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Chapter 5: Electrons in Atoms
Transcript of Chapter 5: Electrons in Atoms
Energy Levels are the fixed energies an electron can have.
A Quantum of energy is the amount of energy required to move an electron from one energy level to another energy level.
The Quantum Mechanical Theory is the modern description of the electrons in atoms and comes from the mathematical solutions to the Schrödinger equation.
An Atomic Orbital is often thought of as a region of space in which there is a high probability of finding an electron. He pounded up materials in his pestle and mortar until he had reduced them to smaller and smaller particles which he called... ATOMA
(Greek for indivisible) Democritus Developed the Idea of Atoms 460 BC John Dalton 1808 Suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them... ATOMS Found that atoms could sometimes release a far smaller negative particle which he called an... Joseph John Thompson 1898 ELECTRON PLUM PUDDING
MODEL like plums surrounded by pudding. 1904 Thompson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge... He fired Helium nuclei at a piece of gold foil which was only a few atoms thick.
He found that although most of them passed through. About 1 in 10,000 hit Ernest Rutherford 1910 He found that while most of the helium nuclei passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back. Rutherford’s new evidence allowed him to propose a more detailed model with a central nucleus.
He suggested that the positive charge was all in a central nucleus. With this holding the electrons in place by electrical attraction
However, this was not the end of the story. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons. Studied under Rutherford Niels Bohr 1913 The electron cloud model is an atom model wherein electrons are no longer depicted as particles moving around the nucleus in a fixed orbit. Instead, as a quantum mechanically-influenced model, we shouldn’t know exactly where they are, and hence describe their probable location around the nucleus only as an arbitrary "cloud". Electron Cloud Model Discovered By Erwin Schrodinger 1926
Electrons can jump from energy level to energy level.
Electrons absorb or emit light energy when they jump from one energy level to another known as...
The energy levels are like the rungs of a ladder but are not equally spaced. Basics of the Quantum Mechanical Model Photons are bundles of light energy that are emitted by electrons as they go from higher energy levels to lower levels. Photons Ground state: the lowest possible energy level an electron be at.
Excited state: an energy level higher than the ground state. Excited State and Ground State Four quantum numbers are required to describe the state of the hydrogen atom. A region in space in which there is high probability of finding an electron. Atomic Orbital: Specify the properties of atomic orbitals and their electrons. Quantum Numbers: Principal Quantum Number
Orbital Quantum Number
Magnetic Quantum Number
Spin Quantum Number Four Quantum Numbers Indicates main energy levels
n = 1, 2, 3, 4…
Each main energy level has sub-levels
The maximum number of electrons in a principal energy level is given by:
Max # electrons = 2n^2
n= the principal quantum number Principal Quantum Number (n)
ℓℓℓℓℓ l = n-1
ℓ l sublevel
4 g Orbital Quantum Number (ℓ)
(Angular Momentum Quantum Number) Atomic Orbital (s) x 1 Atomic Orbitals (p) x 3 Atomic Orbitals (d) x 5 Atomic orbitals (f) x 7 Indicates the orientation of the orbital in space.
Values of ml : integers -l to l
The number of values represents the number of orbitals.
for l= 2, ml = -2, -1, 0, +1, +2 Which sub-level does this represent?
Answer: d Magnetic Quantum Number (ml) Indicates the spin of the electron (clockwise or counterclockwise).
Values of ms: +1/2, -1/2 Electron Spin Quantum Number (ms or s) List the values of the four quantum numbers for orbitals in the 3d sub-level.
l = 2
ml = -2,-1, 0, +1, +2
ms = +1/2, -1/2 for each pair of electrons Example: The higher the electron density, the higher the probability that an electron may be found in that region.
The electron cloud represents positions where there is probability of finding an electron. The Electron Cloud Electrons are located in specific energy levels.
There is no exact path around the nucleus.
The model estimates the probability of finding an electron in a certain position. Quantum Mechanical Model 2s Indicates shape of orbital sublevels n = 4 n = 2 n = 1 n = 3 n = 0 Key Concepts
•Rutherford’s atomic model could not explain the chemical properties of elements.
•Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.
•The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.
•Each energy sub-level corresponds to an orbital of a different shape, which describes where the electron is likely to be found. Electron Configurations are the ways in which electrons are arranged in various orbitals around the nuclei of atoms.
The Aufbau Principle states that electrons occupy the orbitals of lowest energy first.
The Pauli Exclusion Principle states that an atomic orbital may describe at most two electrons.
Hund’s Rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. •Three Rules to electron configuration, The Aufbau Principle, The Pauli Exclusion Principle, and the Hund’s Rule, tell you how to find the electron configurations of atoms.
•Some actual electron configurations differ from those assigned using the Aufbau Principle because half-filled sub-levels are not as stable as filled sub-levels, but they are more stable than other configurations. Vocabulary Key Concepts Electron Configurations Electron configurations tells us in which orbitals the electrons for an element are located.
1. Electrons fill orbitals starting with lowest n and moving upwards (Aufbau);
2. No more than two electrons can fill one orbital with the same spin (Pauli);
3. For degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron (Hand’s rule). Aufbau Principle The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel.
The sublevel is written followed by a superscript with the number of electrons in the sublevel.
If the 2p sublevel contains 2 electrons, it is written 2p2 First, determine how many electrons are in the atom.
For Example Iron has 26 electrons.
Arrange the energy sublevels according to increasing energy:
1s 2s 2p 3s 3p 4s 3d …
Fill each sublevel with electrons until you have used all the electrons in the atom:
Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6
The sum of the superscripts equals the atomic number of... Iron (26) Writing Electron Configurations Pauli Exclusion Principle Hund's Rule Electron Configurations and the Periodic Table The periodic table can be used as a guide for electron configurations.
The period number is the value of n.
Groups 1A and 2A have the s-orbital filled.
Groups 3A - 8A have the p-orbital filled.
Groups 3B - 2B have the d-orbital filled.
The Lanthanides and Actinides have the f-orbital filled. Recall, the electron configuration for Na is:
Na: 1s2 2s2 2p6 3s1
We can abbreviate the electron configuration by indicating the innermost electrons with the symbol of the preceding noble gas.
The preceding noble gas with an atomic number less than sodium is neon, Ne. We rewrite the electron configuration:
Na: [Ne] 3s1 Noble Gas Core Electron Configurations This is a Condensed Electron Configuration
Neon completes the 2p subshell.
Sodium marks the beginning of a new row.
So, we write the condensed electron configuration for sodium as
Na: [Ne] 3s1
[Ne] represents the electron configuration of neon.
Core electrons: electrons in [Noble Gas].
Valence electrons: electrons outside of [Noble Gas]. Exceptions to the Aufbau Principle What it should be Cr [Ar] 4s2 3d4
Cu [Ar] 4s2 3d9 What it actually is Cr [Ar] 4s1 3d5
Cu [Ar] 4s1 3d10 This happens because the d orbitals are unstable and want to be stable, so they take an electron from the s orbital.
The Amplitude of a wave is the wave’s height from zero to the crest (peak).
The Wavelength, represented by the Greek letter lambda, is the distance between the peaks.
Frequency, represented by the Greek letter nu, is the number of wave cycles to pass a given point per unit of time.
Hertz (Hz) is the SI unit of cycles per second.
Electromagnetic Radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet rays, X-rays, and gamma rays.
A Spectrum, specifically the electromagnetic spectrum, shows the separation of different frequencies including color.
The Atomic Emission Spectrum is a certain spectrum that determines the frequencies emitted by an element into discrete lines.
The Ground State of an atom is when that atom is at its lowest energy level.
Photons are light quanta.
The Heisenberg Uncertainty Principle states that it is impossible to know exactly both the velocity and the position of a particle at the same time. •Frequency x Wavelength = Speed of Light (nu x lambda = c)
•The wavelength and frequency of light are inversely proportional to each other.
•When atoms absorb energy, electrons move into higher energy levels. These electrons then lose energy by emitting light when they return to lower energy levels.
•The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron.
•Classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms as waves. Vocabulary Key Concepts The study of light led to the development of the quantum mechanical model.
Light is a kind of electromagnetic radiation.
Electromagnetic radiation includes many kinds of waves
All move at 3.00 x 108 m/s = c Light Parts of a wave Crest To measure where a electron is, we use light.
But the light moves the electron
And hitting the electron changes the frequency of the light. More obvious with the very small It is impossible to know exactly the location and velocity of a particle.
The better we know one, the less we know the other.
Measuring changes the properties.
Instead, analyze interactions with other particles Heisenberg Uncertainty Principle Quantum mechanics explains how the very small behaves.
Classic physics is what you get when you add up the effects of millions of packages.
Quantum mechanics is based on probability The physics of the very small Energy is quantized
Light is energy
Light must be quantized
These smallest pieces of light are called photons.
Energy & frequency: directly related. Light is a Particle Are inversely related
As one goes up the other goes down.
Different frequencies of light are different colors of light.
There is a wide variety of frequencies
The whole range is called a spectrum Frequency and wavelength Origin - the base line of the energy.
Crest - high point on a wave
Trough - Low point on a wave
Amplitude - distance from origin to crest
Wavelength - distance from crest to crest
Is abbreviated by the Greek letter lambda = l
Frequency is the number of waves that pass a given point per second.
Units: cycles/sec or hertz (hz or sec-1) abbreviated by Greek letter nu = n Let’s look at a hydrogen atom Changing the energy When we write electron configurations, we are writing the lowest energy.
The energy level, and where the electron starts from, is called it’s ground state- the lowest energy level. Explanation of Atomic Spectra White light is made up of all the colors of the visible spectrum.
Passing it through a prism separates it. Spectra/Prisms These are called discontinuous spectra, or line spectra
They are unique to each element. May fall down in steps, each with a different energy As the electron falls back to ground state, it gives the energy back as light Heat or electricity or light can move the electron up energy levels (“excited”) Infrared Visible Ultraviolet Further they fall, more energy, higher frequency.
This is simplified
the orbitals also have different energies inside energy levels
All the electrons can move around. By heating a gas with electricity we can get it to give off colors.
Passing this light through a prism does something different. If the light is not white... Each element gives off its own characteristic colors.
Can be used to identify the atom.
How we know what stars are made of. Atomic Spectrum After Photon changes wavelength Electron Changes Velocity Before Photon Moving Electron c = ln Increasing Energy Level Thank You and Chemistry On!!!