What happens to electrons when an atom absorbs energy?

When an atom absorbs energy, electrons become excited and move from the ground state to higher energy levels. According to the Bohr Model, electrons occupy discrete energy states labeled by the quantum number n. An electron can only transition between these specific orbits if the absorbed energy exactly equals the difference between two energy states.

Why do different elements produce different spectral lines?

Each element has unique energy level spacing, so electrons transitioning between states release different amounts of energy. Since the wavelength of emitted light depends on the energy difference between levels, different atoms produce distinct spectral lines. This variation allows scientists to identify elements by their characteristic emission spectra.

What defines the Balmer Series in hydrogen?

The Balmer Series consists of visible spectral lines produced when hydrogen electrons transition from higher energy levels down to n=2. Johann Balmer observed four main lines at 410, 434, 486, and 656 nm, corresponding to transitions from n=6, 5, 4, and 3 respectively. The red h-alpha line at 656 nm results from the n=3 to n=2 transition.

How does the Rydberg equation predict spectral line wavelengths?

The Rydberg equation combines the Bohr Model with empirical observations to calculate wavelengths of emitted light. It uses the Rydberg constant, the initial energy level (n-initial), and final energy level (n-final) to determine the wavelength. For the Balmer Series, n-final equals 2, allowing prediction of all visible hydrogen spectral lines.

What is the relationship between energy difference and emitted light wavelength?

The wavelength of emitted light is directly determined by the energy difference between two energy levels. Larger energy differences produce higher-energy light with shorter wavelengths, while smaller differences produce lower-energy light with longer wavelengths. This relationship allows scientists to calculate energy transitions from observed spectral line positions.

Why does the emission spectrum appear as discrete lines rather than a continuous spectrum?

Pure elemental samples produce discrete spectral lines because electrons can only occupy specific energy levels and transition between them. Each transition releases a precise amount of energy corresponding to a single wavelength. In contrast, mixed molecular samples produce continuous spectra because they contain many different possible transitions.

How did Johann Balmer quantify the visible hydrogen spectral lines?

Balmer developed an empirical formula relating observed wavelengths to quantum numbers, using a constant value and the lower energy level of 2. His formula successfully predicted the four visible hydrogen lines and was later refined by Johannes Rydberg into a more general equation. This mathematical relationship bridged observation and the Bohr Model's theoretical framework.

Balmer Series

Concepts Instructor Prep Student Protocol

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Atoms and molecules exhibit very interesting behavior when they absorb and release energy. The electrons in an atom traditionally exist at their lowest energy state, called the ground state, and labeled as n equal to one. However, when an atom absorbs energy, the electrons become excited and move to a higher energy level. When the electrons relax down to either a lower energy state or to the ground state, the excess energy is released as emitted light.

You may be familiar with these energy states, as they are depicted in the Bohr Model, which describes an atom as a nucleus with orbiting electrons in shells or orbitals. These shells are the same as the energy levels and are labeled by n. The wavelength of the emitted light depends on the difference between the high and low energy levels. High energy emitted light results from electrons relaxing from a higher energy level, and low energy emitted light results from electrons relaxing from a lower energy level.

The emission spectrum is a measure of emitted radiation across a range of wavelengths. With pure elemental species, the emission behavior appears as lines of specific wavelengths rather than a broad spectrum. Since different atoms have different energy levels, these spectral lines vary from element to element and depend on the transitions those electrons make between energy states when excited. For example, there are six named series of spectral lines for hydrogen, one of which is the Balmer Series.

The Balmer Series of spectral lines occurs when electrons transition from an energy level higher than n = 3 back down to n = 2. The visible light spectrum for the Balmer Series appears as spectral lines at 410, 434, 486, and 656 nm. The h alpha line is the red line at 656 nm and occurs due to the transition from n= 3 to n = 2. The cyan, blue, and violet lines correspond to energy level transitions from n = 4, 5, and 6, respectively, back down to n = 2. Additional spectral lines can be measured outside of the visible range.

Johann Balmer quantified the visible lines using the Balmer formula. Here, lambda is the observed wavelength, c is a constant, n is the lower energy level of two, and m is the higher energy level. The combination of the Balmer equation and the Bohr Model gives us the Rydberg equation, which describes the spectral lines of many different elements. In this equation, lambda is the recorded wavelength, and R_H is the Rydberg constant. The initial levels marked by n-initial represent the higher energy level that the electrons excite to and n final is the lower energy level that the electrons relax back to. For the Balmer series, n_final = 2.

In this lab, you will measure and observe the emission spectrums of hydrogen, helium, and neon and use the Rydberg Equation to determine the location of the spectral lines.