Molecular Spectra & Fluorescence Applications - 2019 PDF

Molecular spectra, Jablonski diagram, fluorescence, fluorescence applications. The text under the slides was written by Ágnes Nagyné Szabó 2019 This instructional material was prepared for the biophysics lectures he...

Molecular spectra, Jablonski diagram, fluorescence, fluorescence applications. The text under the slides was written by Ágnes Nagyné Szabó 2019 This instructional material was prepared for the biophysics lectures held by the Department of Biophysics and Cell Biology Faculty of Medicine University of Debrecen Hungary https://biophys.med.unideb.hu During the lecture, after learning about the energy levels and spectra of atoms and molecules, we will study the Jablonski diagram in detail, discuss the phenomenon of absorption and different relaxation processes, and interpret Kasha’s rule. We will learn in more detail about the phenomenon of fluorescence and its parameters, and discuss how to measure fluorescence and absorption. Finally, we review some of the applications of fluorescence. On page 123 of the Medical Biophysics textbook, figure II.27. displays a similar electromagnetic spectrum showing the ranges of electromagnetic radiation as a function of frequency (f) and wavelength (λ). Radiations can be grouped in several ways, and different radiations have common properties. One of the most important among such principles is that energy propagates in every kind of radiation. Radiations usually start from some source and then reach the irradiated body after covering a certain distance, in which they can be partially or completely absorbed, scattered, or even reflected. Taking into account the physical nature of the radiations, the following radiations can be distinguished: 1) electromagnetic radiations: radiowave, microwave, infrared radiation, visible light, ultraviolet radiation, X-ray, gamma radiation 2) particle radiation: alpha-radiation, beta-radiation, proton radiation, etc. Electromagnetic waves do not require an elastic medium for propagation, they also propagate in vacuum. Their propagation speed (c) can be expressed by the formula c    f. Electromagnetic radiation, like light, has dual properties. The "particles" of electromagnetic radiation are photons, their energy (E) can be described by the following formula: c E photon  h  f photon  h  , where h is Planck’s-constant (h=6.63×10-34 Js), see in lecture 1. photon Higher energy photons have a higher frequency and a lower wavelength. The wavelength range of electromagnetic radiation visible to the human eye is between 400 and 750 nm. The first experimental evidence for quantized energy transfer was the Franck-Hertz experiment, in which two German physicists (James Franck and Gustav Ludwig Hertz) investigated the collision between electrons and mercury atoms (Hg) in 1914. The electron tube is filled with low-pressure Hg vapor and includes a third electrode, known as a grid, between the cathode and the anode, through which the electrons can pass. As a result of heating of the cathode, electrons are released (thermal emission) and accelerated by a voltage (U) applied between the cathode and the grid. Electrons could collide (both elastically and inelastically) with Hg atoms on their path, and if they then pass through the grid, the anode is reached only by those electrons that have enough kinetic energy to overcome the braking electric field (counter-field) of the anode and they generate current (I). The lower left figure on the slide shows the current (I) measured as a function of the accelerating voltage (U) applied between the cathode and the grid. As the accelerating  1  voltage increases, the kinetic energy of the electrons increases  eU  mv 2  , they collide  2  elastically with Hg atoms (they do not lose energy), pass through the grid, reach the anode, and, as expected, the current increases (the red line in the lower right figure illustrates this). However, at the accelerating voltage of U1=4.9V, the current drops suddenly, because the kinetic energy of electrons at the grid electrode is 4.9eV, which is completely transferred to Hg atoms near the grid in inelastic collisions, and the electrons are not able to overcome the counter-field and reach the anode (the blue line in the lower right figure illustrates this). As the accelerating voltage is further increased, the current rises again because the inelastic collision of the electrons with the Hg atoms occurs long before reaching the grid, and the electrons can be accelerated over the remaining distance enough to pass through the grid and reach the anode (the green line in the lower right figure illustrates this). When the accelerating voltage is U2=9.8V (2x4.9V), the current decreases again, because the electrons accelerated by 9.8V are able to reach the kinetic energy of 4.9eV twice along their journey. Once exactly halfway between the cathode and the grid, when their full 4.9eV energy is transferred to a Hg atom and they stop (Ekin=0). Electrons will be accelerated again up to 4.9V by the grid, and in the vicinity of the grid they collide with an Hg atom again, they lose their kinetic energy and do not pass through the grid. As the voltage is further increased, the phenomenon continues periodically. What is the explanation for the above phenomenon? When the electrons collided inelastically with the Hg atoms, they transferred their energy, thereby exciting the Hg atoms. However, Hg atoms could only absorb a certain amount of energy (4.9eV) from the electrons that collided with, which is exactly the same as the energy difference between the ground state and the first excited state of the Hg atom (illustrated in the upper left figure): E2  E1  4.9eV After excitation, the Hg atoms return to their ground state and emit a photon with an energy of 4.9eV. The experiment provided important evidence for the quantized atomic states and also supported the Bohr atomic model. In the 19th century, scientists studied absorption and emission spectra and found that glowing gases could only absorb photons of the same wavelength as the light they emit (the photon marked in green in the lower left figure of the slide), while photons of other wavelengths are simply transmitted. In this sense, the absorption spectrum is discrete, as those wavelengths are missing from the transmitted light that correspond to the emitted lines (indicated by the black band in the absorption spectrum in the lower right part of the slide), i.e. materials can only absorb and emit photons of a certain energy. The atom in the lower figure absorbed a photon of a given wavelength and energy ( E  h  f ) (marked in green), so that a lower energy (E1) electron in the first orbit gained enough energy to be placed in the second orbit, a higher energy level (E2), the atom was excited. The energy of the absorbed photon covered the energy difference between the two orbits: E  h  f  E2  E1. The lifetime of the excited state is extremely short, the electron placed in the higher energy orbit returns to the lower energy orbit and radiates the difference between the two energy levels in the form of a photon of the same energy as that absorbed during excitation (green photon in the figure). The process is called relaxation, and only the spectrum of this emitted photon will be visible in the discrete emission spectrum (green bar in the emission spectrum in the lower right figure of the slide). When an electric discharge is passed through a sample containing hydrogen gas, the H2 molecules decompose and the energetically excited H atoms emit light (photons) containing discrete frequency values (discrete energy). The first significant result in the explanation of the resulting line spectrum (the top right figure on the slide) was achieved in 1885 by the Swiss high school teacher Johann Balmer. He derived the Balmer formula, which describes the relationship between the wavelengths (λ) of the spectral lines in the visible range of spectrum 1  1 1  of the hydrogen atom (400-750nm):  RH   2  2  , n  3, 4,5,... where RH is the Rydberg  2 n  constant for the hydrogen atom. The set of spectral lines that can be described by the formula is called the Balmer series (the red lines in the spectrum of the hydrogen atom on the slide). The Balmer formula, which only gives the wavelengths of lines belonging to the Balmer series, 1  1 1  was generalized by Rydberg:  RH   2  2  , where n1 and n2 are the principal   n1 n2  quantum numbers of the orbits involved in the transition. In 1913, Bohr’s atomic model (electrons in an atom can only occupy certain energy orbits around the nucleus) could establish the discrete absorption and emission spectra of the hydrogen atom theoretically (see Lecture 4: X-rays for more details). Depending on which orbits are involved in the electronic transitions (E1, E2, E3,….), we distinguish between Lyman, Balmer, and Paschen series. In the Lyman series, the absorbed or emitted photons are of low wavelength and fall in the invisible ultraviolet range (energy higher than that of visible light). In this case the electron is excited from either the very first orbit (E1) to any higher energy orbit (E2, E3,….) , or relaxes from any of the higher energy orbits to the very first orbit (the green lines on the slide in the spectrum of the hydrogen atom). The photons in the Balmer series are already of lower energy, so they fall in the range of visible light, and the electron is either excited from the second orbit (E2) to any higher energy orbit (E3, E4,….) or relaxes from any higher energy orbit to the second orbit. The Paschen series includes the lowest energy photons in the infrared range, in which case the electron is either excited from the third orbit (E3) to any of the higher energy orbits (E4, E5,….) or relaxes from any of the higher energy orbits to the third orbit (blue lines in the spectrum of the hydrogen atom on the slide). Although the Bohr’s model successfully described the energy levels of the H atom and the spectral lines of all single-electron (“hydrogen-like”) ions, the spectral lines of multi-electron atoms and ions, as well as molecules, could only be explained by the quantum mechanical model. Atoms that make up molecules also move continuously within the molecule, vibrate along a covalent bond, can rotate, and longer molecules bend. These movements have a definite (quantized) energy. The spectrum examined in molecular spectroscopy is formed by the molecule absorbing or emitting a photon while its energy changes. The difference from atomic spectroscopy is that the energy of molecules can change not only through electronic transitions but also through rotational and vibrational transitions leading to the appearance of vibrational and rotational energy levels in the spectrum. Therefore, the spectrum of molecules is more complex than the spectrum of atoms. These transitions are usually summarized in an energy level system (bottom left figure), in which the series of vibrational energy levels (v=0, v=1) are superimposed on the electronic energy levels, and the system of rotational levels (r=0,1,2…) are added to each vibrational level. However, due to selection rules, not all transitions are allowed. Changes in the electron distribution of molecules require energies in the order of magnitude of a few thousand electron volts, so photons emitted or absorbed during such changes fall in the visible (400-750nm) and ultraviolet (200-400nm) regions of the spectrum. Significantly less energy is required for the rearrangement of molecular vibrations (vibrational transition) and even less for the transition between possible rotational movements of the molecule (rotational transition), so the energy range of these transitions falls in the range of infrared radiation (750nm-400μm). Why is the UV-visible spectrum of molecules banded and not discrete? The excitation of electrons is also accompanied by changes in the vibrational and rotational states of the molecule leading to many possible transitions (colored lines in the top left figure). These transitions are characterized by different energies and different wavelengths. However, these narrow spectral lines can only be observed in the gas phase of the molecules (left figure). In solution, interaction of the excited molecules with solvent molecules makes the energies of the different transitions more uncertain (represented by the wide gray bands in the top figures in the middle and on the right). Therefore, many more transitions with different energies are possible, whose energies overlap with each other. In these cases, only the spectral envelope is observed (continuous band) instead of individual spectrum lines (bottom right figure). Fluorescence: the relaxation of molecules with light emission, in which an electron returns from the first excited state (S1) to the ground state (S0) and emits the energy difference between the two states in the form of a fluorescence photon.  In the figure, process 1 is the excitation, which is usually accomplished by absorbing a photon with a certain energy of Eabs. photon  hf abs. photon , and the electron is transferred from the ground state (S0) to an excited state of higher energy (S1’).  Thereafter, through rapid thermal relaxation, which releases heat without photon emission, energy is lost (this is process 2 in the figure).  In process 3, the electron returns from the excited state (S1) to the lower energy ground state (S0) while emitting the energy difference between the two states in the form of a photon with an energy of Eem. photon  hf em. photon.  The energy of the emitted photon is lower than the energy of the absorbed photon due to heat dissipation during thermal relaxation. Therefore, the wavelength of the emitted photon is higher than that of the absorbed one, it is shifted towards the red (higher) wavelength range compared to the wavelength of the absorbed photon. This red shift is the Stokes’ shift Eabs. photon  Eem. photon , which makes the fluorescence a sensitive spectroscopic method (since unabsorbed or scattered excitation light can be separated from the emitted light by optical filters, so it does not interfere with the detection of fluorescence). The slide shows the Jablonski diagram, which summarizes the processes that can take place in a molecule upon absorption of light. In the figure, many horizontal lines belonging to one electronic state (S0, S1, S2 T1) indicate the vibrational energy levels. Rotational energy levels are generally not indicated for simplicity.  What is the difference between singlet (S0, S1, S2,..) and triplet (T1, T2,..) states? In these considerations we always investigate a pair of electrons since molecular orbitals are filled with two electrons. In the singlet state the number of unpaired electrons is zero, the spin of every electron is “neutralized” by another electron with an opposite spin  , so the resulting total spin quantum number of electrons is S=0, and the number of possible orientations of the total spin in a magnetic field is 1 (2S+1=1). In contrast, in the triplet state there are two electrons of the same spin with a resulting spin quantum number of +1  , and in magnetic field three different orientations of the total spin vector are possible (2S+1=3). Slide 35 in “Supplementary materials” shows a diagram of these conditions. The probability of electron transitions between singlet states is much higher than the probability of singlet-triplet transitions, since the latter involves spin inversion, therefore it is called a forbidden transition.  Upon absorption of a photon, the distribution of electrons is rearranged over a period of femtoseconds (10-15s) and the electron is transferred to a vibrational level of a higher energy excited singlet state, depending on the energy of the absorbed photon (black arrows pointing upward in the figure).  Because the excited state does not correspond to thermal equilibrium with the environment, thermal relaxation occurs in picoseconds (10-12s), during which only heat is released, no photon is emitted. In one or more steps the electron relaxes to the lowest vibrational level of the S1 excited singlet state. Vibrational relaxation is a thermal relaxation that occurs between vibrational transitions within the same electronic state (denoted by the kv rate constant in the figure). In contrast, internal conversion is a thermal relaxation that occurs between different singlet states (denoted by the kic rate constant in the figure).  From the lowest vibrational level of the S1 state, there are several possibilities for the electron to return to the S0 ground state depending on the molecule and its environment: o Internal conversion without photon emission, when the excitation energy is completely converted to heat (black dashed arrow in the figure marked with the kic rate constant). o Fluorescence with photon emission, when the electron returns to one of the vibrational levels of the ground state (S0) accompanied by the emission of a photon (red photons labeled by the rate constant kfl in the figure). Fluorescence lifetimes are on the order of nanoseconds (10-9s). According to Kasha’s rule, relaxation with photon emission can only occur from the lowest vibrational level of the S1 excited singlet state, since the relaxation processes leading here are much faster than fluorescence. o Another possibility is the transition between systems, intersystem crossing (characterized by the rate constant kisc in the figure), during which the electron moves from the S1 state to the lower energy excited triplet state (T1), which has a small probability of occurrence (forbidden transition), since it involves a spin flipping. One of the conditions for this to occur is that the lowest vibrational level of S1 overlaps with any vibrational level of T1. Let us see what processes can occur from the T1 excited triplet state:  Phosphorescence occurs if the electron relaxes to the ground state (S0) during the emission of a photon (characterized by the rate constant kph and labeled by the green photon in the figure). The lifetime of phosphorescence is significantly longer (10–6–10s) than that of fluorescence because both intersystem crossing and phosphorescence are forbidden transitions involving spin inversion. In addition, another important difference is that the energy of a photon emitted during phosphorescence is lower than the energy of a photon emitted during fluorescence (since the energy of the lowest vibrational level of T1 is lower than that of the lowest vibrational level of S1), therefore the phosphorescence spectrum shifts towards higher wavelength range (red) compared to the fluorescence spectrum.  Thermal relaxation without photon emission can also occur from the T1 state during internal conversion to the S0 ground state.  Delayed fluorescence occurs when the electron returns from the T1 state to the lowest vibrational level of the S1 state (indicated by the blue dashed arrows with the kisc rate constant in the figure) and then returns to the S0 state during photon emission (blue photons labeled by the kd.fl rate constant in the figure). Photon emission (fluorescence, phosphorescence) from the excited electronic states is collectively called luminescence. A luminescent material can be characterized by several parameters (spectral distributions, quantum efficiency, lifetime, degree of polarization), which will be introduced in more detail in the next few slides. One such parameter is the spectral distribution of absorption and emission:  Emission spectrum: plots the intensity of light emitted as a function of wavelength. The spectrum can be recorded by exciting the molecule at a given wavelength and measuring the intensity of the emitted fluorescence at different wavelengths with a spectrofluorimeter. The emission spectrum of the molecule represents the vibrational levels of the ground state, as these photons are generated when the electron relaxes from the lowest vibrational level of the excited singlet state (S1) to one of the vibrational levels of the S0 ground state.  Excitation spectrum: produced by measuring the light intensity of an emission at a single wavelength while continuously varying the excitation wavelength. The shape of the curve is usually the same as the absorption spectrum, which gives the wavelength dependence of absorbance. The excitation spectrum provides information on the vibrational levels of the excited state. In the case of atoms, the spectrum has a very narrow, discrete structure, but in the case of molecules in solution, it has a wider, banded structure. Tryptophan is a naturally luminescent amino acid whose absorption (purple curve) and emission spectra (blue curve for fluorescence, green curve for phosphorescence) are shown in the figure. In the upper right corner, the Jablonski diagram shows the electron transitions corresponding to the spectra with identical colors.  Due to the Stokes shift, the peak of the emission spectrum is always shifted toward longer wavelength (red shifted) with respect to the peak of the absorption spectrum, since the energy of the emitted photons is always lower than the energy of the absorbed photons.  The peak of the phosphorescence emission spectrum is always shifted toward longer wavelength ranges relative to the peak of the fluorescence spectrum because the energy of the photon emitted during phosphorescence is lower than the energy of the photon emitted during fluorescence.  The shape of the absorption and fluorescence spectra are mirror symmetric to each other.  The phosphorescence spectrum of tryptophan (green curve) was recorded at 77K, therefore it splits up into vibrational states: the first maximum corresponds to the photon energy between the lowest vibrational energy state of T1 and S0, and the other maxima show the transitions when relaxation takes place to one of the higher energy vibrational levels of S0. In the other two spectra, these details are blurred and not visible. Proteins generally have an absorption maximum at 280 nm, while nucleic acids have an absorption maximum at 260 nm. The absorbance of molecules can be determined with a spectrophotometer, the scheme of which is shown in the upper right figure. The instrument measures the intensity of the incident light (I0) and the light passing through the sample (I). The general law of intensity attenuation applies to the attenuation of light intensity during absorption: I  I 0 10  L , where I is the intensity of light after passing through a material, I0 is the intensity of the incident light, μ is the absorption coefficient and L is the thickness or optical path length. If a dilute solution is measured, the absorption coefficient can be expressed by the following formula:   c   , where c is the concentration of the solution (amount of absorbing particles) and ε is the molar absorption coefficient of the substance, whose value depends on the wavelength of light. The light absorption of dilute solutions is described by the Lambert-Beer law: I  I 0 10 cL. Transmittance (T) is the fraction of incident light which is transmitted through the sample: I , which can range from 0 (all light absorbed) to 1 (no light absorption). The absorbance T I0 (A, optical density) can be determined by this formula: A   log T    c  L. The emission of molecules can be measured with a spectrofluorimeter, whose scheme is shown in the figure. Two significant differences can be observed compared to the absorption spectrophotometer. In spectrofluorimeters we have to use two monochromators, since we have to be able to control not only the wavelength of the illuminating light, but we also have to select the desired (characteristic) wavelength component from the emitted light. The other difference is that here the optical elements performing the illumination and the detection are not arranged along one line, but they are most often at right angles to each other. Absorption spectrophotometry in the wavelength range of ultraviolet and visible light is essential in medical practice. It can be used to identify and quantify unknown substances, as well as provide information on their structure and interaction with the environment. Spectrophotometers are suitable instruments for determining absorption spectra. The light from the light source is made parallel by a collimator (a converging lens) and then the light rays obtained pass through a monochromator to select a narrow wavelength range, i.e. photons of certain wavelengths, with which the sample can be excited. The sample is placed into a cuvette of a given thickness (L) in front of the detector, and the intensity of the light passing through the sample is detected after the sample has absorbed some of the light (which is necessary for its excitation). A spectrofluorimeter can be used to determine the excitation or emission spectra of molecules. A spectrofluorimeter has a monochromator not only after the light source to select the photons of the appropriate wavelength, but also in front of the detector, since the photons emitted by the sample are also filtered based on their wavelength. In the most common arrangement the direction of observation of fluorescence is perpendicular to the direction of the excitation light, since we can completely prevent the excitation light from entering the detector in this way. Depending on the setting of the instrument, the excitation and emission spectra of a fluorescent molecule in solution can also be recorded. In addition to spectral distributions, the quantum efficiency (Q) of fluorescence is one of the most important parameters that characterizes the emission ability of a dye and can be determined by the ratio of the number of emitted fluorescence photons and the number of N emitted photon absorbed photons: Q . It can also be expressed by the ratio of the rate constant N absorbed photon of fluorescence and the sum of the rate constants characterizing the different relaxation k fl processes: Q . Its value is always less than 1, because the transition from the k fl  kic  kisc excited state to the ground state can occur not only by photon emission. Its value depends on the molecular structure and environmental parameters (temperature, solvent, viscosity, interaction with surrounding molecules). In case of a dye (fluorophore) which is used as a fluorescent marker, a high quantum efficiency is a requirement. Let us look at an example that we will also use in cell biology practice in the next semester to label nuclei. Propidium iodide is a dye intercalating into DNA. DNA-bound propidium iodide has a much higher quantum efficiency (Q>0.9) than the free dye in aqueous solution (Q=0.05). It means that the dye bound to DNA fluoresces much more intensely because, when bound, it is protected from collision with solvent molecules and thus from quenching of fluorescence. The fluorescence intensity (I) gives the number of photons emitted per unit time and it is proportional to the intensity of the incident light, absorption (cL), and the quantum efficiency. The fluorescence intensity is often affected by the pH of the solution and the quality of the solvent. The figure shows the exponential decrease in fluorescence intensity as a function of time. The number of molecules in the excited state decreases exponentially over time, since after the excitation is over, the molecules return to the ground state and the number of photons  ( k f  kic  kisc ) t emitted decreases over time. At a given time t, the equation N  N 0 e describes the number of molecules in the excited state (N). Fluorescence lifetime (τ) is the time during which the fluorescence intensity decreases to 1 1/e-times its initial value, i.e. to 37% of the initial intensity:  . k f  kic  kisc Since the excited molecule can relax not only by fluorescence, other relaxation processes without photon emission (internal conversion, intersystem crossing) must also be taken into account when calculating the fluorescence lifetime. The graph shows how the fluorescence lifetime can be read from the time-dependent decrease in fluorescence intensity. From the definition, fluorescence lifetime is the time at which the intensity decreases from the initial Imax intensity to 1/e-times Imax (Imax/e). Fluorescence lifetime (τ) is the average time spent by excited molecules in the excited state, which is influenced by other processes leading to relaxation (internal conversion, intersystem 1 crossing):  . Its value is between 10-9-10-7s and it is sensitive to the polarity k f  kic  kisc and pH of the microenvironment. The lifetime of a homogeneous system is a single value, but in the case of heterogeneous systems several components can be measured. During the time spent in the excited state, the molecule can undergo a number of interactions with its environment, resulting in a decrease in fluorescence lifetime:  Collisional quenching: a reduction in the intensity of light emitted by an excited molecule in the presence of a non-fluorescent molecule (quencher) whose electronic structure is suitable for absorbing energy from the excited molecule when colliding with it. The quencher radiates the absorbed energy in some form (e.g., heat). Diffusion greatly affects collisional quenching.  Fluorescence resonance energy transfer (FRET): a molecule in the excited state (donor) transfers its energy to a nearby (2-10nm) and spectrally appropriate molecule (acceptor) in a dipole-dipole interaction (nonradiative process), thus without photon emission returns to the initial state (see later).  Intersystem crossing (isc): the singlet-triplet transition also reduces the fluorescence lifetime (photon emission from the triplet state, but the phosphorescence lifetime is longer than the fluorescence lifetime)  Rotational motion The figure on the left is similar to the figure on slide 17, showing the decrease in fluorescence intensity (I) over time. The figure on the right represents the decrease in the number of undecayed nuclei (N) as a function of time during radioactive decay (see ‘Nuclear physics, nuclear binding energy, radioactivity, law of radioactive decay, radioactive series’ in more detail). In both cases, the fluorescence intensity (I0) and the number of undecayed nuclei (N0) are maximal at time t=0. Comparing the two figures, we can see that both functions can be described by the same exponentially decreasing curve from which the lifetime can be determined (τ and T): during this time the fluorescence intensity (I0/e) or the number of undecayed nuclei (N0/e) decreases to 37% of the initial value. The slope of the exponential curves depends on the quality of the fluorescent molecules and the radioactive nuclei, it decreases more steeply with a shorter lifetime. Although fluorescence and radioactive decay are completely different types of processes, the time dependence of both is described by an exponential function. This is because in both processes, the “event” (emission of a fluorescent photon, radioactive decay) occurs randomly, independently of the other atoms or molecules. Fluorophores are molecules that are able to fluoresce. This fluorescence may be intrinsic, natively present in certain materials that make up living cells, or it may be owed to extrinsic fluorophores selectively bound to cellular constituents. Fluorophores (especially external fluorophores) can be used to study biological systems (e.g., labeling cellular constituents), intra- and intermolecular interactions, molecular motions, or measure molecular distances (FRET), etc.  Intrinsic, native (internal) fluorophores: UV-excitable aromatic amino acids (tryptophan, tyrosine, phenylalanine), NADH, vitamins A, B, C. In most cases, native fluorescence, usually called autofluorescence, is preferred to be eliminated when extrinsic fluorescent dyes are used.  Extrinsic fluorophores: These are used to label various non-fluorescent cellular constituents. Fluorescent molecules can be bound to antibodies (antibodies bind with high affinity to the antigen they recognize), toxins (phalloidin: binds to F-actin, β-scorpion toxin: binds to a voltage-gated K-channel) or any other protein. Dyes are usually conjugated to an amino or sulfhydryl group in the protein. Fluorescent proteins can also be included in this group. The first type of these is the green fluorescent protein (GFP), a natural protein of a jellyfish called Aequorea victoria that gives the animal a greenish glow. GFP is widely used in biological research because its gene can be linked to the gene of any protein (e.g., actin). Thus, the fusion protein contains functional actin and fluorescent GFP. The intracellular location of actin can be examined by introducing the actin-GFP fusion  protein into cells. By modifying GFP, a wide variety of colored fluorescent protein derivatives were produced. More information on fluorophores can be found in the ‘Supplementary materials’ section. The figure summarizes the sizes of fluorophores most commonly used in biological applications. Fluorophores are most commonly used to visualize a biological target molecule. Since we want to study the behavior and distribution of the target molecule, optimally the fluorophore should not affect the properties of the target molecule. Therefore, it is desirable for the size of the fluorophore to be much smaller than the size of the biological target molecule it labels. We can see by how much the fluorescent dyes (Atto565, Cy3, Alexa488,…) are smaller than the antibodies (IgG) to which the dyes are bound, and thus the dye-labeled antibody can be used to label proteins or macromolecules in situ (in cells) or ex vivo (extracted from cells). Quantum dots (QDot) are semiconductor nanocrystals whose emission wavelength ranges depend on their size. They generally have a wide excitation and narrow emission spectrum, they are photostable and have a long fluorescence lifetime, however their size is commensurate with the size of the proteins, usually between 15-20nm. Fluorophores can be used in immunofluorescence (dye-labeled antibody binds to antigen), they may be indicators for intracellular pH or ion concentrations (the excitation or emission maximum of the dye or the fluorescence quantum efficiency of the dye varies depending on pH or ion concentration), but are also suitable for membrane potential measurement and DNA labeling. The figure summarizes the pros and cons of two different approaches for measuring fluorescence. Spectrofluorimetry, discussed in detail earlier, is used to analyze solutions only. Depending on the settings, either the excitation or the emission spectra of the solution containing the fluorescent molecule can be recorded. Photons of the appropriate wavelength are selected with a monochromator. Living cells can be examined with a fluorescence microscope. The emission range corresponding to the fluorophore is usually selected with an optical filter. The excitation wavelength of the fluorophore is either selected from the light of a broad-spectrum lamp with an optical filter, or a monochromatic laser is used (the latter is commonly used in confocal microscopy, see a later presentation). The microscope is suitable for studying the distribution of fluorescence intensity within a cell or tissue. We also use fluorescence microscopes in the biophysics practical to study fluorescent beads and to visualize the cellular components in the cell biology practice next semester. In addition to the above, fluorescence can even be measured in a flow cytometer (discussed in more detail in the ‘Flow cytometry and confocal microscopy’ lecture). The figure shows the schematic structure of a fluorescence microscope. From photons emitted by the light source covering a wide wavelength range the excitation filter is used to select those with which the fluorophore in the sample can be excited (the wavelength of the photons should overlap with the excitation spectrum of the fluorophore). In order to separate the excitation and emitted photons, a dichroic mirror is used, which reflects photons (excitation light) with a wavelength less than a certain value and transmits photons with a wavelength greater than this characteristic wavelength (emitted light). This mirror is positioned at an angle in the microscope to reflect the excitation light directly onto the sample. Because the wavelengths of the photons emitted by the sample are higher (their energy is lower) than the wavelengths of the excitation photons, they pass through the dichroic mirror to the emission filter without reflection, with which we can select which photons can enter the ocular based on their wavelength. Fluorescent molecules are clearly visible due to the dark background, so high-precision measurements can also be performed. For more information, see the ‘Optics, optical microscopy’ lecture. The figure shows how the excitation and emission filters and the dichroic mirror are arranged in a so-called filter cube. Optical filters can be divided into several groups depending on the wavelength range in which the photons are transmitted:  Short Pass filter (SP): photons with a wavelength less than a certain value can pass through the filter and photons with a longer wavelength will be blocked.  Long Pass filter (LP): in contrast to the previous one, photons with a wavelength higher than a certain value can pass through the filter, while it blocks those with a wavelength lower than this value.  Band Pass filter (BP): it allows to pass photons in a specific wavelength range and blocks photons outside this range.  Dichroic mirror (DM): reflects photons below a certain wavelength, while passing the rest (discussed in more detail on the previous slide). These images represent fluorescently labeled cellular components recorded by a fluorescence microscope. Photobleaching is one of the most common problems in microscopic imaging. Depending on the properties of the fluorophore, it can only go through a certain number of excitation- emission cycles, i.e. after a while it is no longer able to absorb photons. Photobleaching is a photochemical reaction in which the fluorophore irreversibly loses its absorbing capacity and its ability to fluoresce. The process of photobleaching starts from the excited state of the fluorophore in the same way as fluorescence. However, photobleaching often occurs after intersystem crossing, so it also includes the triplet state of the fluorophore. The process of photobleaching is influenced by the intensity of the light used during excitation and the duration of the excitation: the longer the illumination lasts, and the higher the intensity of the light is, the sooner photobleaching of the fluorophores occurs and the more fluorophore molecules are affected. The process is usually faster in the presence of oxygen, but it certainly depends on the properties of the fluorophore. In Figures a and b at the bottom right of the slide, the manifestations of photobleaching can be seen. (a) When fluorophores are exposed to short-term (