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

Átomos e moléculas exibem um comportamento muito interessante quando absorvem e liberam energia. Os elétrons em um átomo tradicionalmente existem em seu estado de energia mais baixo, chamado de estado fundamental, e rotulados como n igual a um. No entanto, quando um átomo absorve energia, os elétrons ficam excitados e se movem para um nível de energia mais alto. Quando os elétrons relaxam para um estado de energia mais baixo ou para o estado fundamental, o excesso de energia é liberado como luz emitida.
Você pode estar familiarizado com esses estados de energia, pois eles são descritos no Modelo de Bohr, que descreve um átomo como um núcleo com elétrons em órbita em camadas ou orbitais. Essas conchas são iguais aos níveis de energia e são rotuladas por n. O comprimento de onda da luz emitida depende da diferença entre os níveis de energia altos e baixos. A luz emitida de alta energia resulta do relaxamento dos elétrons de um nível de energia mais alto, e a luz emitida de baixa energia resulta do relaxamento dos elétrons de um nível de energia mais baixo.
O espectro de emissão é uma medida da radiação emitida em uma faixa de comprimentos de onda. Com espécies elementares puras, o comportamento de emissão aparece como linhas de comprimentos de onda específicos, em vez de um amplo espectro. Como átomos diferentes têm níveis de energia diferentes, essas linhas espectrais variam de elemento para elemento e dependem das transições que esses elétrons fazem entre os estados de energia quando excitados. Por exemplo, existem seis séries nomeadas de linhas espectrais para hidrogênio, uma das quais é a Série Balmer.
A série de linhas espectrais de Balmer ocorre quando os elétrons fazem a transição de um nível de energia superior a n = 3 de volta para n = 2. O espectro de luz visível para a série Balmer aparece como linhas espectrais em 410, 434, 486 e 656 nm. A linha h alfa é a linha vermelha a 656 nm e ocorre devido à transição de n = 3 para n = 2. As linhas ciano, azul e violeta correspondem às transições de nível de energia de n = 4, 5 e 6, respectivamente, de volta para n = 2. Linhas espectrais adicionais podem ser medidas fora da faixa visível.
Johann Balmer quantificou as linhas visíveis usando a fórmula de Balmer. Aqui, lambda é o comprimento de onda observado, c é uma constante, n é o nível de energia mais baixo de dois e m é o nível de energia mais alto. A combinação da equação de Balmer e do Modelo de Bohr nos dá a equação de Rydberg, que descreve as linhas espectrais de muitos elementos diferentes. Nesta equação, lambda é o comprimento de onda registrado e RH é a constante de Rydberg. Os níveis iniciais marcados por n-inicial representam o nível de energia mais alto para o qual os elétrons se excitam e n final é o nível de energia mais baixo para o qual os elétrons relaxam de volta. Para a série de Balmer, nfinal = 2.
Neste laboratório, você medirá e observará os espectros de emissão de hidrogênio, hélio e neônio e usará a Equação de Rydberg para determinar a localização das linhas espectrais.

série balmer

O modelo de Bohr
Niels Bohr propôs um modelo para o átomo de hidrogênio em 1913 que descrevia que estados de energia discretos estão associados a uma órbita eletrônica fixa ao redor do núcleo. É importante ressaltar que um átomo não pode descarregar energia enquanto seus elétrons estão em estados estacionários. Um elétron só pode emitir energia alterando os estados de energia. Para mudar os estados de energia, um elétron deve se mover de uma órbita para outra, absorvendo ou emitindo energia. Essa mudança só pode ocorrer se a energia absorvida ou emitida for igual à diferença entre os dois estados. Os elétrons não podem existir entre as órbitas.
O número quântico, n, é usado para rotular os diferentes estados de energia. O estado de energia mais baixo é o estado fundamental, que é n igual a um. Os estados excitados são rotulados como n igual a 2, 3, 4 e assim por diante. Quando um elétron no estado fundamental absorve um fóton cuja energia é igual à diferença entre o estado fundamental e o segundo estado, o elétron fica excitado e faz a transição do estado fundamental para o estado excitado n = 2. Se a energia do fóton for igual à diferença entre o estado fundamental e o terceiro estado, o elétron se move para o estado n = 3.
De acordo com o modelo de Bohr, a energia potencial de um elétron noenésimo nível pode ser calculada usando a seguinte equação:
<img src="/CDNSource/lm/labs/54/54_Concepts_1.jpg" alt="Fórmula do nível de energia do modelo de Bohr, En = -Rhc / n², relacionada às transições eletrônicas, equação." data-altupdatedat="1754907201000">
onde En é a energia potencial, R é a constante de Rydberg (1,0974 × 107 m-1), h é a constante de Planck (6,62607004 × 10-34 m2·kg/s) e c é a velocidade da luz (~ 3 × 108 m/s). Os elétrons também podem retornar espontaneamente ao estado fundamental ou a qualquer outro estado excitado inferior. Quando isso acontece, o excesso de energia é liberado na forma de um fóton emitido. A energia do fóton é igual à diferença de energia entre os estados de energia superior e inferior. Essa energia corresponde aos comprimentos de onda da luz. Como cada tipo de átomo tem diferentes níveis de energia, a luz emitida por cada transição varia para cada átomo. Para uma amostra de moléculas mistas, a luz emitida contém uma faixa de comprimentos de onda no que é chamado de espectro contínuo. Para uma amostra contendo átomos de um único elemento, a luz emitida contém apenas certos comprimentos de onda, que podem ser vistos como linhas discretas, uma vez separadas por um prisma.
<h4>O átomo de hidrogênio</h4>
Olhando especificamente para o átomo de hidrogênio, a excitação de seus elétrons requer a absorção de energia suficiente para dividir a ligação na molécula diatômica H2. Como mais energia é usada para dividir a molécula do que o necessário, os elétrons no átomo de hidrogênio absorvem o excesso de energia e são excitados para um nível de energia mais alto. Quando os elétrons retornam espontaneamente a um nível de energia mais baixo, a luz é emitida, o que corresponde à diferença de energia entre o nível excitado e o nível inferior.
Ao discutir a emissão de energia, o nível de energia mais alto é considerado o nível inicial, ou ni, enquanto o nível mais baixo é considerado o nível final, ou nf. Os comprimentos de onda da luz emitida dependem, em última análise, da diferença de energia entre os dois níveis.
Em uma amostra pura de gás hidrogênio, o espectro de emissão aparece como linhas distintas de comprimentos de onda discretos que são específicos do elemento hidrogênio. Algumas dessas linhas estão na faixa visível do espectro eletromagnético, enquanto outras estão na faixa ultravioleta ou infravermelha.
<h4>A série Balmer</h4>
A série de linhas visíveis no espectro do átomo de hidrogênio é chamada de série de Balmer. Esta série de linhas de emissão espectral ocorre quando o elétron faz a transição de um nível de alta energia para o nível de energia mais baixo de n = 2. Johann Balmer observou essas linhas espectrais em 410,2 nm, 434,1 nm, 486,1 nm e 656,3 nm, que correspondem a transições dos níveis de energia n = 6, n = 5, n = 4 e n = 3 para o nível n = 2, respectivamente.
Balmer foi capaz de relacionar esses comprimentos de onda da luz emitida usando a fórmula de Balmer.
<img src="/CDNSource/lm/labs/54/54_Concepts_2.jpg" alt="Fórmula de cálculo do comprimento de onda λ=C(m²/(m²-n²)), equação para análise de espectroscopia." data-altupdatedat="1754907201000">
Aqui, λ é o comprimento de onda observado, C é uma constante (364,50682 nm), n é o nível de energia mais baixo com um valor de 2 e m é o nível de energia mais alto, que tem um valor maior que 3. Esta observação foi então refinada por Johannes Rydberg, onde R é a constante de Rydberg.
<img src="/CDNSource/lm/labs/54/54_Concepts_3.jpg" alt="Equação da fórmula de Rydberg; calcula linhas espectrais em espectroscopia atômica." data-altupdatedat="1754907201000">
Lembre-se de que esta equação descreve a luz emitida, portanto, o nível de energia mais alto é considerado o nível inicial, ou ni, enquanto o nível mais baixo é considerado o nível final, ou nf. No caso da série de Balmer, nf é igual a 2. Esta equação foi combinada com o modelo de Bohr para calcular a energia necessária para mover um elétron entre seus níveis de energia inicial e final, ΔE.
<img src="/CDNSource/lm/labs/54/54_Concepts_4.jpg" alt="Equação de transição do nível de energia \(\Delta E = Rhc(\frac{1}{n_f^2} - \frac{1}{n_i^2})\); fórmula física." data-altupdatedat="1754907201000">
Mais tarde, outras séries espectrais para o átomo de hidrogênio foram descobertas. Por exemplo, a série de Lyman contém linhas de emissão com energias na região ultravioleta.
<h4>Referências</h4>
<ol>
	<li style="font-size:17px">Kotz, J.C., Treichel Jr, P.M., Townsend, J.R. (2012). Química e Reatividade Química. Belmont, CA: Brooks / Cole, Cengage Learning.</li>
	<li style="font-size:17px">Silderberg, M.S. (2009). Química: A Natureza Molecular da Matéria e da Mudança. Boston, MA: Colina McGraw.</li>
</ol>

Série Balmer

O modelo de Bohr
Niels Bohr propôs um modelo para o átomo de hidrogênio em 1913 que descrevia que estados de energia discretos estão associados a uma órbita fixa de elétrons ...

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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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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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