydrogen to oxygen is always approximately 1:8. This means that in any sample of water:
If you have 10 grams of water, it will contain about 1.1 grams of hydrogen and 8.9 grams of oxygen.
If you have 100 grams of water, it will contain about 11.1 grams of hydrogen and 88.9 grams of oxygen.
No matter how large or small the sample, this ratio will never change for a pure substance. If it does, then the substance is not pure or is not water.
Analogy: Recipes
Think of a recipe for making chocolate chip cookies:
If the recipe requires 2 cups of flour, 1 cup of sugar, and 1 cup of chocolate chips, that ratio must remain the same, regardless of how many batches you make.
If you alter the ratio, you're not making the same cookie.
Similarly, in chemistry, changing the ratio of elements in a compound creates a different compound. The fixed ratio defines the chemical identity of the compound.
Empirical Formula
The empirical formula of a compound gives the simplest whole-number ratio of atoms of each element present in the compound.
It is derived using the principles of the law of definite proportions and is extremely useful when determining the identity of unknown compounds.
Empirical vs Molecular Formula
The molecular formula shows the actual number of atoms of each element in a molecule.
Example: Glucose has a molecular formula of C6H12O6.
The empirical formula simplifies that to the smallest ratio of atoms:
Glucose becomes CH2O.
Image Courtesy of scienceabc
Although different molecules may have the same empirical formula, their molecular formulas and structures can be very different.
Determining Empirical Formula: Step-by-Step Process
To find an empirical formula, you are typically given the percent composition of a compound. Follow these steps:
Step 1: Convert Percent to Grams
Assume a 100-gram sample so that percent values can be treated as grams.
Example:
33.3% C becomes 33.3 g C,
7.4% H becomes 7.4 g H.
To find the remaining percentage for oxygen, subtract:
100% - 33.3% - 7.4% = 59.3% O,
so you have 59.3 g O.
Step 2: Convert Grams to Moles
Use the molar mass of each element:
Carbon = 12.01 g/mol
Hydrogen = 1.008 g/mol
Oxygen = 16.00 g/mol
Calculate moles:
C: 33.3 g / 12.01 g/mol = 2.773 mol
H: 7.4 g / 1.008 g/mol = 7.34 mol
O: 59.3 g / 16.00 g/mol = 3.706 mol
Step 3: Divide by the Smallest Number of Moles
This simplifies the ratio to the smallest whole numbers:
C: 2.773 / 2.773 = 1
H: 7.34 / 2.773 ≈ 2.65
O: 3.706 / 2.773 ≈ 1.34
Step 4: Multiply to Eliminate Decimals
To get whole numbers, multiply all by 3:
C = 1 × 3 = 3
H = 2.65 × 3 ≈ 8
O = 1.34 × 3 ≈ 4
Final empirical formula: C3H8O4
This process illustrates how experimental data (like percent composition) is used to determine the simplest elemental ratio in a compound.
Important Terms to Know
Element: A pure substance made of only one kind of atom.
Compound: A substance made of two or more elements chemically bonded in fixed ratios.
Pure Substance: Material with a definite composition and consistent properties.
Formula Unit: Simplest ratio of ions in an ionic compound.
Law of Definite Proportions: States that a pure compound always has the same mass ratio of elements.
Empirical Formula: Simplest whole-number ratio of elements in a compound.
Molecular Formula: Actual number of atoms of each element in a molecule.
States of Matter: Physical forms of matter—solid, liquid, gas.
Ionic Compound: Compound made of cations and anions held together by ionic bonds.
Molar Mass: The mass of one mole of a substance, expressed in grams per mole.
FAQ
Empirical formulas are essential because they provide the simplest ratio of elements in a compound, which is critical for identifying unknown substances and analyzing experimental data. In many laboratory settings, especially when analyzing new or unknown compounds, percent composition is often the first data available. From this, scientists can determine the empirical formula, which may help narrow down possible identities or assist in determining molecular formulas later when molar mass is known.
Empirical formulas are foundational for stoichiometric calculations.
They help identify patterns and relationships between different compounds.
They are especially useful in combustion analysis and elemental analysis techniques.
In biochemical contexts, empirical formulas allow simplified representation of complex molecules.
Yes, different compounds can share the same empirical formula but differ in their molecular formula and structure. These are called isomers or chemically distinct substances with the same element ratio.
For example, both glucose (C6H12O6) and acetic acid (C2H4O2) have the empirical formula CH2O.
They differ in molar mass, structure, and function.
To distinguish between them:
Determine the molar mass using methods like mass spectrometry.
Analyze chemical properties and functional groups through spectroscopy.
Study reactivity patterns and physical properties such as boiling point or solubility.
Empirical formulas alone are not sufficient to uniquely identify a compound, but they provide a valuable starting point.
Temperature and pressure do not affect the elemental composition of a pure substance. The ratio of elements in a pure substance is fixed and defined by its chemical bonds, not by external physical conditions. However, temperature and pressure can change the state of matter or the physical form of the substance.
Water (H2O) remains H2O whether it is in ice, liquid, or vapor form.
Changes in temperature may cause phase transitions, but the chemical identity remains unchanged.
Only a chemical reaction can change the elemental composition, such as electrolysis breaking H2O into hydrogen and oxygen.
This is a fundamental concept distinguishing physical changes from chemical changes.
Percent composition by mass refers to the proportion of each element's mass in a compound relative to the total mass of the compound. Elemental composition by atom count refers to the actual number of atoms of each element in a molecule or formula unit.
For H2O:
Atom count: 2 hydrogen atoms, 1 oxygen atom.
Mass percent: Hydrogen = ~11.1%, Oxygen = ~88.9%.
Key differences:
Percent composition helps with empirical formula calculations.
Atom count is used for writing molecular and structural formulas.
Percent mass varies by atomic weight, while atom count is a simple ratio.
Both are critical for understanding quantitative and qualitative aspects of chemical compounds.
Elemental composition plays a vital role in identifying and analyzing biological macromolecules like proteins, nucleic acids, and carbohydrates. Scientists use elemental analysis to determine the ratio of elements such as carbon, hydrogen, nitrogen, oxygen, phosphorus, and sulfur.
In proteins, the ratio of C:H:N:O:S provides clues about the amino acid content.
In DNA, the composition of C, H, N, O, and P helps identify nucleotide structure and sequence composition.
Elemental analysis is combined with techniques like mass spectrometry and chromatography to deduce molecular formulas and infer structures.
This information helps in:
Identifying unknown biomolecules
Comparing homologous sequences
Studying metabolic products
Practice Questions
A compound is found to be composed of 36.0% carbon, 6.0% hydrogen, and 58.0% oxygen by mass. Determine the empirical formula of the compound and explain the steps used to arrive at your answer.
To determine the empirical formula, first convert each percentage to grams, assuming a 100 g sample: 36.0 g C, 6.0 g H, and 58.0 g O. Then convert grams to moles using molar masses: C = 3.00 mol, H = 6.00 mol, O = 3.625 mol. Divide each mole value by the smallest (3.00): C = 1, H = 2, O ≈ 1.21. Multiply all ratios by 5 to get whole numbers: C = 5, H = 10, O = 6. The empirical formula is C5H10O6, showing the simplest whole-number ratio of elements in the compound.
Explain how the law of definite proportions supports the identification of pure substances and give an example involving a common compound.
The law of definite proportions states that a pure compound always contains the same elements in a fixed ratio by mass. This law supports the identification of pure substances because any deviation from the expected ratio suggests the substance is either a mixture or a different compound. For example, in water (H2O), the mass ratio of hydrogen to oxygen is always about 1:8. If a sample claiming to be water has a different ratio, it cannot be pure water. This predictable composition allows scientists to confirm the identity and purity of substances through quantitative analysis.
