Some Basic Concepts of Chemistry · Lesson 3 of 11

Properties, Measurement, and Units

“Describe physical quantities with meaningful units and connect mass, volume, density, and temperature.”

Learning Objectives

• Distinguish physical properties from chemical properties. • Identify the seven SI base quantities and recognise derived units and prefixes. • Explain mass, weight, volume, density, and common measuring apparatus. • Convert temperature scales and related volume units. • Explain why measurement standards and calibration are needed.

Properties become useful when we can compare them

A substance has characteristic properties that help us describe and identify it. Colour, melting point, and density describe observable behaviour, while combustibility and reactions with acids describe how it can form new substances. Careful measurements allow different observers to make meaningful comparisons.

Definition

Physical property

A property observed or measured without changing the chemical identity of the substance, such as density, colour, melting point, or boiling point.

Definition

Chemical property

A characteristic describing how a substance undergoes chemical change, such as combustibility or reactivity with acids or bases.

A quantitative observation needs both a number and a unit. Saying a room is 6 long is incomplete; 6 metres and 6 centimetres describe very different lengths. A unit is an agreed reference amount against which the measured quantity is compared.

A common measurement language

Different local units make comparisons difficult. The decimal metric system made multiples and subdivisions convenient, and the International System of Units, abbreviated SI, provides a shared system. Seven base quantities form the foundation; quantities such as volume, speed, and density are expressed using combinations of base units.

Base quantity	Quantity symbol	SI unit	Unit symbol
Length	$l$	metre	m
Mass	$m$	kilogram	kg
Time	$t$	second	s
Electric current	$I$	ampere	A
Thermodynamic temperature	$T$	kelvin	K
Amount of substance	$n$	mole	mol
Luminous intensity	$I_{v}$	candela	cd

The unit definitions are anchored to fixed values of physical constants so that a standard can be realised through experiments. The metre is tied to the speed of light; the kilogram to Planck’s constant; and the mole to the exact number of specified entities. You need not reproduce the experimental realisation to use the units correctly.

Unit	Defining constant or reference	Fixed numerical value used in SI
second	Caesium-133 transition frequency	$9192631770 Hz$
metre	Speed of light in vacuum	$299792458 m s^{- 1}$
kilogram	Planck constant	$6.62607015 \times 1 0^{- 34} J s$
ampere	Elementary charge	$1.602176634 \times 1 0^{- 19} C$
kelvin	Boltzmann constant	$1.380649 \times 1 0^{- 23} J K^{- 1}$
mole	Avogadro constant	$6.02214076 \times 1 0^{23} mo l^{- 1}$
candela	Luminous efficacy of specified monochromatic radiation	$683 lm W^{- 1}$ at $540 \times 1 0^{12} Hz$

A reference standard is used to calibrate an instrument: its readings are checked against a known reference. National metrology institutes maintain and compare measurement standards; in India this role is performed by the National Physical Laboratory in New Delhi. Older standards used physical platinum–iridium objects for the kilogram and metre. Those objects belong to the history of measurement; the current kilogram definition uses Planck’s constant.

Prefixes change the scale of a unit by a power of ten. For example, milli means one-thousandth, whereas kilo means one thousand. Letter case matters: m and M do not mean the same prefix. The following table contains the prefix range used in the source chapter.

Power of ten	Prefix	Symbol
$1 0^{- 24}$	yocto	y
$1 0^{- 21}$	zepto	z
$1 0^{- 18}$	atto	a
$1 0^{- 15}$	femto	f
$1 0^{- 12}$	pico	p
$1 0^{- 9}$	nano	n
$1 0^{- 6}$	micro	µ
$1 0^{- 3}$	milli	m
$1 0^{- 2}$	centi	c
$1 0^{- 1}$	deci	d
$1 0^{1}$	deca	da
$1 0^{2}$	hecto	h
$1 0^{3}$	kilo	k
$1 0^{6}$	mega	M
$1 0^{9}$	giga	G
$1 0^{12}$	tera	T
$1 0^{15}$	peta	P
$1 0^{18}$	exa	E
$1 0^{21}$	zetta	Z
$1 0^{24}$	yotta	Y

Mass, weight, and volume

Mass is a measure of the amount of matter in a sample in the sense used in this chapter. Weight is the gravitational force on that sample. Moving an object to a place with a different gravitational field changes its weight, while its mass remains the same. An analytical balance compares mass with high precision; laboratory amounts are often expressed in grams rather than kilograms.

Measuring mass against a reference— The source illustrates a mechanical analytical balance. Modern electronic balances serve the same mass-measurement role with different mechanisms.

Volume is the space occupied by a sample. Since volume involves three lengths, its SI unit is $m^{3}$ . Smaller laboratory volumes are commonly given in $c m^{3}$ , $d m^{3}$ , millilitres, or litres. Cubing a length conversion requires cubing the numerical factor as well.

Connecting common volume units— Each edge matters: 10 cm on three edges gives 1000 cubic centimetres.

Laboratory volume relationsLaTeX

1 L = 1 d m^{3} = 1000 c m^{3} = 1000 mL

Therefore

1 mL = 1 c m^{3}

and

1 L = 1 0^{- 3} m^{3}

. Litre is accepted for use with SI, but the coherent SI volume unit is the cubic metre.

Apparatus has different measuring roles— A volumetric flask sets the final solution volume; it is not simply a container for adding a stated volume of water.

A graduated cylinder measures a range of volumes. A burette measures the volume delivered by comparing readings, while a suitable pipette transfers a known volume. A volumetric flask is used to prepare a solution to its specified final volume. Read the liquid level at eye height; for a clear aqueous liquid the bottom of the curved meniscus is normally used.

Example — Cubing a unit conversion

Problem
Express $1 m^{3}$ in cubic centimetres.

1.One metre is $100 cm$ . A cubic metre has that length along all three perpendicular edges.
2. $1 m^{3} = (100 cm)^{3} = 1 0^{6} c m^{3}$ .
3.Using only a factor of 100 would convert a length, not a volume.

Density relates mass to volume

Equal volumes of different substances need not have equal masses. Density compares mass for each unit of volume, allowing us to compare samples of different sizes. For a uniform substance under the same conditions, doubling the sample size doubles both mass and volume, so their ratio stays the same.

DensityLaTeX

ρ = \frac{m}{V}

ρ

is density,

m

is mass, and

V

is volume. Common units are

kg m^{- 3}

and

g c m^{- 3}

Density reflects both particle masses and the space they occupy. It is not a direct measurement of particle spacing alone: a material can be denser because its particles are heavier, more closely arranged, or both. Density can also depend on temperature and physical state.

Example — Finding a sample volume

Problem
A uniform liquid has density $0.800 g m L^{- 1}$ . What volume corresponds to $20.0 g$ ?

1.Rearrange $ρ = m / V$ to $V = m / ρ$ because volume is the unknown.
2. $V = (20.0 g) / (0.800 g m L^{- 1}) = 25.0 mL$ .
3.The gram units cancel. A lower-density liquid needs more than $20.0 mL$ to have a mass of $20.0 g$ , which agrees with this answer.

Temperature scales use different zero points

Celsius, Fahrenheit, and kelvin describe temperature with different origins and, in the case of Fahrenheit, different interval sizes. Under standard atmospheric pressure, water freezes near 0 degrees Celsius and boils near 100 degrees Celsius. These correspond to 32 and 212 degrees Fahrenheit.

Celsius and FahrenheitLaTeX

T_{F} = \frac{9}{5} T_{C} + 32

T_{C}

is the numerical temperature in degrees Celsius;

T_{F}

is the numerical temperature in degrees Fahrenheit.

Celsius and kelvinLaTeX

T_{K} = T_{C} + 273.15

T_{K}

is the numerical temperature in kelvin. A kelvin has the same interval size as one degree Celsius; its zero point is different.

Comparing temperature scales— A difference of one Celsius degree equals one kelvin, but their zero points differ.

Example — A temperature on three scales

Problem
Convert $25.00^{\circ} C$ to kelvin and degrees Fahrenheit.

1. $T_{K} = 25.00 + 273.15 = 298.15$ , so the temperature is $298.15 K$ .
2. $T_{F} = (9/5) (25.00) + 32 = 77.00$ , so it is $77.00^{\circ} F$ .
3.The same physical temperature has three numerical descriptions. The conversion does not change the sample.

Unit symbols and temperature intervals

Write K, not degree K. A temperature rise of $10^{\circ} C$ is a rise of $10 K$ , but a temperature of $10^{\circ} C$ is $283.15 K$ . For ordinary systems discussed here, absolute zero is the lower limit of the kelvin scale.

Quick check

Which is a chemical property?

Quick check

Which statement about mass and weight is correct?

Quick check

What is $1 mL$ equal to?

Quick check

A $10.0 c m^{3}$ sample has mass $25.0 g$ . Its density is:

Quick check

Which apparatus is designed to set a known final solution volume?

Practice Problems

Distinguish boiling point from combustibility as types of property.
Write the SI units of mass, time, amount of substance, and temperature.
Convert $250 mL$ to litres and cubic metres.
Find the mass of $15.0 c m^{3}$ of a substance with density $2.00 g c m^{- 3}$ .
Convert $- 10.0^{\circ} C$ to kelvin and degrees Fahrenheit.
Explain why an instrument should be checked against a reference standard.
An object moves from Earth to the Moon without losing material. Explain which of mass and weight changes.

Key Takeaways

A quantitative measurement requires a number and a unit. SI has seven base quantities; derived units are combinations of base units. Mass and gravitational weight are different quantities. Volume conversions require powers of length factors, and density is mass per volume. Temperature conversions change the numerical scale, not the sample temperature.

Previous · Lesson 2

States and Classification of Matter

Next · Lesson 4

Scientific Notation and Reliable Calculations

Properties become useful when we can compare them

A common measurement language

Mass, weight, and volume

Density relates mass to volume

Temperature scales use different zero points

Quiz

Practice Problems

Key Takeaways