Matter and Its Properties

Our primary goal in chemistry is to understand the nature of matter and the changes it undergoes. To do this, we must first build a vocabulary for describing the “stuff” of the universe. This chapter lays the foundation for that vocabulary. We will learn how chemists classify matter based on its physical state and chemical composition, and we will explore the key properties, such as density and temperature, that we use to characterize it. Mastering this language of classification and measurement is the first step toward understanding the more complex chemical transformations to come.

Classifying Matter

ImportantA Critical Distinction: Mass vs. Weight

Before we begin, we must define matter itself. Matter is anything that has mass and occupies space. Mass is the fundamental property of matter that measures the amount of “stuff” an object contains. It is distinct from weight, which is the force of gravity on that matter. In chemistry, we are almost always concerned with an object’s mass, an intrinsic property that is the same everywhere in the universe.

The most fundamental way to describe matter is by its composition and its physical state. This two-tiered classification system allows us to create a hierarchical map for organizing and understanding all the substances in the universe.

The Three Common States of Matter

The first and most familiar way to classify matter is by its physical state. Matter typically exists in one of three states: solid, liquid, or gas. The state of a substance is a macroscopic property that is determined by the microscopic arrangement and motion of its constituent particles.

Solid

In a solid, particles are packed tightly together in a fixed arrangement. They vibrate in place but do not generally move past one another. Solids have a definite shape and a definite volume.

Liquid

In a liquid, particles are in close contact but are free to move past one another. A liquid has a definite volume but takes the shape of its container.

Gas

In a gas, particles are far apart and move randomly and rapidly. A gas has no definite shape or volume; it expands to fill its container completely. Gases are highly compressible.


These are macroscopic descriptions of behavior that arise directly from the particulate level. The addition or removal of energy can drive a substance to undergo phase transitions from one state of matter to another.



TipBeyond the Basics: Other States of Matter

While solid, liquid, and gas are the states we encounter most often on Earth, they aren’t the only ones!

  • Plasma: If you add enough energy to a gas, the atoms themselves can be torn apart into a “soup” of charged electrons and ions. This is plasma. It is the most abundant state of matter in the universe, making up stars, lightning, and the aurora borealis.

  • Supercritical Fluids: Under very high temperature and pressure, the distinction between liquid and gas can disappear. A substance in this state, called a supercritical fluid, has properties of both: it can diffuse through solids like a gas and dissolve substances like a liquid. This unique behavior is used commercially to decaffeinate coffee beans.


Composition: Pure Substances vs. Mixtures

Beyond its physical state, the most fundamental classification of matter is based on its chemical composition. Every substance can be classified as either a pure substance or a mixture. The distinction hinges on a simple question: “Can the substance be separated into simpler components by physical means?”

Pure Substances

A pure substance has a uniform and definite composition throughout. It cannot be separated into other substances by physical processes like filtration or evaporation. It cannot be separated into other kinds of matter by any physical process, such as filtration or evaporation. Pure substances are further divided into two categories.

  • An element is the simplest form of matter, consisting of only one type of atom. Elements cannot be broken down into simpler substances by any chemical reaction. Examples include iron (Fe), oxygen gas (O2), and helium (He).
  • A compound is a pure substance composed of two or more different elements chemically bonded together in a fixed ratio, as described by the Law of Definite Proportions. Compounds can be broken down into their constituent elements only by a chemical reaction. Examples include water (H2O) and sodium chloride (NaCl).

Mixtures

A mixture consists of two or more pure substances that are physically combined but not chemically bonded. The components of a mixture retain their individual chemical properties and can be separated from one another by physical means.

  • A homogeneous mixture has a uniform composition and appearance throughout. The individual components are mixed at the molecular level and are not visually distinguishable. Homogeneous mixtures are also called solutions. Examples include saltwater, clean air, and alloys like brass.
  • A heterogeneous mixture does not have a uniform composition. Its components are often visually distinguishable, and its properties can vary from one part of the mixture to another. Examples include sand and water, oil and vinegar, and a piece of granite.

Practice

A student is given four unlabeled beakers, each containing a different substance. Based on the following descriptions, classify each substance as an element, a compound, a homogeneous mixture, or a heterogeneous mixture. Provide a brief justification for each classification.

  1. Beaker A contains a clear, blue liquid. When a portion of the liquid is evaporated, a white crystalline solid remains.
  2. Beaker B contains a silvery, liquid metal. Heating the metal in a sealed container causes it to boil, but upon cooling, the original silvery liquid is reformed. No other substances are produced.
  3. Beaker C contains a cloudy white liquid. After being left undisturbed for an hour, a fine white powder settles at the bottom of the beaker.
  4. Beaker D contains a colorless, odorless gas that, when analyzed, is found to be composed of 85.6 % carbon and 14.4 % hydrogen by mass. The gas can be decomposed into two simpler substances.
Solution

  1. Homogeneous Mixture: The substance is a uniform (clear, blue) liquid, but it is composed of at least two different substances (the blue liquid and the white solid) that can be separated by a physical process (evaporation). This fits the definition of a homogeneous mixture, or solution.

  2. Element: The substance is a single liquid that cannot be broken down into simpler substances by physical changes (boiling). Since it is in its simplest form, it is an element. (This describes mercury).

  3. Heterogeneous Mixture: The substance is not uniform in composition, as evidenced by the solid settling out over time. Its components are visually distinguishable and can be separated by physical means like filtration or decantation.

  4. Compound: The substance consists of more than one element (carbon and hydrogen) chemically bonded in a fixed ratio. It can be decomposed into simpler substances only by chemical means. This fits the definition of a compound. (This describes an alkene like butene, C4H8).


Characterizing Matter: Properties

Once we have classified a substance, we can describe it in more detail using its properties. These properties are the characteristics that give each substance its unique identity. We can group these properties into several key categories.

Physical and Chemical Properties

The most fundamental way to characterize a property is to determine whether observing it changes the substance’s chemical identity.

  • A physical property is a characteristic that can be observed and measured without changing the substance into a new one.
  • A chemical property describes the ability of a substance to undergo a chemical change, transforming it into a new substance.
TipPhyiscal, Chemical, or both?

Some properties like solubility or color can be classified as both physical and chemical changes. Why is that? Read The Blurry Line Between Physical and Chemical Properties article to find out!

Intensive and Extensive Properties

Another way to classify properties is based on whether they depend on the amount of matter present.

  • An intensive property is a property of matter that is independent of the amount of the substance. It is a characteristic of the substance itself.
  • An extensive property is a property that depends on the amount of the substance present.

For example, the mass and volume of a gold bar are extensive properties; a larger bar has more mass and more volume. However, the density of the gold is an intensive property; it is the same for a small gold nugget as it is for a large gold bar.

Qualitative vs. Quantitative

  • Qualitative properties are descriptive and non-numerical. Examples: color (red), odor (sweet), state (liquid).
  • Quantitative properties are measured and expressed with a number. Examples: mass (10.5 g), temperature (25 °C).

A Closer Look at Key Quantitative Properties

Now that we have the full vocabulary for describing matter, let’s take a closer look at two of the most important quantitative, intensive, and physical properties we will use throughout this course: density and temperature.

Density

Density (ρ) is a fundamental property that relates the mass of a substance to the amount of space it occupies. It is defined as the ratio of an object’s mass (m) to its volume (V). \[ \rho = \frac{m}{V} \]

TipThe Symbol for Density

Density is commonly symbolized as d to give

\[ d = \dfrac{m}{V} \] The IUPAC recommended symbol is ρ which this textbook will adhere to.

Because density is an intensive property, it is a characteristic of the substance itself. A small gold nugget and a large gold bar have the same density. This makes density an invaluable tool for identifying unknown substances.

A common laboratory technique for determining the density of an irregularly shaped object is volume displacement. By measuring the change in the liquid level in a graduated cylinder after submerging an object, we can find the object’s volume, which can then be used with its mass to calculate its density.

Example: Density of an Unknown Metal

A student is tasked with identifying an unknown metal. They measure the mass of a piece of the metal to be 68.60 g. They then add the metal to a graduated cylinder containing water, which causes the water level to rise from an initial volume of 15.0 mL to a final volume of 22.5 mL.

What is the density of the unknown metal in g mL−1?

Step 1: Calculate the Volume of the Metal

First, we determine the volume of the metal by calculating the volume of water it displaces. The precision of the subtraction is limited by the tenths place of the volume readings. \[ \begin{align*} V(\mathrm{metal}) &= V_{\mathrm{final}} - V_{\mathrm{initial}} \\[1.5ex] &= 22.\bar{5}~\mathrm{mL} - 15.\bar{0}~\mathrm{mL} \\[1.5ex] &= 7.\bar{5}~\mathrm{mL} \end{align*} \] The resulting volume has 2 significant figures.

Step 2: Calculate the Density of the Metal

Next, we use the measured mass and the calculated volume to find the density. \[ \begin{align*} d(\mathrm{metal}) &= m(\mathrm{metal}) ~ V(\mathrm{metal})^{-1} \\[1.5ex] &= \left( 68.6\bar{0}~\mathrm{g} \right) \left( \frac{1}{7.\bar{5}~\mathrm{mL}} \right) \\[1.5ex] &= 9.\bar{1}46~\mathrm{g~mL^{-1}} \\[1.5ex] &= 9.1~\mathrm{g~mL^{-1}} \end{align*} \] The density of the metal is 9.1 g mL−1.


Practice

An analytical chemist is tasked with identifying an unknown liquid. She performs two experiments.

  1. In the first experiment, she measures out a 15.0 mL sample of the liquid and finds it has a mass of 11.9 g. In the second experiment, she uses a different balance and measures a 45.0 mL sample, which has a mass of 35.7 g. Based on this data, is density an intensive or extensive property? Justify your answer.

  2. A common solvent, dichloromethane (CH2Cl2), has a density of 1.33 g mL−1 at room temperature. What would be the mass, in kilograms, of a 2.50 L container of dichloromethane?

Solution

(a) Intensive vs. Extensive Property

First, we calculate the density from each experiment.

Experiment 1: \[ \rho = \frac{m}{V} = \frac{11.\bar{9}~\text{g}}{15.\bar{0}~\text{mL}} = 0.79\bar{3}33~\text{g mL}^{-1} = 0.793~\text{g mL}^{-1} \]

Experiment 2: \[ \rho = \frac{m}{V} = \frac{35.\bar{7}~\text{g}}{45.\bar{0}~\text{mL}} = 0.79\bar{3}33~\text{g mL}^{-1} = 0.793~\text{g mL}^{-1} \]

Density is an intensive property. Although the mass and volume (which are extensive properties) were changed between the experiments, the ratio of mass to volume (density) remained constant. This shows that the property is independent of the amount of substance present.

(b) Mass Calculation

We rearrange the density formula to solve for mass, m = ρ × V. We must first ensure our units are consistent. The density is given in g mL–1, but the volume is in L. We will convert liters to milliliters.

\[ V = \left ( \dfrac{2.5\bar{0}~\text{L}}{} \right ) \left ( \frac{1000~\text{mL}}{1~\text{L}} \right ) = 25\bar{0}0~\text{mL}\]

Now we can calculate the mass in grams.

\[ \begin{align*} \rho &= \dfrac{m}{V} \longrightarrow \\[1.5ex] m &= \rho V \\[1.5ex] &= \left( \frac{1.3\bar{3}~\text{g}}{\text{mL}} \right) \left ( \dfrac{25\bar{0}0~\text{mL}}{} \right ) \\[1.5ex] &= 33\bar{2}5~\text{g} \end{align*} \]

Finally, we convert the mass from grams to the requested unit of kilograms.

\[ m = \left ( \frac{33\bar{2}5~\text{g}}{} \right ) \left ( \frac{1~\text{kg}}{1000~\text{g}} \right ) = 3.3\bar{2}5~\text{kg} = 3.32~\text{kg} \]


Temperature

Temperature is another fundamental intensive property. While we intuitively understand it as a measure of “hotness” or “coldness,” its scientific definition is that temperature (T or t) is a measure of the average kinetic energy of the atoms or molecules in a system. This kinetic energy arises from the constant, random motion of particles—their translation, rotation, and vibration.

Temperature and heat are not the same thing. Heat (q) is the transfer of thermal energy between objects, while temperature is the property that determines the direction of that energy transfer.

TipTemperature is Not Heat

Imagine a tiny spark from a welder and a large bathtub of warm water.

  • The spark has an extremely high temperature (perhaps over 1000 °C). Its individual particles are moving incredibly fast, so their average kinetic energy is very high. However, because there are so few particles, the total amount of thermal energy is very small.
  • The bathtub has a low temperature (perhaps 40 °C). Its individual water molecules are moving much more slowly. However, because there is a vast number of molecules, the total amount of thermal energy it contains is immense.

Temperature determines the direction of spontaneous energy transfer. Energy will always be transferred as heat from the object with the higher temperature to the object with the lower temperature.

There are three common temperature scales used in science: Celsius (°C), Fahrenheit (°F), and Kelvin (K). The Kelvin (K) scale is the SI base unit and is an absolute temperature scale, meaning its zero point (0 K) is absolute zero, the theoretical temperature at which a system has reached its lowest possible energy state.

Temperature Conversion Formulas

The relationship between different temperature scales is not a simple ratio, so we cannot use a standard dimensional analysis conversion factor. Instead, the scales are related by linear transformation equations. A simple conversion factor (like 1 kg / 1000 g) only accounts for a change in scale, but these temperature equations must account for both a change in scale (the size of a degree is different) and a change in the zero point (0 °C is not the same as 0 °F).

Because these formulas include addition/subtraction to account for different zero points, we can’t use simple unit cancellation. Instead, we use these conversion equations directly, substituting the numerical temperature value and its unit..

The formal notation t/°C is read as “the numerical value of the temperature t when expressed in degrees Celsius.” This notation allows us to write the correct mathematical relationship between the numbers on each scale.

The relationships are therefore formally defined as equations relating these numerical values (exact values in bold): \[ t/^\circ\mathrm{C} = (t/^\circ\mathrm{F} - \mathbf{32}) / \mathbf{1.8} \] \[ t/^\circ\mathrm{F} = (\mathbf{1.8} \times t/^\circ\mathrm{C}) + \mathbf{32} \] \[ T/\mathrm{K} = t/^\circ\mathrm{C} + \mathbf{273.15} \]

NoteA Note on Temperature Symbols

To maintain clarity and adhere to scientific convention, we use specific symbols for temperature scales and temperature differences.

  • T (Capital Letter): This symbol is reserved for the absolute temperature scale, which is the Kelvin (K) scale.
  • t (Lowercase Letter): This symbol is used for all other common, non-absolute temperature scales, such as Celsius (°C) and Fahrenheit (°F).
  • ΔT (Temperature Change): A change in temperature has the same magnitude in both Kelvin and Celsius. A change of 1 K is identical to a change of 1 °C. Therefore, the symbol ΔT is used for a temperature difference, whether that difference is calculated from Kelvin or Celsius values.
TipUnderstanding Temperature Scales

Ever wonder why the temperature conversion equations are what they are? It has to do with how these temperature scales relate. Read More Than a Formula: Understanding Temperature Scales to find out more!


Practice

The boiling point of liquid nitrogen is −196 ;°C. What is this temperature in °F?

Solution

The relationship between Celsius and Fahrenheit is \(t/^\circ\text{C} = (t/^\circ\text{F} - 32)/1.8\).

\[ \begin{align*} t/^\circ\text{C} &= (t/^\circ\text{F} - 32)/1.8 \longrightarrow \\[1.5ex] t/^\circ\text{F} &= ( 1.8 \times t/^\circ\text{C} ) + 32 \\[1.5ex] &= \left ( 1.8 \times -19\bar{6} \right ) + 32 \\[1.5ex] &= -35\bar{2}.8 + 32 \\[1.5ex] &= -32\bar{0}.8~^\circ\text{F} \\[1.5ex] &= -321~^\circ\text{F} \end{align*} \]


Energy in Physical and Chemical Changes

We have now seen how to classify matter and describe its properties. The final concept that connects all of these ideas is energy. Energy is formally defined as the capacity to cause change. In thermodynamics, energy can perform work (causing physical motion) or transfer as heat (thermal energy transfer). More simply, energy is the driver of all change.

Every physical and chemical change, from the melting of an ice cube to the combustion of a fuel, is accompanied by a change in energy. To understand chemistry, we must understand the forms this energy can take at the molecular level.

  • Potential Energy (Ep): This is the stored energy that results from a particle’s position or composition. In chemistry, we are primarily concerned with chemical potential energy. This energy is stored within the chemical bonds that hold atoms together in a molecule and in the intermolecular forces that hold molecules near each other. When a chemical reaction occurs, old bonds are broken and new, more stable bonds are formed, often releasing this stored potential energy.

  • Kinetic Energy (Ek): This is the energy of motion. For the particles in a sample of matter, this includes translational (point-to-point), rotational, and vibrational motion. The temperature of a substance is a direct measure of the average kinetic energy of its constituent atoms or molecules.

The total energy of a system at rest, its internal energy (U), is the sum of the potential and kinetic energies of all the particles within it.

The Takeaway: The properties and changes we observe in matter are all manifestations of the energy stored within its bonds and the energy of its moving particles. Understanding this relationship between energy and matter is the key that unlocks the next topic in our study of chemistry: thermochemistry, the quantitative study of energy changes in chemical reactions.