Photoelectric Effect, Atomic Spectra, and Bohr Model Flashcards

A comprehensive vocabulary review of the photoelectric effect, atomic emission spectra, and Bohr's atomic model based on the lecture notes.

Photoelectric Effect

The phenomenon where light with a frequency of $\nu \ge \nu_0$ ejects electrons from a metal surface.

Threshold Frequency ($\nu_0$)

The minimum frequency of incident light required to eject electrons from a metal surface.

Work Function ($\Phi$)

The minimum energy required to remove an electron from a metal surface, defined as $\Phi = h\nu_0$.

Photon

A packet or quantum of light energy defined by the equation $E_{\text{photon}} = h\nu$.

Photoelectric Effect Equation

The energy conservation equation $h\nu = \frac{1}{2}m_e u^2 + \Phi$, relating incoming photon energy to the electron's kinetic energy and the metal's work function.

Spectroscopy

The study of the interaction between light and matter.

Emission Spectra

Distinct, non-continuous lines of light radiated by excited gas particles that are unique to and characteristic of each element.

Rutherford's Atomic Model

An atomic model featuring a positively charged nucleus surrounded by electrons, which failed to explain why electrons do not spiral into the nucleus or why atomic spectra are discrete.

Balmer's Equation

An empirical formula, $\lambda = B \frac{m^2}{m^2 - n^2}$, developed to describe the observed spectral lines of hydrogen.

Rydberg's Equation

The equation $\frac{1}{\lambda} = R Z^2 \left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)$ used to calculate the wavelengths of atomic spectral lines for single-electron species.

Rydberg Constant ($R$)

A fundamental physical constant equal to $1.096776 \times 10^7\,m^{-1}$.

Wave Number

The reciprocal of wavelength, represented mathematically as $\frac{1}{\lambda}$.

Stationary States

Specific allowed energy levels in an atom where an electron can orbit without emitting or losing energy.

Ground State

The lowest energy state of an atom where the electron resides in the orbit closest to the nucleus ($n = 1$).

Excited State

Any state in which an electron occupies an orbit with a principal quantum number $n > 1$, higher in energy than the ground state.

Bohr Orbit Energy Equation

The formula $E = -2.18 \times 10^{-18} \frac{Z^2}{n^2}$ used to calculate the energy of an electron in orbit $n$ for a one-electron species.

UV Region Spectral Transitions

Electron energy emission transitions in hydrogen that end at $n = 1$, producing spectral lines in the ultraviolet region.

Visible Region Spectral Transitions

Electron energy emission transitions in hydrogen that end at $n = 2$, producing spectral lines in the visible range.

IR Region Spectral Transitions

Electron energy emission transitions in hydrogen that end at $n = 3$, producing spectral lines in the infrared region.

The energy difference between the n=1 and n=2 orbital is

The largest. Followed by n=2 to n=3 and so on.

What does z represent in Rydberg’s equation?

Atomic number. Hydrogen is 1, helium is 2, etc.

The energy of each orbit is

Quantized (depends on n)

The energy difference between two orbits is

À discrete value and results in a discrete spectral line

Emission

If an electron goes from a higher to lower energy level, it emits a photo whose energy = difference between the two levels

Absorption

electron absorbs à photon whose energy equals the difference between lower and higher energy levels. Electron moves to the outer (higher energy level)

Boar’s model

Applies to H atom

H atom is only allowed certain energy levels where electrons can reside (stationary states) and moves between states by absorbing or emitting a photon

Higher energy level - further from nucleus