Chapter 5 Electrons in Atoms. Section 1 Light and Quantized Energy

Chapter 5

Electrons in Atoms

Section 1

Light and Quantized Energy

Light, a form of electromagnetic radiation, has characteristics of both a wave and a particle.

Section 1: Light and Quantized Energy

KWhat I Know

WWhat I Want to Find Out

LWhat I Learned

The Atom and Unanswered Questions

• Recall that in Rutherford's model, the atom’s mass is concentrated in the nucleus and electrons move around it.

– The model doesn’t explain how the electrons were arranged around the nucleus.

– The model doesn’t explain why negatively charged electrons aren’t pulled into the positively charged nucleus.

• In the early 1900s, scientists observed certain elements emitted visible light when heated in a flame.

• Analysis of the emitted light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms.


Flame test

http://www.youtube.com/watch?v=jJvS4uc4TbU




The Wave Nature of Light

• Visible light is a type of electromagnetic radiation, a form of energy that exhibits wave-like behavior as it travels through space.

• All waves can be described by several characteristics.

• The wavelength (λ) is the shortest distance between equivalent points on a continuous wave.

• The frequency (ν) is the number of waves that pass a given point per second.

• The amplitude is the wave’s height from the origin to a crest.



• The speed of light (3.00 × 108 m/s) is the product of it’s wavelength and frequency c = λν.



• Sunlight contains a continuous range of wavelengths and frequencies.

• A prism separates sunlight into a continuous spectrum of colors.

• The electromagnetic spectrum includes all forms of electromagnetic radiation.


Electromagnetic Spectrum

https://www.youtube.com/watch?v=m4t7gTmBK3g



Properties of Waves• Wave length:

– The length of each wave• λ = c/v

Constant: speed of light (c) 3.00 × 108 m/s

Frequency (v) Hz

Wavelength (λ) m

c

vλ

CalculationsMicrowaves are used to cook food and transmit information. What is the wavelength of a microwave that has a frequency of 3.44 × 109 Hz?KNOWN UNKNOWN

ν = 3.44 × 109 Hz

λ = ? m

c = 3.00 × 108 m/s

λ = c/v

Calculations

• What is the frequency of a microwave that has a wavelength of 6.72 × 10-9 m?KNOWN UNKNOWN

λ = ν =

c =

λ = c/v

The Particle Nature of Light

• The wave model of light cannot explain all of light’s characteristics. Some examples include:

– Why heated objects emit only certain frequencies of light at a given temperature.

– Why some metals emit electrons when light of a specific frequency shines on them.


The Particle Nature of Light

The Quantum Concept

• In 1900, German physicist Max Planck (1858-1947) began searching for an explanation of this phenomenon as he studied the light emitted by heated objects.

• Planck’s study led him to a startling conclusion:

– Matter can gain or lose energy only in small, specific amounts called quanta.

– A quantum is the minimum amount of energy that can be gained or lost by an atom.

– Planck’s constant has a value of 6.626 × 10–34 J ● s.


The Particle Nature of Light The Photoelectric Effect

• The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it.


Photoelectric Effect

http://www.youtube.com/watch?v=kcSYV8bJox8




The Particle Nature of Light Light’s Dual Nature

• Albert Einstein proposed in 1905 that light has a dual nature.

– A beam of light has wavelike and particle-like properties.

– A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.


Planck

http://www.youtube.com/watch?v=AjnBGWLAoZY




Energy of a Photon• Ephoton = hv

Ephoton energy

Frequency (v) Hz

H (Planck’s contstant) 6.626 × 10–34 J ● s.

Ephoton

vh

CalculationsEvery object gets its color by reflecting a certain portion of incident light. The color is determined by the wavelength of the reflected photons, thus by their energy. What is the energy of a photon from the violet portion of the Sun’s light if it has a frequency of 7.230 × 1014 s-1?

KNOWN UNKNOWN

ν = 7.230 × 1014 s-1 Ephoton = ? J

h = 6.626 × 10-34 J•s

Ephoton = hv

Atomic Emission Spectra

• Light in a neon sign is produced when electricity is passed through a tube filled with neon gas and excites the neon atoms.

• The excited atoms return to their stable state by emitting light to release energy.


Atomic Emission Spectra

• The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by the atoms of the element.

• Each element’s atomic emission spectrum is unique.


Essential Questions• How do the wave and particle natures of light compare?

• How is a quantum of energy related to an energy change of matter?

• How do continuous electromagnetic spectra and atomic emission spectra compare and contrast?


Section 2

Quantum Theory and the Atom

Wavelike properties of electrons help relate atomic emission spectra, energy states of atoms, and atomic orbitals.

Section 2: Quantum Theory and the Atom

KWhat I Know


LWhat I Learned

Bohr’s Model of the Atom

• Einstein’s theory of light’s dual nature accounted for several unexplainable phenomena but not why atomic emission spectra of elements were discontinuous rather continuous.

• In 1913, Niels Bohr, a Danish physicist working in Rutherford’s laboratory, proposed a quantum model for the hydrogen atom that seemed to answer this question.

• Bohr correctly predicted the frequency lines in hydrogen’s atomic emission spectrum.

• The lowest allowable energy state of an atom is called its ground state.

• When an atom gains energy, it is in an excited state.



• Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits.



• Each orbit was given a number, called the quantum number.

• Hydrogen’s single electron is in the n = 1 orbit in the ground state.

• When energy is added, the electron moves to the n = 2 orbit.


Atomic orbitals• Orbital- is a region of space around the nucleus where an electron is

likely to be found.• An electron cloud is a good approximation of how electrons behave in

their orbitals• The level in which an electron has the least energy—the lowest energy

level—has only one orbital. Higher energy levels have more than one orbital

Electron Configuration

• The most stable electron configuration is the one in which the electrons are in orbitals with the lowest possible energies.

Energy Level

Number of Orbitals

Max. # of Electrons

1 1 2

2 4 8

3 9 18

4 16 32

Energy Levels, Orbitals, and Electrons

DRAWING BOHR MODELS OF ATOMS

• Draw the nucleus and label with #p & #n• Draw electron orbitals

–1st orbital can have 2 electrons only–2nd and 3rd ring can each have

_8___electrons–Fill – _First ring FIRST______– Inner rings must be filled first before any

electron enters a higher ring!!!!!–Let’s Practice:

Bohr Model

http://www.youtube.com/watch?feature=player_embedded&v=nVW1zDPPZGM

http://www.youtube.com/watch?v=URbtiIpyAYY



Bozeman Science: Bohr

https://www.youtube.com/watch?v=GhAn8xZQ-d8




Bohr Model - Hydrogen

Bohr Model - Oxygen

Bohr Model - Sulfur

Bohr’s Model of the AtomThe limits of Bohr’s model

• Bohr’s model explained the hydrogen’s spectral lines, but failed to explain any other element’s lines.

• The behavior of electrons is still not fully understood, but substantial evidence indicates they do not move around the nucleus in circular orbits.


The Quantum Mechanical Model of the Atom

• Heisenberg showed it is impossible to take any measurement of an object without disturbing it.

• The Heisenberg uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

• The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.


The Quantum Mechanical Model of the Atom

• Schrödinger treated electrons as waves in a model called the quantum mechanical model of the atom.

• Schrödinger’s equation applied equally well to elements other than hydrogen.

• The wave function predicts a three-dimensional region around the nucleus called the atomic orbital.


Hydrogen Atomic Orbitals

• Principal quantum number (n) indicates the relative size and energy of atomic orbitals.

• n specifies the atom’s major energy levels, called the principal energy levels.

• Energy sublevels are contained within the principal energy levels.



• Each energy sublevel relates to orbitals of different shape.




Essential Questions• How do the Bohr and quantum mechanical models of the

atom compare?

• What is the impact of de Broglie’s wave-particle duality and the Heisenberg uncertainty principle on the current view of electrons in atoms?

• What are the relationships among a hydrogen atom’s energy levels, sublevels, and atomic orbitals?


Section 3


Three rules are used to determine electron arrangement in an atom.

Section 3: Electron Configuration

KWhat I Know


LWhat I Learned

Bozeman Science: Electron Configuration

https://www.youtube.com/watch?v=2AFPfg0Como




Ground-State Electron Configuration

• The arrangement of electrons in the atom is called the electron configuration.

• The aufbau principle states that each electron occupies the lowest energy orbital available.





• The aufbau diagram can be used to write correct ground-state electron configurations for all elements up to and including Vanadium, atomic number 23.

• The electron configurations for certain transition metals, like chromium and copper, do not follow the aufbau diagram due to increased stability of half-filled and filled sets of s and d orbitals.


Electron Configuration• Electron Configuration - a representation of the arrangement of electrons

in an atom


• Examples of electron Configuration– 1. Li 1s22s1

– 2. C 1s22s22p6

principleazimuthal

# of e- in that shell


• Take note that after 4s is filled, 3d is than filled before 4p.

• …… 6s than 4f than 5d than 6p• When writing out the electron configuration,

always write your numbers in numerical order– Y 1s22s22p63s23p64s23d104p65s24d1 – NO!

– Y 1s22s22p63s23p63d104s24p64d15s2


• Examples:• Be

• O

• Ca

• Mn


• Examples• Pb

• Os


• Hund’s Rule (better known as the Bus Rule)– Before any second electron can be placed in a sub level, all the

orbitals of that sub level must contain at least one electron – spread out the e- before pairing them up.

• Pauli Exclusion Principle - electrons occupying the same orbital must have opposite spin.

• See a good online illustration at http://www.avogadro.co.uk/light/aufbau/aufbau.htm

http://www.avogadro.co.uk/light/aufbau/aufbau.htm

Orbital Notation

http://www.youtube.com/watch?v=SoNIQjW5Zxs




Orbital Notation

H1s

1s 2s 2p

F

Orbital Notation

• Examples:• Li F

• Na Sc

Valence Electrons

• Valence electrons are defined as electrons in the atom’s outermost orbitals—those associated with the atom’s highest principal energy level.

• Electron-dot structure consists of the element’s symbol representing the nucleus, surrounded by dots representing the element’s valence electrons.


Significance of Electron Configurations

• Valence shell electrons - outermost electrons involved with bonding• no atom has more than 8 valence electrons• Noble gases - 8 valence electrons – least reactive of all elements• Lewis Dot structures: NSEW (cheating) also show correct way, count to 8

Lewis Dot Structures

Essential Questions• How are the Pauli exclusion principle, the aufbau

principle, and Hund’s rule used to write electron configurations?

• How do electron-dot structures represent an atom’s valence electrons?


Documents

Chapter 5 Electrons in Atoms. Section 1 Light and Quantized Energy