Electrons in Atoms.

Electrons in Atoms

Atomic Model Timeline

Dalton’s Atomic Theory
John Dalton ( ) had four theories All elements are composed of submicroscopic indivisible particles called atoms Atoms of the same element are identical. The atoms of anyone element are different from those of any other element Atoms of different elements can physically mix together or can chemically combine w/ one another in simple whole-number ratios to form compounds Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another elements as a result of a chemical reaction

Thomson’s Atomic Model
Thomson though electrons were like plums embedded in a positively charged “pudding”, so his model was called the “plum pudding” model

Thomson’s Theory Thomson stated: The atom had negatively charged electrons stuck into a lump of positively charged protons. Thomson never explained Number of protons and neutrons The arrangement of the particles in the atom The ease with which atoms are stripped of electrons to form ions

Rutherford Model Rutherford used the Gold Foil Experiment
Rutherford proposed the following: Thomson model was incorrect Most of the mass of the atom and all of its positive charge reside in a very small, extremely dense region, which he called the nucleus Most of the total volume of the atom is empty space in which electrons move around the nucleus

Rutherford’s Model Discovered dense positive piece at the center of the atom Nucleus Electrons moved around Mostly empty space

Bohr Model Bohr changed the Rutherford model and explained how the electrons travel. In 1913, Bohr introduced his atomic model based on the simplest atom, hydrogen (only 1 electron) Bohr explained the following in his model: Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus. Electrons don’t fall into the nucleus because electrons in particular path have fixed energy and don’t lose energy His model was patterned after the motion of the planets around the sun. It is often called the Planetary model.

IMPLICATIONS Classical Theory: (old)
Any amount of energy could be absorbed or emitted Quantum Theory: (new) Hydrogen atoms can only have EXACT energy levels

Bohr Model Cont.

Bohr’s Model Nucleus Electron Orbit Energy Levels

The Bohr Model Each electron has a fixed energy = an energy level.
Electrons can jump from one energy level to another. Electrons can not be or exist between energy levels. A quantum of energy is the amount of energy needed to move an electron from one energy level to another energy level.

Bohr Model To move from one level to another, the electron must gain or lose the right amount of energy. The higher the energy level, the farther it is from the nucleus. Gain energy to move to higher energy levels (away from nucleus) Lose energy to move to lower energy levels (closer to nucleus)

The Bohr Model The amount of energy required to go from one energy level to another is the not same for the electrons. Higher energy levels are closer together. This means it takes less energy to change levels in the higher energy levels. The Bohr model was tested with the hydrogen element but failed to explain the energies absorbed and emitted by atoms with more than one electron.

} Bohr’s Model Fifth Fourth Increasing energy Third Second First
Further away from the nucleus means more energy. There is no “in between” energy Energy Levels Fifth Fourth Third Increasing energy Second First Nucleus

Bohr’s Model cont. Energy levels are not equally spaced.
Energy levels more closely spaced further from the nucleus Higher energy level occupied by an electron, the more energetic that electron is. Amount of energy gained or lost by an electron is not always the same amount.

Bohr Model

An Explanation of Atomic Spectra
5.3 An Explanation of Atomic Spectra In the Bohr model, the lone electron in the hydrogen atom can have only certain specific energies. When the electron has its lowest possible energy, the atom is in its ground state. Excitation of the electron by absorbing energy raises the atom from the ground state to an excited state. A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level.

5.3 Atomic Spectra 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.

An Explanation of Atomic Spectra
5.3 An Explanation of Atomic Spectra 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.

The degree to which they move from level to level determines the frequency of light they give off.
Figure 5-14

Did you know that an element can be identified by its emission spectra?
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. Mercury Nitrogen

Fingerprints of certain atoms
Figure 5-8

Quantum Theory Bohr explained how electrons were moving via Quantum Theory Key Terms: Energy Levels- Regions around the nucleus where the electron is likely moving Quantum- Amount of energy required to move an electron from one energy level to the next Quantum Leap- Abrupt Change

Limitations of the Bohr Model
Can only treat atoms or ions with one electron Does not predict the intensities of the spectral lines Inconsistent with the uncertainty principle (see later!) Does not predict the observed splitting of the spectral lines

From Bohr model to Quantum mechanics
Bohr’s theory was a great accomplishment and radically changed our view of matter. But problems existed with Bohr theory — theory only successful for the H atom. introduced quantum idea artificially. So, we go on to QUANTUM or WAVE MECHANICS

Quantum or Wave Mechanics
Light has both wave & particle properties de Broglie (1924) proposed that all moving objects have wave properties. For light: E = hn = hc / l For particles: E = mc2 (Einstein) L. de Broglie ( ) Therefore, mc = h / l and for particles (mass)x(velocity) = h / l l for particles is called the de Broglie wavelength

Louis de Broglie ( ) Electrons should be considered waves confined to the space around an atomic nucleus

The Quantum Mechanical Model
Rutherford’s and Bohr’s model focused on describing the path of the electron around the nucleus like a particle (like a small baseball). Austrian physicist Erwin Schrödinger (1887–1961) treated the electron as a wave. The modern description of the electrons in atoms, the quantum mechanical model, comes from the mathematical solutions to the Schrödinger equation.

I don’t believe that God plays dice!
Erwin Schrödinger Schrödinger introduced the wave function, a function of position and time whose absolute value squared is related to the probability of finding an electron near a specific point in space and time. I don’t believe that God plays dice!

Erwin Schrödinger In this theory, the electron can be thought of as being spread out over a large volume and there are places where it is more likely to be found than others! This can be thought of as an electron cloud. Rubbish!

I’m used to the idea of waves, so I like using Schrödinger’s model
Beware! This wave function is only a mathematical model that fits very well. It also links well with the idea of wave particle duality (electron as wave and particle). But it is only one mathematical model of the atom. Other more elegant mathematical models exist that don’t refer to waves, but physicists like using the wave model because they are familiar with waves and their equations. We stick with what we are familiar! I’m used to the idea of waves, so I like using Schrödinger’s model .

Erwin Schrödinger (1926) Quantum mechanics
electrons can only exist in specified energy states Electron cloud model orbital: region around the nucleus where e- are likely to be found

Schrödinger’s Equation
The wave function is a F(x, y, z) Actually F(r,θ,φ) Solutions to the equation are called orbitals. These are not Bohr orbits. Each solution is tied to a certain energy These are the energy levels Animation

Quantum or Wave Mechanics
Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms. Solution to WAVE EQUATION gives set of mathematical expressions called WAVE FUNCTIONS, Y Each describes an allowed energy state of an e- Quantization introduced naturally. E. Schrodinger

WAVE FUNCTIONS, Y • Y is a function of distance and two angles.
• For 1 electron, Y corresponds to an ORBITAL — the region of space within which an electron is found. • Y does NOT describe the exact location of the electron. • Y2 is proportional to the probability of finding an e- at a given point.

Uncertainty Principle
Problem of defining nature of electrons in atoms solved by W. Heisenberg. Cannot simultaneously define the position and momentum (= m•v) of an electron. Dx. Dp = h At best we can describe the position and velocity of an electron by a PROBABILITY DISTRIBUTION, which is given by Y2 W. Heisenberg

Werner Heisenberg Heisenberg Uncertainty Principle – It is impossible to determine the position and velocity of an electron simultaneously.

Heisenberg uncertainty principle
In order to observe an electron, one would need to hit it with photons having a very short wavelength. Short wavelength photons would have a high frequency and a great deal of energy.

Heisenberg uncertainty principle
If one were to hit an electron, it would cause the motion and the speed of the electron to change. Lower energy photons would have a smaller effect but would not give precise information.

Electrons as Waves EVIDENCE: DIFFRACTION PATTERNS VISIBLE LIGHT

5.1 The Quantum Mechanical Model The propeller blade has the same probability of being anywhere in the blurry region, but you cannot tell its location at any instant. The electron cloud of an atom can be compared to a spinning airplane propeller. The quantum model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus. The electron cloud of an atom is compared here to photographs of a spinning airplane propeller. a) The airplane propeller is somewhere in the blurry region it produces in this picture, but the picture does not tell you its exact position at any instant. b) Similarly, the electron cloud of an atom represents the locations where an electron is likely to be found.

Things that are very small behave differently from things big enough to see. The quantum mechanical model is a mathematical solution It is not like anything you can see.

5.1 The Quantum Mechanical Model The probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud. The cloud is more dense where the probability of finding the electron is high. The electron cloud of an atom is compared here to photographs of a spinning airplane propeller. a) The airplane propeller is somewhere in the blurry region it produces in this picture, but the picture does not tell you its exact position at any instant. b) Similarly, the electron cloud of an atom represents the locations where an electron is likely to be found.

5.1 Atomic Orbitals (fuzzy cloud) = An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron.

Radial Distribution Curve
Quantum Mechanics Orbital (“electron cloud”) Region in space where there is 90% probability of finding an e- Orbital Radial Distribution Curve

The atom is found inside a blurry “electron cloud” A area where there is a chance of finding an electron. Draw a line at 90 %

Probability cloud

Atomic Orbitals There are the region of space which there is a high probability of finding an electron Within each energy level the complex math of Schrodinger’s equation describes several shapes. These are called atomic orbitals Quantum Numbers- numbers that specify the properties of atomic orbitals and their electrons

Quantum Numbers There are 4 types of Quantum Numbers
Principal – distance from the nucleus Angular Momentum- Orbital Shape Magnetic- Orbital position with respect to the X, Y, & Z axes. Spin- Has only two values (+1/2 or –1/2) and is needed to specify 1 of 2 positional orientations of an electron

Principal Quantum Number
Symbolized by the letter N, indicates the main energy levels surrounding the nucleus There are 7 principal quantum numbers A.K.A. – Shells Value of N is a whole number ex. 1,2,3 ect.. Main Energy Level – N=1; closest to the nucleus or ground state Ground State- state of the lowest energy of the atom. As N increases, the distance from the nucleus increases and the energy increases

Angular Momentum Quantum Number
Indicates the shape of the orbital. Within each main energy level beyond the first, orbitals with different shapes occupy different regions A.K.A. – Sublevels or Subshells The number of sublevels = Value of the Principal Quantum Number

Magnetic Quantum Number
Indicates the orientation of a orbital about the nucleus There are 4 types of orbital orientation S Orbital P Orbital D Orbital F Orbital

Spin Quantum Number Has only two possible values: +1/2 or –1/2. These values indicate two possible states of an electron in an orbital Spin Quantum # is significant because each single orbital can hold no more than two electrons, which must have opposite spin.

Larger atom— More electrons take up more space. s-orbitals are spherically shaped. Smaller atom Smaller atom— Fewer electrons take up less space.

S orbitals 1 s orbital for every energy level Spherical shaped
Each s orbital can hold 2 electrons Called the 1s, 2s, 3s, etc.. orbitals.

p-orbitals are “dumbell” shaped. z-axis

p-orbitals are “dumbell” shaped. x-axis

p-orbitals are “dumbell” shaped. y-axis

p-orbitals together x, y, & z axes.

P orbitals Start at the second energy level 3 different directions
3 different shapes Each can hold 2 electrons

D orbitals Start at the third energy level 5 different shapes
Each can hold 2 electrons

F orbitals Start at the fourth energy level
Have seven different shapes 2 electrons per shape

F orbitals

Table 5-1

Shells and Orbitals and Atomic Structure
f d p s Shells of an atom contain a number of stacked orbitals 4 3 2 1

Atomic Orbitals 5.1 Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped. Four of the five d orbitals have the same shape but different orientations in space. The d orbitals are illustrated here. Four of the five d orbitals have the same shape but different orientations in space. Interpreting Diagrams How are the orientations of the dxy and dx2 – y2 orbitals similar? How are they different?

Summary # of shapes Max electrons Starts at energy level s 1 2 1 p 3 6
d 5 10 3 7 14 4 f

5.1 Atomic Orbitals The numbers and kinds of atomic orbitals depend on the energy sublevel. Energy Level, n # of sublevels Letter of sublevels # of orbitals per sublevel # of electrons in each orbital Total electrons in energy level

5.1 Atomic Orbitals The numbers and kinds of atomic orbitals depend on the energy sublevel. Energy Level, n # of sublevels Letter of sublevels # of orbitals per sublevel # of electrons in each orbital Total electrons in energy level 1 s 2 s p 3 6 8 s p d 5 10 18 4 s p d f 7 14 32

By Energy Level First Energy Level only s orbital only 2 electrons 1s2
Second Energy Level s and p orbitals are available 2 in s, 6 in p 2s22p6 8 total electrons

By Energy Level Third energy level s, p, and d orbitals
2 in s, 6 in p, and 10 in d 3s23p63d10 18 total electrons Fourth energy level s,p,d, and f orbitals 2 in s, 6 in p, 10 in d, ahd 14 in f 4s24p64d104f14 32 total electrons

By Energy Level Any more than the fourth and not all the orbitals will fill up. You simply run out of electrons The orbitals do not fill up in a neat order. The energy levels overlap Lowest energy fill first.

Question for You How many principal quantum numbers are there?
What is the maximum number of electrons that can fill the 3rd energy level? How many orbitals are in the sublevel F? What is the total number of orbitals for the 3rd main energy level?

Electron Configuration
The way electrons are arranged in atoms There are three rules which help dictate how electrons are arranged in the atoms. Aufbau Principle- electrons occupy the orbitals of the lowest energy first Hund’s Rule- Orbitals of equal energy are each occupied by one electron before any one orbital is occupied by a second electron. All electrons in a single occupied orbital must have the same spin.

Electron Configuration cont.
Pauli Exclusion Principle- No two electrons may occupy any given orbital without having opposite spin. No two electrons in the same atom can have the same set of four quantum numbers. Let’s determine electron configuration. Let’s start with Phosphorus. Need to account for all 15 electrons

Electron Configurations
Distribution of all electrons in an atom Consist of Number denoting the energy level

Distribution of all electrons in an atom Consist of Number denoting the energy level Letter denoting the type of orbital

Distribution of all electrons in an atom. Consist of Number denoting the energy level. Letter denoting the type of orbital. Superscript denoting the number of electrons in those orbitals.

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

Orbital Diagrams Each box represents one orbital.
Half-arrows represent the electrons. The direction of the arrow represents the spin of the electron.

Hund’s Rule “For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”

Exceptions to Electron Configuration

Orbitals fill in order Lowest energy to higher energy.
Adding electrons can change the energy of the orbital. Half filled orbitals have a lower energy. Makes them more stable. Changes the filling order

Write these electron configurations
Titanium - 22 electrons 1s22s22p63s23p64s23d2 Vanadium - 23 electrons 1s22s22p63s23p64s23d3 Chromium - 24 electrons 1s22s22p63s23p64s23d4 is expected But this is wrong!!

Chromium is actually 1s22s22p63s23p64s13d5 Why?
This gives us two half filled orbitals. Slightly lower in energy. The same principal applies to copper.

Copper’s electron configuration
Copper has 29 electrons so we expect 1s22s22p63s23p64s23d9 But the actual configuration is 1s22s22p63s23p64s13d10 This gives one filled orbital and one half filled orbital. Remember these exceptions

Shortcuts for Electron Configuration
There are two short handed methods of writing the electron configuration. The 1st method is called the outer-level configuration. That tells you the outer-most configuration for that element. The 2nd method is called the Noble Gas Notation. This tells you the complete notation using Noble Gases. Let’s start with outer-level notation!!!

S- block s1 s2 Alkali metals all end in s1
Alkaline earth metals all end in s2 really have to include He but it fits better later. He has the properties of the noble gases.

The P-block p1 p2 p3 p4 p6 p5

Transition Metals -d block

F - block f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13
inner transition elements f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13

Summary Outer-Level Configuration

Figure 7.23

Writing Electron configurations the easy way
Yes there is a shorthand

Electron Configurations repeat
The shape of the periodic table is a representation of this repetition. When we get to the end of the column the outermost energy level is full. This is the basis for our shorthand.

The Shorthand Write the symbol of the noble gas before the element.
Then the rest of the electrons. Aluminum - full configuration. 1s22s22p63s23p1 Ne is 1s22s22p6 so Al is [Ne] 3s23p1

More examples Ge = 1s22s22p63s23p64s23d104p2 Ge = [Ar] 4s23d104p2
Hf=1s22s22p63s23p64s23d104p65s2 4d105p66s24f145d2 Hf=[Xe]6s24f145d2

The Shorthand Again Sn- 50 electrons The noble gas before it is Kr
Takes care of 36 Next 5s2 Then 4d10 Finally 5p2 [ Kr ] 5s2 4d10 5p2

Quantum Numbers Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies. Each orbital describes a spatial distribution of electron density. An orbital is described by a set of three quantum numbers.

The Wave-like Electron
The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves. Louis deBroglie

Heisenberg Uncertainty Principle
It is impossible to know exactly the position and velocity (momentum) of a particle. The better we know one, the less we know the other. The act of measuring changes the properties. More precisely the velocity is measured, less precise is the position (vice versa).

A totally new approach De Broglie said matter could be like a wave. De Broglie said they were like standing waves. The vibrations of a stringed instrument

What’s possible? You can only have a standing wave if you have complete waves. There are only certain allowed waves. In the atom there are certain allowed waves called electrons. 1925 Erwin Schroedinger described the wave function of the electron Much math, but what is important are the solutions

Schrödinger’s Equation
The wave function is a F(x, y, z) Actually F(r,θ,φ) Solutions to the equation are called orbitals. These are not Bohr orbits. Each solution is tied to a certain energy These are the energy levels Animation

What does the wave Function mean?
nothing. it is not possible to visually map it. The square of the function is the probability of finding an electron near a particular spot. best way to visualize it is by mapping the places where the electron is likely to be found.

Probability Distance from nucleus

Sum of all Probabilities
Distance from nucleus

Defining the size The nodal surface.
The size that encloses 90% to the total electron probability. NOT at a certain distance, but a most likely distance. For the first solution it is a a sphere.

Quantum Mechanics Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. It is known as quantum mechanics. © 2009, Prentice-Hall, Inc.

Quantum Mechanics The wave equation is designated with a lower case Greek psi (). The square of the wave equation, 2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time. © 2009, Prentice-Hall, Inc.

Quantum Numbers Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies. Each orbital describes a spatial distribution of electron density. An orbital is described by a set of three quantum numbers. © 2009, Prentice-Hall, Inc.

Quantum Numbers There are many solutions to Schrödinger’s equation
Each solution can be described with quantum numbers that describe some aspect of the solution. Principal quantum number (n) size and energy of an orbital Has integer values >0

s Orbitals Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n−1 nodes, or regions where there is 0 probability of finding an electron. © 2009, Prentice-Hall, Inc.

p Orbitals The value of l for p orbitals is 1.
They have two lobes with a node between them. © 2009, Prentice-Hall, Inc.

d Orbitals The value of l for a d orbital is 2.
Four of the five d orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center. © 2009, Prentice-Hall, Inc.

Quantum numbers Angular momentum quantum number l shape of the orbital
integer values from 0 to n-1 l = 0 is called s l = 1 is called p l =2 is called d l =3 is called f l =4 is called g

Magnetic Quantum Number (ml)
The magnetic quantum number describes the three-dimensional orientation of the orbital. Allowed values of ml are integers ranging from -l to l: −l ≤ ml ≤ l. Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc. © 2009, Prentice-Hall, Inc.

Magnetic Quantum Number (ml)
Orbitals with the same value of n form a shell. Different orbital types within a shell are subshells. © 2009, Prentice-Hall, Inc.

Pauli Exclusion Principle
No two electrons in the same atom can have exactly the same energy. Therefore, no two electrons in the same atom can have identical sets of quantum numbers. © 2009, Prentice-Hall, Inc.

Electrons in Atoms.

Similar presentations

Presentation on theme: "Electrons in Atoms."— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

Electrons in Atoms.

Similar presentations

Presentation on theme: "Electrons in Atoms."— Presentation transcript:

Similar presentations

About project

Feedback