Source: CHEM 106 Lab Report, New York University
Tags: density, volume, mass, graduated cylinder, g/mL, ethanol, water, measurement, CHEM 106
Difficulty: Introductory | Prerequisites: Basic arithmetic, familiarity with lab equipment (beakers, graduated cylinders, scales)
This lab sits right at the start of general chemistry and builds the foundational skill of measuring physical properties of matter. Density links two quantities you can measure in any lab (mass and volume) into a single value that identifies a substance. If you can measure density accurately, you can distinguish between substances, check the purity of a sample, and predict how liquids will layer or mix. You should already be comfortable reading a graduated cylinder and using a balance.
Density is mass divided by volume, measured in g/mL. You find it by weighing a known volume of liquid in a graduated cylinder, subtracting the cylinder's mass, then dividing. Water's accepted density is about 1.00 g/mL; ethanol's is roughly 0.789 g/mL. Lab error shows up as deviation from those values, and learning to spot (and explain) that deviation is the real point of the exercise.
Density
The mass of a substance per unit volume, typically expressed in grams per millilitre (g/mL) for liquids. Think of it as how tightly packed the matter is in a given space.
Volume
The amount of three-dimensional space a substance or object occupies. In this lab, measured in millilitres (mL) using graduated cylinders.
Mass
The quantity of matter in an object, measured in grams (g) on a balance. Unlike weight, mass does not change with gravity.
Graduated cylinder
A tall, narrow piece of lab glassware marked with volume graduations, used to measure liquid volumes precisely. In this experiment, 10 mL and 25 mL cylinders were used.
Pasteur pipet (transfer pipet)
A small glass or plastic dropper used to add or remove small amounts of liquid. Handy for adjusting a liquid to an exact volume line in a graduated cylinder.
Miscible
Describes two liquids that mix completely in all proportions to form a single uniform phase. Water and ethanol are miscible. In simple terms, if you pour one into the other, you get one layer, not two.
Tare (zeroing)
Resetting the balance to zero with the empty container on the pan so that subsequent readings give only the mass of the contents. This is how you isolate the mass of the liquid from the mass of the cylinder.
Density tells you how much mass is packed into a given volume. A denser substance has more mass per mL.
For any pure substance at a fixed temperature and pressure, density is constant. This makes it an identifying (intensive) property.
If you double the volume of a sample, you double its mass, but the density stays the same.
Weigh the empty graduated cylinder on the balance. Record the mass.
Add a measured volume of liquid (e.g. 8 mL or 10 mL) using a Pasteur pipet to hit the mark precisely.
Weigh the cylinder again with the liquid inside.
Mass of liquid = (mass of cylinder + liquid) minus (mass of empty cylinder).
Density = mass of liquid divided by volume of liquid.
Four trials were run at volumes of 8 mL and 10 mL.
Calculated densities ranged from 0.77 to 0.91 g/mL (one anomalous value of 0.10 g/mL, likely a recording error).
The accepted density of water at room temperature is approximately 1.00 g/mL, so all experimental values were lower than expected.
This systematic under-reading likely reflects measurement error: incomplete filling, parallax when reading the meniscus, or residual water on the outside of the cylinder.
Four trials at volumes of 8 mL and 10 mL.
Calculated densities ranged from 0.46 to 0.77 g/mL.
The accepted density of ethanol at room temperature is about 0.789 g/mL.
Variation between trials was wider than for water, possibly because ethanol evaporates more quickly, losing mass between weighing steps.
Record mass of empty cylinder: e.g. 24.32 g.
Add liquid to the desired volume mark: e.g. 10 mL.
Weigh cylinder + liquid: e.g. 32.12 g.
Subtract: 32.12 g minus 24.32 g = 7.80 g (mass of water).
Divide: 7.80 g / 10 mL = 0.78 g/mL.
\rho = \frac{m}{V}Where ρ (rho) = density in g/mL, m = mass in grams, V = volume in mL.
Rearranged forms:
m = \rho \times VV = \frac{m}{\rho}To find the mass of a liquid in a container:
m_{\text{liquid}} = m_{\text{cylinder + liquid}} - m_{\text{empty cylinder}}These four expressions cover every calculation in this lab. If you know any two of mass, volume, and density, you can find the third.
Density measurement is how quality-control labs verify that a batch of liquid (fuel, cooking oil, pharmaceutical solution) meets its specification. If the density is off, the composition is off.
Brewers and winemakers measure the density of their product at every stage of fermentation to track how much sugar has converted to alcohol. A hydrometer, which floats higher or lower depending on liquid density, is the standard tool.
Students often think a heavier object must be denser. It does not follow. A large block of wood is heavier than a small iron nail, but iron is far denser than wood. Density depends on mass per unit volume, not mass alone.
Students sometimes confuse mass with weight. In a general-chemistry lab at sea level the difference is negligible, but they are conceptually different: mass is the amount of matter, weight is the gravitational force on that matter.
A common error is reading the graduated cylinder at the top of the meniscus instead of the bottom. For water and most aqueous solutions, you read at the bottom of the curved surface. Parallax (looking from an angle rather than at eye level) makes this worse.
Some students assume density changes when you change the amount of sample. It does not. Density is an intensive property: it stays the same regardless of how much material you have.
⚠️ You will almost certainly be asked to calculate density from mass and volume data. Practise until the formula and the subtraction step are automatic.
⚠️ Know the accepted densities of water (~1.00 g/mL) and ethanol (~0.789 g/mL) at room temperature. Exam questions often give you an experimental value and ask you to calculate percent error.
⚠️ Be ready to explain why experimental density might differ from the accepted value. Common acceptable answers: parallax error, incomplete transfer of liquid, evaporation (especially ethanol), and balance calibration.
⚠️ Understand that density is an intensive property. This distinction between intensive (density, temperature, boiling point) and extensive (mass, volume, energy) properties is a classic exam question.
True or false: Doubling the volume of a water sample doubles its density. (False. Density is intensive; it stays the same.)
Fill in the blank: Density equals ______ divided by ______. (mass; volume)
True or false: If a liquid's density is less than 1.00 g/mL, it will float on water. (True, assuming the two are immiscible.)
Fill in the blank: The accepted density of ethanol at room temperature is approximately ______ g/mL. (0.789)
True or false: You should read a water meniscus at its highest point in a graduated cylinder. (False. Read at the bottom of the meniscus.)
Q: A graduated cylinder has a mass of 25.00 g when empty. After adding 10.0 mL of an unknown liquid, it has a mass of 33.50 g. What is the density of the liquid?
A: Mass of liquid = 33.50 g minus 25.00 g = 8.50 g. Density = 8.50 g / 10.0 mL = 0.85 g/mL.
Q: A student measures the density of water as 0.78 g/mL. The accepted value is 1.00 g/mL. What is the percent error?
A: Percent error = |0.78 minus 1.00| / 1.00 x 100 = 22%. This is quite large and suggests a systematic measurement error.
Q: Is density an intensive or extensive property? Explain why.
A: Intensive. It does not depend on the amount of substance present. A cup of water and a swimming pool of water have the same density at the same temperature.
Q: You have 15.0 mL of ethanol (density 0.789 g/mL). What is the mass of the ethanol?
A: m = ρ x V = 0.789 g/mL x 15.0 mL = 11.8 g.
Q: Name two sources of error in a density-by-graduated-cylinder experiment and explain how each would affect the result.
A: (1) Parallax error when reading the meniscus could cause the volume to be recorded too high or too low, directly affecting the calculated density. (2) Evaporation of a volatile liquid like ethanol reduces the mass between filling and weighing, giving a density that is lower than the true value.
This connects to the concept of significant figures and measurement uncertainty, which you will revisit in nearly every lab going forward. The precision of your graduated cylinder (typically ±0.1 mL) limits how many significant figures your density can have.
Density also ties directly into unit conversions and dimensional analysis: if you know the density of a substance, you can convert between mass and volume in a single step, a technique used constantly in stoichiometry.
In later topics on solutions and concentrations, you will use density to convert between molarity and mass percent, especially for concentrated acids and bases.
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