Atomic Emission Spectroscopy and Flame Tests Revision Notes for HSC SSCE Chemistry

Atomic Emission Spectroscopy and Flame Tests

What is atomic emission spectroscopy?

Under normal conditions, atoms don't emit light on their own, whether they exist in isolation or as part of compounds. However, when we give atoms extra energy—for example, by heating them—they can be made to emit light. This process forms the basis of atomic emission spectroscopy, a powerful analytical technique.

How atoms emit light

When atoms are heated to temperatures above 1500°C, something interesting happens to their electrons. Some electrons absorb this thermal energy and get "excited"—they jump from their normal energy levels (called orbitals) to higher energy levels. This normal state is called the ground state.

However, excited electrons don't stay in these higher energy levels for long. After a brief moment, they fall back down to their ground state. As they do this, the extra energy they absorbed is released as light. This light can be in the form of:

Visible light (what we can see)
Ultraviolet (UV) radiation
Infrared (IR) radiation

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Energy Conservation Principle

The energy emitted when an electron falls back to its ground state is exactly the same as the energy that was absorbed when it was raised to the excited state. This is a fundamental law of physics—energy cannot be created or destroyed, only transformed.

The diagram above shows how this works. Electron A absorbs energy $E_1$ to reach an excited state, then emits light corresponding to $E_1$ when it falls back. Similarly, Electron B absorbs energy $E_2$ and later emits light corresponding to $E_2$ . Each of these energy transitions produces a separate bright line in the emission spectrum.

The relationship between energy and wavelength

There's a mathematical relationship between the energy released by an excited electron and the wavelength of light produced:

$\Delta E = \frac{hc}{\lambda}$

Where:

$\Delta E$ is the energy released
$h$ is Planck's constant (a fundamental constant of nature)
$c$ is the speed of light
$\lambda$ (lambda) is the wavelength of the emitted radiation

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Inverse Relationship: Energy and Wavelength

This equation reveals a crucial relationship: the greater the energy released, the shorter the wavelength of light emitted. This inverse relationship means:

Large energy jumps produce UV light (short wavelengths)
Medium energy jumps produce visible light (medium wavelengths)
Small energy jumps produce IR radiation (long wavelengths)

Understanding emission spectra

When we take the light emitted by heated atoms and break it into its different wavelength components (for example, by passing it through a prism or using a spectroscope), we discover something remarkable: the emissions occur at just a few discrete (separate or distinct) wavelengths, not at all wavelengths.

This pattern of lines at specific wavelengths is called an emission spectrum. It appears as a set of bright or coloured lines on a black background—each line represents a specific electron transition in the atom.

The diagram above shows the emission spectra of four common elements: hydrogen, helium, sodium, and mercury. Notice how different each pattern is—this is crucial for identification purposes.

Why emission spectra are unique

Each wavelength in an emission spectrum corresponds to the energy required to excite a particular electron from its ground state to an excited state. Since different elements have different arrangements of electrons and different energy sublevels (such as 2p, 3p, or 3d), each element has its own unique emission spectrum. It's like a fingerprint for elements!

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The Power of Spectroscopic Identification

This uniqueness is incredibly useful. By examining the emission spectrum of an unknown sample, scientists can identify which elements are present. In fact, the element helium was first discovered by analyzing emission spectra from the Sun—hence its name, which comes from "Helios," the Greek word for Sun.

Flame tests for element identification

While emission spectroscopy uses sophisticated equipment, there's a simpler way to identify certain metallic elements: flame tests. When compounds containing some metallic elements are heated in a flame, they give the flame a distinctive colour.

How flame tests work

The high temperature of a Bunsen burner flame (typically 1000 to 1500°C) provides enough energy to:

Decompose the compound into its component elements
Excite electrons in atoms of one of the elements
Cause these excited electrons to emit light of characteristic wavelengths

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For flame tests to work, the element must be present as a compound that can vaporise at the flame's temperature. All elements that produce distinctive flame colours are metals.

Common flame colours

Different metallic elements produce different flame colours. Here are the most common ones you should know:

Element	Flame colour
Lithium	Carmine (dull red)
Sodium	Yellow
Potassium	Light purple (lilac)
Calcium	Brick-red (orange-red)
Strontium	Scarlet (deep red)
Barium	Pale green (apple green)
Copper	Blue-green

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Exam tip: A helpful mnemonic to remember these elements is "Little Sally Put Calcium Stones By Creek" (Lithium, Sodium, Potassium, Calcium, Strontium, Barium, Copper).

Why some elements produce distinctive colours

Elements that produce distinctive flame colours have a special characteristic: their atoms have one electron transition that occurs much more frequently than any other transition. This means that when the element is excited, it predominantly emits light of one particular wavelength, and therefore one colour.

Elements without this characteristic produce a mixture of many wavelengths, which results in white or unclear colours that aren't useful for identification.

Investigation 3.1: Conducting flame tests

Aim

To observe the colours produced by metal ions when placed in a Bunsen flame.

Safety considerations

What are the risks?	How to stay safe
Bunsen burner flame is hot and could cause burns	Set up on a heat-proof mat and place any used equipment onto the mat
Use of chemicals	Use the program RiskAssess or equivalent to determine safe handling procedures for each chemical

Key materials needed

Bunsen burner and heat-proof mat
Nichrome wire loop mounted in glass tubing
1 mol L $^{-1}$ hydrochloric acid
Distilled or demineralised water
Various metal salt samples (copper sulfate, calcium nitrate, barium chloride, potassium chloride, sodium chloride, strontium nitrate, etc.)

Basic procedure

Clean the nichrome loop by dipping it in distilled water, then in hydrochloric acid
Heat the loop in the blue Bunsen flame until the flame returns to its normal blue colour (this confirms the loop is clean)
Dip the clean loop into a solid metal salt sample
Place the loop with the sample into the blue Bunsen flame
Observe and record the flame colour
Repeat for each metal salt sample

What to look for

When analyzing results, compare the colours produced by:

Different metal salts containing the same metal (e.g., copper sulfate vs copper chloride)
Different salts with the same anion (e.g., different nitrates, chlorides, or sulfates)

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Key Finding

This helps you determine whether it's the cation (positive metal ion) or the anion (negative ion) that produces the distinctive colour. You'll find it's the metal cation that produces the colour.

Investigation 3.2: Observing emission spectra

Aim

To observe the emission spectra of different elements using discharge tubes.

Safety considerations

What are the risks?	How to stay safe
The induction coil (high-voltage source) emits X-rays	Have a teacher demonstrate equipment setup. Stand at least 2 metres away from the equipment

Key materials needed

Darkened room (essential for observation)
Discharge tubes containing different elements (hydrogen, helium, etc.)
Induction coil (high-voltage source)
Power pack
Spectroscope or digital spectrometer
Retort stand with clamp

Basic procedure

Set up the discharge tube in a retort stand
Connect the tube to the induction coil and power pack (teacher demonstration)
Darken the room and turn on the power pack
Look through the spectroscope to observe the spectrum
Record the wavelengths of spectral lines and their intensities (bright or faint)
Repeat with discharge tubes containing different elements

Important observations to make

The colour of light emitted from each discharge tube
The specific wavelengths of spectral lines in each spectrum
The gaps between spectral lines
Whether lines are faint or intense
How the spectra differ between elements

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Each element will show a unique pattern of lines, confirming that emission spectra are element-specific.

Practical applications: Fireworks

The distinctive flame colours produced by metallic elements have a spectacular practical application—fireworks! Pyrotechnicians use compounds containing specific metals to create different colours in fireworks displays.

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Connecting Theory to Practice

The beautiful colours we see in fireworks are actually the same atomic emission phenomenon we observe in flame tests and emission spectra—just on a much larger and more dramatic scale!

For example:

Red colours come from strontium compounds
Green colours come from barium compounds
Blue colours come from copper compounds
Yellow colours come from sodium compounds

Remember!

bookmarkSummary

Key Points to Remember:

Atoms emit light when energized: Under normal conditions, atoms don't emit light. However, when heated to high temperatures (above 1500°C), electrons absorb energy and move to excited states. When they fall back to ground states, they release this energy as light.
Each element has a unique emission spectrum: The pattern of spectral lines is like a fingerprint for each element because different elements have different electron energy levels. This makes emission spectroscopy a powerful tool for identifying elements.
Energy and wavelength are inversely related: The relationship $\Delta E = \frac{hc}{\lambda}$ shows that larger energy releases produce shorter wavelengths (UV light), while smaller energy releases produce longer wavelengths (IR radiation).
Flame tests identify metallic elements: Certain metals produce characteristic flame colours—lithium gives red, sodium gives yellow, potassium gives lilac, calcium gives brick-red, strontium gives scarlet, barium gives pale green, and copper gives blue-green.
One dominant transition creates distinctive colours: Elements that produce distinctive flame colours do so because they have one electron transition that occurs much more frequently than others, resulting in predominantly one wavelength (colour) of light being emitted.

Atomic Emission Spectroscopy and Flame Tests (HSC SSCE Chemistry): Revision Notes