Physical & Chemical Properties
Every substance possesses unique or characteristic properties that distinguish it from all others. NCERT classifies these into two categories. Physical properties include colour, odour, melting point, boiling point and density. Chemical properties include composition, combustibility and reactivity with acids and bases.
The defining test between them lies in what happens to the substance during measurement. A physical property can be measured or observed without changing the identity or composition of the substance — you can read a melting point or weigh out a density without converting the material into anything else. A chemical property, by contrast, can only be observed when a chemical change occurs: to know that magnesium is combustible, you must actually burn it.
Consider rusting. The grey lustre of iron is a physical property; its tendency to react with oxygen and moisture, $\ce{4Fe + 3O2 -> 2Fe2O3}$, is a chemical property, observable only because iron is consumed in the process. Chemists describe, interpret and predict the behaviour of substances using both classes of property, each determined by careful measurement and experimentation.
| Aspect | Physical property | Chemical property |
|---|---|---|
| Identity of substance | Unchanged on measurement | Changed — a new substance forms |
| Chemical change required? | No | Yes |
| Typical examples | Colour, odour, melting point, boiling point, density | Acidity/basicity, combustibility, reactivity with acids and bases |
| Reversibility | Often reversible (melting, dissolving) | Usually involves a new product |
Measurement & the SI System
Scientific investigation demands quantitative measurement. Many properties of matter — length, area, volume — are quantitative in nature, and any such observation is expressed as a number followed by a unit. The length of a room written as "6 m" carries two pieces of information: the number 6, and the unit m (metre) in which length is measured. A bare number is meaningless in chemistry.
Historically two systems competed — the English System and the Metric System. The metric system, which originated in France in the late eighteenth century, was the more convenient because it was built on the decimal system. As the scientific community grew, the need for a single common standard became pressing. That system, the International System of Units (SI, from the French Le Système International d'Unités), was established in 1960 by the 11th General Conference on Weights and Measures (CGPM).
The SI system rests on seven base units pertaining to the seven fundamental scientific quantities. All other physical quantities — speed, volume, density and the rest — are derived from these seven base quantities.
| Base physical 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 | Iv | candela | cd |
Figure 1 — The seven base units (teal) anchor the SI system; derived quantities such as volume, speed and density (purple) are combinations of them.
The modern definitions of these base units are fixed by universal physical constants rather than physical artefacts. The metre is defined through the fixed value of the speed of light in vacuum; the kilogram through the Planck constant $h = 6.62607015\times10^{-34}\ \text{J s}$; the second through the caesium-133 hyperfine transition frequency; and the kelvin through the Boltzmann constant. The mole is defined as exactly $6.02214076\times10^{23}$ elementary entities — the Avogadro number. In India, the National Physical Laboratory (NPL), New Delhi, maintains these national standards of measurement.
The system of units is not frozen. As the accuracy of measurement of a particular unit is enhanced by adopting new principles, the member nations of the Metre Convention (signed in Paris in 1875) agree to revise the formal definition of that unit. This is why the kilogram, once tied to a platinum-iridium cylinder, is now defined through the Planck constant. For NEET, the practical takeaway is simpler: you must recognise each base quantity, recall its unit and symbol, and know that quantities like volume and density are not fundamental but built up from these seven.
SI Prefixes
Chemistry routinely deals with quantities far larger or smaller than the base unit. The SI system therefore permits prefixes to denote decimal multiples and submultiples of a unit, so that a length of $10^{-9}\ \text{m}$ becomes a more readable 1 nanometre (nm). Mastery of the common prefixes converts cumbersome powers of ten into single, memorable symbols.
| Multiple | Prefix | Symbol | Multiple | Prefix | Symbol |
|---|---|---|---|---|---|
10⁻¹⁵ | femto | f | 10³ | kilo | k |
10⁻¹² | pico | p | 10⁶ | mega | M |
10⁻⁹ | nano | n | 10⁹ | giga | G |
10⁻⁶ | micro | µ | 10¹² | tera | T |
10⁻³ | milli | m | 10⁻² | centi | c |
10⁻¹ | deci | d | 10¹ | deca | da |
The lowercase/uppercase prefix clash
Case is not decorative in SI symbols. Lowercase m means milli ($10^{-3}$), but capital M means mega ($10^{6}$) — a factor of $10^{9}$ apart. Likewise k (kilo) is always lowercase, while temperature kelvin uses an uppercase K with no degree sign.
Read every symbol's case before substituting. "mm" is millimetre; "Mm" would be megametre.
Mass vs Weight
Mass and weight are among the most commonly confused pair of terms in physical chemistry, and NEET examiners exploit the slip. Mass is the amount of matter present in a substance; it is a constant for a given body, identical on Earth, on the Moon, or in deep space. Weight is the force exerted by gravity on that body, and so it can vary from place to place as the gravitational field changes.
The mass of a substance is determined accurately in the laboratory using an analytical balance. The SI unit of mass is the kilogram (kg), but because chemical reactions involve small amounts of material, the laboratory ordinarily quotes mass in the fractional unit gram, where $1\ \text{kg} = 1000\ \text{g}$.
"Weighing" a chemical does not measure its weight
When you place a sample on an analytical balance, you obtain its mass, not its weight, even though the everyday verb is "to weigh". A statement claiming that the mass of a reagent changes when carried to a higher altitude is false — only its weight would change.
Mass is invariant; weight depends on $g$. Stoichiometry always uses mass.
Once you can record a measurement with the right unit, the next skill is reporting it honestly. See Uncertainty & Significant Figures.
Volume & Its Units
Volume is the amount of space occupied by a substance. Because it has the dimensions of (length)3, its SI unit is the cubic metre ($\text{m}^3$). The cubic metre is inconveniently large for laboratory chemistry, so volumes are more often expressed in cubic centimetres ($\text{cm}^3$) or cubic decimetres ($\text{dm}^3$).
A widely used but non-SI unit for the volume of liquids is the litre (L). The litre relates to SI units through the chain $$1\ \text{L} = 1000\ \text{mL}, \qquad 1\ \text{dm}^3 = 1000\ \text{cm}^3 = 1\ \text{L}.$$ In the laboratory, liquid and solution volumes are measured with graduated cylinders, burettes and pipettes, while a volumetric flask is used to prepare a known volume of a solution.
| Unit | Symbol | SI? | Equivalence |
|---|---|---|---|
| Cubic metre | m³ | SI (derived) | 1 m³ = 1000 dm³ = 1000 L |
| Cubic decimetre | dm³ | SI (derived) | 1 dm³ = 1 L = 1000 cm³ |
| Litre | L | Non-SI (common) | 1 L = 1000 mL = 1 dm³ |
| Cubic centimetre | cm³ | SI (derived) | 1 cm³ = 1 mL |
Density
The two properties already discussed — mass and volume — combine to give a third, derived property. Density is the mass of a substance per unit volume: $$\text{Density} = \frac{\text{Mass}}{\text{Volume}}.$$ Working out its SI unit directly from this definition gives mass in kilograms over volume in cubic metres, that is $\text{kg m}^{-3}$. Because that unit is large for everyday materials, the chemist commonly expresses density in gram per cubic centimetre, $\text{g cm}^{-3}$, with mass in grams and volume in cm³.
Density encodes how closely the particles of a substance are packed: a higher density means more tightly packed particles. It is also a physical property, since it can be measured without altering the chemical identity of the material.
A rectangular block of a metal has a mass of $\ce{52.0\ g}$ and occupies a volume of $\ce{6.5\ cm^3}$. Find its density in $\text{g cm}^{-3}$ and convert it to SI units.
Apply the definition $\rho = \dfrac{m}{V}$:
$$\rho = \frac{52.0\ \text{g}}{6.5\ \text{cm}^3} = 8.0\ \text{g cm}^{-3}.$$
To express in SI units, use $1\ \text{g} = 10^{-3}\ \text{kg}$ and $1\ \text{cm}^3 = 10^{-6}\ \text{m}^3$:
$$8.0\ \frac{\text{g}}{\text{cm}^3} \times \frac{10^{-3}\ \text{kg}}{1\ \text{g}} \times \frac{1\ \text{cm}^3}{10^{-6}\ \text{m}^3} = 8.0 \times 10^{3}\ \text{kg m}^{-3}.$$
Answer: the density is $8.0\ \text{g cm}^{-3}$, equivalently $8.0\times10^{3}\ \text{kg m}^{-3}$ — a value consistent with a dense metal whose particles are closely packed.
Temperature Scales
Three scales are in common use for temperature: degree Celsius (°C), degree Fahrenheit (°F) and kelvin (K). Of these, the kelvin is the SI unit. Thermometers calibrated on the Celsius scale run from 0° to 100°, marking the freezing and boiling points of water, while the Fahrenheit scale places these same fixed points at 32° and 212°.
The Celsius and Fahrenheit scales are related by $$^{\circ}\text{F} = \frac{9}{5}\left(^{\circ}\text{C}\right) + 32,$$ and the kelvin scale is related to the Celsius scale by the simple offset $$K = {^{\circ}\text{C}} + 273.15.$$ A notable consequence is that while negative temperatures are perfectly valid on the Celsius scale, the kelvin scale admits no negative values — 0 K is the absolute lowest temperature attainable.
Figure 2 — The freezing and boiling points of water aligned across the three scales. Note that 100 Celsius degrees span the same interval as 100 kelvins but 180 Fahrenheit degrees.
A reaction mixture is held at $\ce{37\ ^\circ C}$ (human body temperature). Express this in kelvin and in degree Fahrenheit.
Celsius to kelvin — add 273.15:
$$K = 37 + 273.15 = 310.15\ \text{K}.$$
Celsius to Fahrenheit — apply $^{\circ}\text{F} = \tfrac{9}{5}(^{\circ}\text{C}) + 32$:
$$^{\circ}\text{F} = \frac{9}{5}(37) + 32 = 66.6 + 32 = 98.6\ ^{\circ}\text{F}.$$
Answer: $37\ ^{\circ}\text{C} = 310.15\ \text{K} = 98.6\ ^{\circ}\text{F}$, the familiar value for body temperature.
Gas-law problems need kelvin, not Celsius
In ideal-gas calculations, $PV = nRT$, the temperature $T$ must always be in kelvin. Substituting a Celsius value — for instance using 27 instead of 300 K for $27\ ^{\circ}\text{C}$ — is one of the most frequent silent errors in numericals. Always convert first using $K = {^{\circ}\text{C}} + 273.15$ (commonly approximated as $+\,273$).
When you see °C inside a gas, kinetic-theory or thermodynamics problem, convert to K before anything else.
Properties of Matter & Measurement in one glance
- Physical properties (colour, m.p., b.p., density) are measured without changing identity; chemical properties (combustibility, reactivity) need a chemical change.
- The SI system has seven base units: m, kg, s, A, K, mol, cd — all other quantities are derived.
- Prefixes scale a unit by powers of ten; mind the case (m = milli, M = mega).
- Mass (amount of matter) is constant; weight (force of gravity) varies with location.
- Volume SI unit is m³; the litre is non-SI: $1\ \text{L} = 1000\ \text{mL} = 1\ \text{dm}^3$.
- Density $=$ mass/volume; SI unit kg m⁻³, lab unit g cm⁻³.
- Temperature: $K = {^{\circ}\text{C}} + 273.15$ and $^{\circ}\text{F} = \tfrac{9}{5}(^{\circ}\text{C}) + 32$; kelvin is the SI unit and never negative.