Matter and Measurement: Particles, Properties, and Units

What “matter” means in chemistry

In chemistry, matter is anything that has mass and takes up space (has volume). This sounds simple, but it becomes powerful when you connect it to measurement: if something is matter, you can usually measure how much of it you have (mass, volume, amount), how tightly packed it is (density), and how it behaves under different conditions (temperature, pressure).

Two important ideas sit underneath this definition:

• Mass is related to how much “stuff” is present. In everyday life we often say “weight,” but in science mass is the more fundamental quantity.
• Volume describes how much space the matter occupies.

Not everything you encounter is matter. For example, light and sound carry energy but do not have mass and volume in the same way ordinary substances do. Chemistry focuses on substances that are matter and on the changes they undergo.

Particles: the microscopic picture behind measurements

Chemistry connects the visible world to an invisible particle level. Matter is made of tiny particles. Depending on the situation, “particles” can mean atoms, molecules, or ions. You do not need to see particles directly to reason about them; you infer their behavior from measurements you can make.

Why the particle model matters for beginners

Many properties of matter make more sense when you imagine particles:

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• Gases expand to fill a container because their particles are far apart and moving freely.
• Liquids flow because their particles are close together but can slide past one another.
• Solids keep a fixed shape because their particles are packed and held in place.

When you measure mass, you are indirectly measuring how many particles (and what kind) are present. When you measure volume, you are measuring how much space those particles collectively occupy under the current conditions.

Physical properties vs. chemical properties

A property is a characteristic you can observe or measure. Chemistry separates properties into two categories because they answer different questions.

Physical properties (no new substance formed)

Physical properties can be observed without changing the identity of the substance. Examples include color, melting point, boiling point, density, electrical conductivity, and solubility.

• Density helps identify a substance and predict whether it will float or sink.
• Melting point and boiling point help you choose safe heating/cooling conditions in experiments.
• Solubility tells you how much of a substance can dissolve in a solvent at a given temperature.

Chemical properties (new substance formed)

Chemical properties describe how a substance can change into a different substance. Examples include flammability, tendency to rust, and reactivity with acids.

Measurement still plays a role: you might measure how fast a reaction happens, how much heat is released, or how much gas is produced. But the key difference is that chemical properties involve a change in composition.

Extensive properties vs. intensive properties

Another useful classification is based on whether a property depends on the amount of material.

Extensive properties (depend on amount)

Extensive properties change when you change the sample size. Examples: mass, volume, total energy, and length.

If you cut a piece of clay in half, each half has about half the mass and half the volume. That is typical extensive behavior.

Intensive properties (do not depend on amount)

Intensive properties are the same no matter how much you have (as long as it is the same substance under the same conditions). Examples: density, temperature, color, melting point, and boiling point.

If you pour a large bottle of pure water into two smaller bottles, the density of the water stays the same. That is why intensive properties are often used to identify substances.

Measurement: why units and tools matter

Chemistry is a measurement-based science. A statement like “I added some salt” is not very useful in a lab. “I added 2.50 g of salt” is useful because it is specific, repeatable, and can be checked by someone else.

Good measurement has three parts:

• A quantity (what you are measuring, such as mass).
• A number (how much).
• A unit (the scale used, such as grams).

Without units, numbers are ambiguous. “10” could mean 10 g, 10 mL, 10 kg, or 10 °C—very different situations.

The SI system: the common language of units

Most chemistry uses the International System of Units (SI). SI provides base units and derived units. You do not need to memorize every unit at once, but you should be comfortable with the most common ones used in introductory chemistry.

Common SI base units used in chemistry

• Mass: kilogram (kg) is the SI base unit, but grams (g) are commonly used in labs.
• Length: meter (m).
• Time: second (s).
• Temperature: kelvin (K) is the SI unit; degrees Celsius (°C) is also widely used.
• Amount of substance: mole (mol).

Common derived units in chemistry

• Volume: cubic meter (m³) in SI, but liters (L) and milliliters (mL) are common in labs.
• Density: kg/m³ in SI; g/mL or g/cm³ are common for liquids and solids.
• Pressure: pascal (Pa), often kilopascal (kPa); atmospheres (atm) and millimeters of mercury (mmHg) are also used.

Metric prefixes: scaling units without changing the idea

Metric prefixes let you express very large or very small quantities conveniently. Instead of writing many zeros, you adjust the unit with a prefix.

• kilo- (k) = 10³ = 1000
• centi- (c) = 10^-2 = 0.01
• milli- (m) = 10^-3 = 0.001
• micro- (µ) = 10^-6
• nano- (n) = 10^-9

Being fluent with prefixes is essential because lab equipment is often labeled in mL, g, mg, or µL, and chemical data tables may use kPa or MPa.

Step-by-step: converting with metric prefixes

Method A: use powers of ten.

• Example: convert 250 mL to liters.
• Know that 1 mL = 10^-3 L.
• Compute: 250 mL × (10^-3 L / 1 mL) = 0.250 L.

Method B: use a simple ratio.

• Example: convert 3.6 kg to grams.
• Know that 1 kg = 1000 g.
• Compute: 3.6 kg × (1000 g / 1 kg) = 3600 g.

In both methods, the unit you do not want cancels, leaving the unit you do want. This “unit cancellation” habit is one of the most reliable ways to avoid mistakes.

Temperature scales used in chemistry

Temperature affects volume, solubility, reaction rates, and many other properties. Chemistry commonly uses Celsius and Kelvin.

Celsius (°C) and Kelvin (K)

• Celsius is convenient for everyday conditions (water freezes at 0 °C and boils at 100 °C at standard pressure).
• Kelvin is the SI unit and is used in many formulas because it starts at absolute zero.

Step-by-step: converting °C to K

• Use: K = °C + 273.15
• Example: 25.0 °C → K = 25.0 + 273.15 = 298.15 K

Notice that Kelvin is written without a degree symbol. Also, in many lab contexts you keep a reasonable number of decimal places based on the measurement.

Volume in the lab: liters, milliliters, and reading glassware

Volume is commonly measured with graduated cylinders, pipettes, burets, and volumetric flasks. Each tool is designed for a different level of accuracy.

Choosing the right tool

• Beaker: good for mixing and rough volumes; not for precise measurement.
• Graduated cylinder: moderate accuracy; good for measuring liquids in many general procedures.
• Volumetric flask: high accuracy for preparing a solution to a specific volume.
• Pipette/buret: high accuracy for delivering measured volumes, especially in titrations.

Step-by-step: reading the meniscus correctly

Many liquids form a curved surface in glassware called a meniscus. Water typically forms a concave meniscus.

• Place the cylinder on a level surface.
• Bring your eye level with the liquid surface (not above or below).
• For a concave meniscus, read the volume at the lowest point of the curve.
• Record the value with the correct number of digits based on the smallest scale marking.

This prevents parallax error (a reading error caused by viewing from an angle).

Mass measurement: balances and good technique

Mass is measured using a balance. In beginner labs you often use a digital top-loading balance, and in more advanced work an analytical balance.

Step-by-step: measuring mass of a solid properly

• Check that the balance is clean and reads 0.000 g (or the appropriate display) before starting.
• Place a container (weigh boat, paper, or beaker) on the balance.
• Press tare/zero to set the container mass to zero.
• Add the solid to the container until you reach the desired mass.
• Record the mass with all digits shown by the balance (those digits reflect the instrument’s precision).

Taring avoids having to subtract the container mass later and reduces arithmetic errors.

Density: linking mass and volume to identify substances

Density is a key intensive property defined as mass per unit volume:

density = mass / volume

Common units: g/mL for liquids, g/cm³ for solids, and kg/m³ in SI. Because 1 mL = 1 cm³, g/mL and g/cm³ are numerically equivalent.

Practical example: density of a liquid

Suppose you measure 10.00 mL of a liquid and find its mass is 7.90 g (after taring the container).

density = 7.90 g / 10.00 mL = 0.790 g/mL

This value can be compared to reference densities to help identify the liquid or check purity.

Step-by-step: density of an irregular solid using water displacement

For a solid that does not have an easy geometric shape (like a small metal bolt), you can find volume by displacement.

• Partially fill a graduated cylinder with water and record the initial volume (V₁).
• Carefully add the solid so it is fully submerged and record the new volume (V₂).
• The solid’s volume is V_solid = V₂ − V₁.
• Measure the mass of the solid on a balance (m).
• Compute density = m / V_solid.

Example numbers: V₁ = 25.0 mL, V₂ = 31.5 mL, so V_solid = 6.5 mL. If m = 52.0 g, then density = 52.0 g / 6.5 mL = 8.0 g/mL (reported with appropriate significant figures based on the measurements).

Significant figures: reporting what you actually know

Every measurement has uncertainty. Significant figures are a way to communicate the precision of a measured value. The idea is: record all certain digits plus one estimated digit.

Common rules you will use

• Nonzero digits are significant (23.7 has 3 significant figures).
• Zeros between nonzero digits are significant (1002 has 4 significant figures).
• Leading zeros are not significant (0.0045 has 2 significant figures).
• Trailing zeros are significant only if a decimal point is shown (2.50 has 3 significant figures; 2500 without a decimal is ambiguous).

Significant figures in calculations (practical rules)

• Multiplication/division: the result has the same number of significant figures as the factor with the fewest significant figures.
• Addition/subtraction: the result is limited by the least precise decimal place.

Example (multiplication/division): 7.90 g / 10.00 mL = 0.790 g/mL. The limiting value is 7.90 (3 sig figs), so the density is reported with 3 sig figs: 0.790 g/mL.

Example (addition): 12.11 mL + 0.3 mL = 12.41 mL, but you report 12.4 mL because 0.3 mL is only precise to the tenths place.

Accuracy vs. precision: two different measurement goals

Accuracy describes how close a measurement is to the true or accepted value. Precision describes how close repeated measurements are to each other.

• You can be precise but not accurate (measurements cluster tightly but are all offset due to a calibration error).
• You can be accurate but not precise (average is close to true value, but measurements vary widely).

In lab work, you aim for both by using appropriate tools, careful technique, and repeated trials when needed.

Dimensional analysis: a step-by-step method for unit conversions

Dimensional analysis (also called the factor-label method) is a systematic way to convert units and check that equations make sense. The core idea is to multiply by conversion factors that equal 1, so you change units without changing the quantity.

Step-by-step: converting volume units

Convert 2.75 L to mL.

• Use the relationship: 1 L = 1000 mL.
• Set up the conversion so liters cancel:

2.75 L × (1000 mL / 1 L) = 2750 mL

Because 2.75 has 3 significant figures, you would typically report 2750 mL as 2.75 × 10³ mL if you want to make the significant figures explicit.

Step-by-step: converting density units (g/mL to kg/m³)

Convert 0.790 g/mL to kg/m³.

• Use: 1 g = 10^-3 kg and 1 mL = 10^-6 m³.

0.790 g/mL × (10^-3 kg / 1 g) × (1 mL / 10^-6 m^3) = 0.790 × 10^3 kg/m^3 = 790 kg/m^3

This conversion is common because many reference tables and physics-based formulas use kg/m³.

Common lab pitfalls and how to avoid them

Mixing up mass and volume

Mass (g) and volume (mL) are different quantities. You cannot convert between them without additional information such as density. If someone asks for “10 mL of ethanol,” you need a volume tool. If they ask for “10 g of ethanol,” you need a balance.

Forgetting temperature dependence

Some properties depend strongly on temperature, especially volume and density. A liquid’s volume expands when warmed, so density often decreases with temperature. When you compare your measured density to a reference value, check that the temperatures match or are close.

Using the wrong number of digits

Recording too few digits throws away information; recording too many suggests false precision. Let the instrument determine the digits you record. For example, if a graduated cylinder has markings every 1 mL, you typically estimate one more digit (to 0.1 mL). A digital balance shows the digits you should record.

Practice set: apply particles, properties, and units together

Scenario 1: identifying an unknown metal sample

You are given a small metal cylinder and told it is either aluminum or steel. You measure mass and volume to compute density.

• Measure mass: m = 21.6 g.
• Measure volume by displacement: V₁ = 15.0 mL, V₂ = 23.0 mL, so V = 8.0 mL.
• Compute density: 21.6 g / 8.0 mL = 2.7 g/mL.

A density near 2.7 g/mL is consistent with aluminum. The reasoning connects a macroscopic measurement (mass and volume) to an intensive property (density) that reflects particle-level packing and atomic mass.

Scenario 2: checking whether a liquid is likely water

You measure 50.0 mL of a clear liquid at room temperature and find its mass is 49.8 g.

density = 49.8 g / 50.0 mL = 0.996 g/mL

This is close to water’s density near room temperature (about 1.00 g/mL, depending on temperature). You would still consider other evidence (smell, boiling point, conductivity), but density provides a strong first check.