Name: hydrogen emission lines
Uploaded: 2020-03-26T07:22:25.000Z
Duration: 3 min 53 s
Channel: JoVE

K12 Standards

Texas Essential Knowledge and Skills

High School

Chemistry and Physics

IPC.7

IPC.7.E

Balmer Series

Atomy i cząsteczki wykazują bardzo interesujące zachowanie, gdy pochłaniają i uwalniają energię. Elektrony w atomie tradycyjnie istnieją w swoim najniższym stanie energetycznym, zwanym stanem podstawowym i oznaczonym jako n równe jeden. Jednak gdy atom pochłania energię, elektrony zostają wzbudzone i przenoszą się na wyższy poziom energetyczny. Kiedy elektrony rozluźniają się do niższego stanu energetycznego lub do stanu podstawowego, nadmiar energii jest uwalniany w postaci emitowanego światła.
Być może jesteś zaznajomiony z tymi stanami energetycznymi, ponieważ są one przedstawione w modelu Bohra, który opisuje atom jako jądro z orbitującymi elektronami w powłokach lub orbitalach. Te powłoki są takie same jak poziomy energii i są oznaczone przez n. Długość fali emitowanego światła zależy od różnicy między wysokim i niskim poziomem energii. Światło emitowane o wysokiej energii wynika ze zrelaksowania się elektronów z wyższego poziomu energetycznego, a światło emitowane o niskiej energii wynika ze zrelaksowania elektronów z niższego poziomu energetycznego.
Widmo emisyjne jest miarą emitowanego promieniowania w zakresie długości fal. W przypadku czystych gatunków pierwiastkowych zachowanie emisyjne pojawia się jako linie o określonych długościach fal, a nie w szerokim spektrum. Ponieważ różne atomy mają różne poziomy energii, te linie widmowe różnią się w zależności od pierwiastka i zależą od przejść, jakie te elektrony tworzą między stanami energetycznymi po wzbudzeniu. Na przykład istnieje sześć nazwanych serii linii widmowych dla wodoru, z których jedną jest seria Balmera.
Seria linii widmowych Balmera występuje, gdy elektrony przechodzą z poziomu energii wyższego niż n = 3 z powrotem w dół do n = 2. Widmo światła widzialnego dla serii Balmer pojawia się jako linie widmowe przy 410, 434, 486 i 656 nm. Linia h alfa jest czerwoną linią o długości fali 656 nm i występuje w wyniku przejścia od n = 3 do n = 2. Niebieskozielone, niebieskie i fioletowe linie odpowiadają przejściom poziomu energii odpowiednio od n = 4, 5 i 6 z powrotem w dół do n = 2. Dodatkowe linie widmowe mogą być mierzone poza zakresem widzialnym.
Johann Balmer określił ilościowo widoczne linie za pomocą formuły Balmera. Tutaj lambda jest obserwowaną długością fali, c jest stałą, n jest niższym poziomem energii dwójki, a m jest wyższym poziomem energii. Połączenie równania Balmera i modelu Bohra daje nam równanie Rydberga, które opisuje linie widmowe wielu różnych pierwiastków. W tym równaniu lambda jest zarejestrowaną długością fali, a RH jest stałą Rydberga. Początkowe poziomy oznaczone n-początkiem reprezentują wyższy poziom energii, do którego wzbudzają się elektrony, a n końcowy to niższy poziom energetyczny, do którego elektrony się rozluźniają. Dla serii Balmera nkońcowe = 2.
W tym laboratorium będziesz mierzyć i obserwować widma emisyjne wodoru, helu i neonu oraz używać równania Rydberga do określenia położenia linii widmowych.

seria balmer

Model Bohra
Niels Bohr zaproponował model atomu wodoru w 1913 roku, który opisywał dyskretne stany energetyczne związane ze stałą orbitą elektronową wokół jądra. Co ważne, atom nie może rozładowywać energii, gdy jego elektrony znajdują się w stanach stacjonarnych. Elektron może emitować energię tylko poprzez zmianę stanów energetycznych. Aby zmienić stany energetyczne, elektron musi przemieszczać się z jednej orbity na drugą, pochłaniając lub emitując energię. Zmiana ta może nastąpić tylko wtedy, gdy pochłonięta lub wyemitowana energia jest równa różnicy między tymi dwoma stanami. Elektrony nie mogą istnieć między orbitami.
Liczba kwantowa, n, służy do oznaczania różnych stanów energetycznych. Najniższy stan energetyczny to stan podstawowy, który jest równy n jeden. Stany wzbudzone są oznaczone jako n równe 2, 3, 4 i tak dalej. Kiedy elektron w stanie podstawowym absorbuje foton, którego energia jest równa różnicy między stanem podstawowym a drugim, elektron ulega wzbudzeniu i przechodzi ze stanu podstawowego do stanu wzbudzonego n= 2. Jeśli energia fotonu jest równa różnicy między stanem podstawowym a trzecim, elektron przechodzi w stan n=3.
Zgodnie z modelem Bohra energię potencjalną elektronu nan-tym poziomie można obliczyć za pomocą następującego równania:
<img src="/CDNSource/lm/labs/54/54_Concepts_1.jpg" alt="Model Bohra wzór na poziom energetyczny, En=-Rhc/n², związany z przejściami elektronowymi, równanie." data-altupdatedat="1754907201000">
gdzie En jest energią potencjalną, R jest stałą Rydberga (1,0974 ×10 7 m-1), h jest stałą Plancka (6,62607004 × 10-34 m2·kg/s), a c jest prędkością światła (~ 3 × 108 m/s). Elektrony mogą również spontanicznie powrócić do stanu podstawowego lub dowolnego innego niższego stanu wzbudzonego. Kiedy tak się dzieje, nadmiar energii jest uwalniany w postaci wyemitowanego fotonu. Energia fotonu jest równa różnicy energii między wyższym i niższym stanem energetycznym. Energia ta odpowiada długościom fal światła. Ponieważ każdy typ atomu ma inny poziom energii, światło emitowane z każdego przejścia jest różne dla każdego atomu. W przypadku próbki mieszanych cząsteczek emitowane światło zawiera zakres długości fal w tak zwanym widmie ciągłym. W przypadku próbki zawierającej atomy pojedynczego pierwiastka, emitowane światło zawiera tylko pewne długości fal, które można postrzegać jako dyskretne linie oddzielone pryzmatem.
<h4>Atom wodoru</h4>
Patrząc konkretnie na atom wodoru, wzbudzenie jego elektronów wymaga absorpcji energii wystarczającej do rozszczepienia wiązania w dwuatomowej cząsteczceH2. Ponieważ do rozszczepienia cząsteczki zużywa się więcej energii niż jest to potrzebne, elektrony w atomie wodoru pochłaniają nadmiar energii i są wzbudzane do wyższego poziomu energetycznego. Kiedy elektrony spontanicznie wracają na niższy poziom energetyczny, emitowane jest światło, które odpowiada różnicy energii między poziomem wzbudzonym a niższym.
Omawiając emisję energii, wyższy poziom energii jest uważany za poziom początkowy, czyli ni, podczas gdy niższy poziom jest uważany za poziom końcowy, czyli nf. Długości fal emitowanego światła ostatecznie zależą od różnicy energii między tymi dwoma poziomami.
W czystej próbce gazowego wodoru widmo emisyjne pojawia się jako wyraźne linie dyskretnych długości fal, które są specyficzne dla pierwiastka wodoru. Niektóre z tych linii znajdują się w zakresie widzialnym widma elektromagnetycznego, podczas gdy inne znajdują się w zakresie ultrafioletu lub podczerwieni.
<h4>Seria Balmer</h4>
Seria widzialnych linii w widmie atomów wodoru nazywana jest szeregiem Balmera. Ta seria widmowych linii emisyjnych występuje, gdy elektron przechodzi z wysokiego poziomu energetycznego na niższy poziom energetyczny n = 2. Johann Balmer obserwował te linie widmowe przy 410,2 nm, 434,1 nm, 486,1 nm i 656,3 nm, które odpowiadają przejściom od poziomów energetycznych n=6, n=5, n=4 i n=3 do poziomu n=2.
Balmer był w stanie powiązać te długości fal emitowanego światła za pomocą wzoru Balmera.
<img src="/CDNSource/lm/labs/54/54_Concepts_2.jpg" alt="Wzór na obliczenie długości fali λ=C(m²/(m²-n²)), równanie do analizy spektroskopowej." data-altupdatedat="1754907201000">
Tutaj λ jest obserwowaną długością fali, C jest stałą (364,50682 nm), n jest niższym poziomem energii o wartości 2, a m jest wyższym poziomem energii, który ma wartość większą niż 3. Obserwacja ta została następnie udoskonalona przez Johannesa Rydberga, gdzie R jest stałą Rydberga.
<img src="/CDNSource/lm/labs/54/54_Concepts_3.jpg" alt="równanie wzoru Rydberga; Oblicza linie widmowe w spektroskopii atomowej." data-altupdatedat="1754907201000">
Pamiętaj, że to równanie opisuje emitowane światło, więc wyższy poziom energii jest uważany za poziom początkowy, czyli ni, podczas gdy niższy poziom jest uważany za poziom końcowy, czyli nf. W przypadku szeregu Balmera nf jest równe 2. Równanie to połączono z modelem Bohra w celu obliczenia energii potrzebnej do przemieszczenia elektronu między jego początkowym i końcowym poziomem energii, ΔE.
<img src="/CDNSource/lm/labs/54/54_Concepts_4.jpg" alt="Równanie przejścia poziomu energii \(\Delta E = Rhc(\frac{1}{n_f^2} - \frac{1}{n_i^2})\); wzór fizyczny." data-altupdatedat="1754907201000">
Później odkryto inne szeregi widmowe dla atomu wodoru. Na przykład seria Lymana zawiera linie emisyjne o energiach w zakresie ultrafioletu.
<h4>Odwołania</h4>
<ol>
	<li style="font-size:17px">Kotz, J.C., Treichel Jr, P.M., Townsend, J.R. (2012). Chemia i reaktywność chemiczna. Belmont, Kalifornia: Brooks/Cole, Cengage Learning.</li>
	<li style="font-size:17px">Silderberg, M.S. (2009). Chemia: molekularna natura materii i zmiana. Boston, Massachusetts: McGraw Hill.</li>
</ol>

Seria Balmer

Model Bohra
Niels Bohr zaproponował model atomu wodoru w 1913 roku, który opisywał dyskretne stany energetyczne związane ze stałą orbitą elektronów ...

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hydrogen emission lines

Hydrogen emission lines show how electrons move between fixed energy levels in the Bohr model. Niels Bohr proposed this model for the hydrogen atom in 1913. In this model, an electron can only occupy certain orbits, and it cannot exist between them.

These energy states are labeled with the quantum number n. The ground state is n = 1. Higher values, such as n = 2, 3, 4, and so on, are excited states. An electron changes state only by absorbing or emitting energy equal to the difference between two levels. A photon with the right energy can move an electron from one orbit to another.

Bohr’s model also links each energy level to a specific amount of potential energy. That value can be calculated with an equation that uses the Rydberg constant, Planck’s constant, and the speed of light. When an electron returns to the ground state or another lower excited state, it releases the extra energy as a photon. The photon energy matches the gap between the two states, so it also determines the wavelength of the emitted light.

Because each atom has its own energy levels, each element gives off its own pattern of light. A mixed sample can produce a continuous spectrum with many wavelengths. A pure sample of one element produces only certain wavelengths. When that light passes through a prism, it appears as separate lines instead of one smooth band.

Hydrogen shows this pattern clearly. Exciting hydrogen atoms requires enough energy to split the H2 molecule, and the extra energy excites the hydrogen electrons to higher levels. When those electrons drop back down, they emit light at wavelengths set by the energy difference between the initial level, n i, and the final level, n f. In pure hydrogen gas, this emission spectrum contains distinct visible, ultraviolet, and infrared lines that are unique to hydrogen.

The visible lines in hydrogen are called the Balmer series. They occur when an electron falls to n = 2 from a higher level. Johann Balmer observed lines at 410.2 nm, 434.1 nm, 486.1 nm, and 656.3 nm, which match transitions from n = 6, 5, 4, and 3 to n = 2. He described these wavelengths with the Balmer formula, where λ is the observed wavelength, C is a constant of 364.50682 nm, n is the lower level, and m is the higher level.

Later, Johannes Rydberg refined this idea using the Rydberg constant. For the Balmer series, the final level n f equals 2. The same relationship, combined with the Bohr model, can be used to calculate the energy change, ΔE, for electron transitions. Other hydrogen spectral series were later identified as well, including the Lyman series, which lies in the ultraviolet region.

Hydrogen Emission Lines

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uv-vis spectroscopy of dyes

How Dyes Absorb Visible Light

Dyes absorb visible light because their molecular structure lets them interact with specific wavelengths. When light hits a substance, some light is absorbed and the rest is reflected or transmitted. The color we see depends on the wavelength that is reflected. A blue substance reflects blue light, around 430 to 480 nm, and absorbs the complementary orange region, around 590 to 630 nm. Compounds that do not absorb in the visible range appear colorless to the human eye. Light also has energy and...

identifying alcohols

Alcohol Tests for Primary, Secondary, and Tertiary

Alcohol tests can help sort alcohols by structure and reactivity. Alcohols are organic molecules with a hydroxyl group attached to an alkyl or aryl group, written as ROH. If the carbon with the hydroxyl group has one R group, the alcohol is primary. If it has two R groups, it is secondary. If it has three R groups, it is tertiary. Alcohols can also be aromatic when they contain a benzene ring. Phenol is the simplest aromatic alcohol. To identify an alcohol, chemists can use its chemical...

chemistry

Comprehensive, curriculum-focused videos for both instructors and students in introductory chemistry lab courses.

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Explore how hydrogen electrons jump between energy levels and produce discrete emission lines, including the visible Balmer series and other hydrogen spectra.

Hydrogen Emission Lines | JoVE K12

Louisiana Student Standards for Science

motion-and-stability-forces-and-interactions

Motion And Stability: Forces And Interactions

HS-PS2-6

Tennessee Academic Standards

waves-and-their-applications-in-technologies-for-information-transfer

Waves and Their Applications in Technologies for Information Transfer

CHEM1.PS4.1

Next Generation Science Standards

Earth and Space Science

Earth’s Place in the Universe

HS-ESS1-2

south-carolina-college-and-career-ready-science-standards

South Carolina College and Career Ready Science Standards

Matter and Its Interactions

C-PS1-4

Georgia Standards of Excellence

9-12

SC1.

Light Emission and Electron Movement

Przygotuj się do nauczania Seria Balmer dzięki temu 1-minutowemu przewodnikowi dla instruktorów. Obejmuje kluczowe kroki przygotowawcze dla laboratorium Chemistry na JoVE Lab Manual.

Przygotowanie instruktora do Seria Balmer w Chemistry - Ponad 2.0K wyświetleń | JoVE Lab Manual

Postępuj zgodnie z procedurami Seria Balmer krok po kroku w tym 4-minutowym filmie z protokołem studenckim. Część kolekcji Chemistry na JoVE Lab Manual.

Protokół studencki dla Seria Balmer w Chemistry - Ponad 5.1K wyświetleń | JoVE Lab Manual

The Atomic Theory of Matter

The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers. Leucippus and Democritus argued that all matter ...

Subatomic Particles

Dalton was only partially correct about the particles that make up matter. All matter is composed of atoms, and atoms are composed of three smaller ...

Elements: Chemical Symbols and Isotopes

A chemical symbol is an abbreviation used to indicate an element or an atom of an element. For example, the symbol for mercury is Hg. The same symbol is ...

Ions and Ionic Charges

In ordinary chemical reactions, the nucleus &mdash; which contains the protons and neutrons of each atom and thus identifies the element &mdash; remains ...

The Periodic Table

As early chemists discovered more elements, they realized that various elements could be grouped by their similar chemical behaviors. One such grouping ...

Atomic Mass

Atoms &mdash; and the protons, neutrons, and electrons that compose them &mdash; are extremely small. For example, a carbon atom weighs less than 2 ...

Molar Mass

The identity of a substance is defined not only by the types of atoms or ions it contains but by the quantity of each type of atom or ion. For example, ...

Periodic Classification of the Elements

The periodic table arranges atoms based on increasing atomic number so that elements with the same chemical properties recur periodically. When their ...

Types of Chemical Bonds

Chemical bonding theories were pioneered by American chemist Gilbert&#xA0;N. Lewis. He developed a model called the Lewis model to explain the type and ...

Ionic Bonding and Electron Transfer

Ions are atoms or molecules bearing an electrical charge. A cation (a positive ion) forms when a neutral atom loses one or more electrons from its valence ...

The Born-Haber Cycle

Lattice Energy&#xA0;

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a ...

trends-in-lattice-energy-ion-size-and-charge

Trends in Lattice Energy: Ion Size and Charge

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a compound is a measure ...

Covalent Bonding and Lewis Structures

Compared to ionic bonds, which results from the transfer of electrons between metallic and nonmetallic atoms, covalent bonds result from the mutual ...

Electronegativity

Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity.&#xA0;

Electronegativity values of ...

bond-polarity-dipole-moment-and-percent-ionic-character

Bond Polarity, Dipole Moment, and Percent Ionic Character

Bond Polarity

The absolute value of the difference in electronegativity (&Delta;EN) of two bonded atoms provides a rough measure of the polarity to be ...

lewis-structures-of-molecular-compounds-and-polyatomic-ions

Lewis Structures of Molecular Compounds and Polyatomic Ions

To draw Lewis structures for complicated molecules and molecular ions, it is helpful to follow a step-by-step procedure as outlined:


	Determine the ...

Bond Energies and Bond Lengths

Stable molecules exist because covalent bonds hold the atoms together. The strength of a covalent bond is measured by the energy required to break it, ...

Bonding in Metals

Metallic bonds are formed between two metal atoms. A simplified model to describe metallic bonding has been developed by Paul Dr&#xFC;de called the ...

Model basic atomic structure and relate an element's atomic structure to its bonding, reactivity, and placement on the Periodic Table.

IPC.7.A

physical-and-chemical-properties-of-matter

Physical and Chemical Properties of Matter

The characteristics that enable us to distinguish one substance from another are called properties.

Physical Properties of Matter

A physical property is ...

atomic-radii-and-effective-nuclear-charge

Atomic Radii and Effective Nuclear Charge

The elements in groups of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number ...

Ionic Radii

Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it ...

Ionization Energy

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy ...

Electron Affinity

The electron affinity (EA) is the energy change for adding an electron to a gaseous atom to form an anion (negative ion).



This process can be either ...

Alkali Metals

Group 1 elements are soft and shiny metallic solids. They are malleable, ductile, and good conductors of heat and electricity. The melting points of the ...

Halogens

Group 17 elements, known as halogens, are nonmetals. At room temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine a solid. ...

Use patterns within the Periodic Table to predict the relative physical and chemical properties of elements.

IPC.7.B

Science concepts. The student knows that relationships exist between the structure and properties of matter. The student is expected to:

Hydrogen Emission Lines

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