electron configuration chemistry. chapter 4 – section 2 the quantum model of the atom

Electron ConfigurationChemistry

Chapter 4 – Section 2

• The Quantum Model of the Atom

I. Introduction• Bohr’s Model did not make sense

to most scientists, who felt that the electron should exist at any distance from the nucleus, depending on energy, rather than in certain levels

II. Electrons as Waves• A. De Broglie proposed that an

electron, like light, might exist as a particle and a wave

• 1. A wave confined to a certain space can only have certain frequencies

• a. This corresponded to Bohr’s orbits• b. It also corresponded to the specific

frequencies produced in line-emission spectrums

• 2. Beams of electrons can also be bent or refracted, or experience interference, just as light does

• a. Diffraction – the bending of a wave around the edge of an object or opening

• b. Interference – the overlapping of waves, which either increases or decreases each the waves’ energy

III. Heisenberg Uncertainty Principle

• A. The dual wave-particle nature of an electron made many scientists uneasy, because they could not pin point where the electron would be located in the atom

• B. Heisenberg proposed that electrons and photons have about the same energy

• 1. Since photons are used to knock electrons off metal (and thus, detect the electrons), the exact location of an electron cannot be determined with certainty

• 2. The Heisenberg Uncertainty Principle – it is impossible to determine the location and speed of the electron or photon at the same time

IV. The Schrodinger Wave Equation

• A. Schrodinger developed an equation to show that electrons behave as waves (only certain frequencies would solved the equation)

• B. Together, Heisenberg and Schrodinger laid the foundation for our modern quantum theory – the mathematical description of the wave properties of electrons and other small particles

• 1.Solutions to the equations are called wave functions

• 2.They do not locate the electron, but merely give the probability of locating it

• 3.This showed that electrons do not travel in neat orbits around the nucleus as Bohr described, but exist in certain regions called orbitals

• 4.Orbital – a three-dimensional region around the nucleus that indicates the probable location of the electron

• 5.Orbitals can have • different shapes and sizes

V. Atomic Orbitals and Quantum Numbers• A. Quantum numbers are used to describe

orbitals and the properties of electrons in them

• 1.The first three quantum numbers indicate:• a.Main energy level of the orbital• b.Shape of orbital• c.Orientation of orbital

• 2.The fourth quantum number is the spin quantum # and indicates the fundamental spin state of the electron

• B.Principal Quantum Number (n)• 1.This is the main energy level of the electron• 2.As (n) increases, electron energy and distance

from the nucleus increases• 3.Electrons with the same (n) value will be in the

same energy level• 4.(n2) is the total number of • orbitals in each energy level

• C.Angular Momentum Quantum Number• 1.These sublevels are orbitals of different shapes• 2.(l) indicates the shape of the orbital and is

called the (l) value• 3.If the (l) value is s, then the shape is spherical• 4.If the (l) value is p, then the shape is a dumbbell• 5.If the (l) value is d, then the shape is a clover

leaf• 6.If the (l) value is f, then the shape is a flower

petals

p d f

D. Magnetic Quantum Number• 1.Orbitals can be the same shape, but have

different orientations (axes)• 2.Magnetic Quantum Number (m) – indicates the

orientation (or axis) of the orbital• 3.The (m) value = + l to -l

E. Spin Quantum Number• 1.These values can be +1/2 or –1/2• 2.This indicates the spin state of the

electron• 3.Any orbital can hold 2 electrons, but they

must have opposite spin states

Chapter 4 – Section III

•Electron Configurationelectron

neutron

proton

I. Electron Configuration• A. Bohr’s model of the atom only

described electron arrangement in the Hydrogen atom

• B.The quantum model of the atom describes electron arrangement for all atoms

• 1. Electron Configuration – the arrangement of electrons in an atom

• 2. Electrons always assume the lowest possible energy state

II. Rules Governing Electron Configurations

• A. Aufbau principle – an electron occupies the lowest energy orbital that has a space for it

• B. Pauli exclusion principle – no two electrons in the same atom can have the same 4 quantum numbers

• 1. No two electrons can be in the same place at the same time

• 2. If two electrons occupy the same orbital, then they must have opposite spin states

• C. Hund’s rule – orbitals of equal energy will have one electron each until all orbitals have one before a second electron enters an occupied orbital

• 1. This minimizes repulsion between electrons

• 2. This allows the electron to have the lowest energy possible electron

neutron

proton

III. Representing Electron Configurations

• A. There are 3 methods of electron configuration notation

• 1. Two are used for elements in the 1st two periods

• 2. The third is used for elements in the 3rd period and higher

1s2 2s2 2p6 3s2 3p5

OR [Ne]3s2 3p5

1s2 2s2 2p6 3s2 3p6 3d10 4s1

OR [Ar]3d10 4s1

• B. Orbital Notation • 1. A blank line (___) represents an orbital

with no electrons• 2. An orbital with 1 electron is represented

with___• An orbital with its maximum number of two

electrons is represented with _____(shows the opposite spins)

• 4. Each line is labeled with the (n) and (l) value

Carbon is paramagnetic

Neon is diamagnetic

• C. Electron Configuration Notation – each orbital letter has a superscript to show how many electrons are in the level

• D. Elements of the 2nd Period• 1. Highest occupied energy level – the

energy level farthest from the nucleus that contains electrons

• 2. Inner shell electrons – those electrons not in the highest energy level

• 3. Octet of electrons – highest energy level is filled with 8 electrons

• E. Elements of the 3rd Period• 1. 1st 10 electrons of any atom in the period are the

same as Neon• 2. Noble Gases – the Group VIII (or 18) elements• 3. Noble Gas Notation – also called abbreviated

notation• a.Use the symbol for the noble gas in brackets [Ne]

that is previous to the element• b.Then continue with the electron configuration

notation• 4. Noble Gas Configuration – the outermost

energy level containing 8 electrons (Helium is the exception)

• F. Elements of the 4th Period• 1. 4s fills 1st

• 2. 3d fills 2nd

• 3. 4p fills 3rd

• G. 5th Period Elements• 1. 5s fills 1st

• 2. 4d fills 2nd

• 3. 5p fills 3rd

• H. Elements of the 6th and 7th Periods• 1. 6s fills 1st

• 2. 5d fills 2nd in La• 3. 4f fills 3rd

• 4. The rest of 5d fills 4th

• 5. 6p fills 5th

• 6. 7s fills 6th

• 7. 6d fills 7th in Ac• 8. 5f fills 8th

• 9. Rest of 6d fills 9th

1s

2s 2px 2py 2pz

3s 3px 3py 3pz

4s 4px 4py 4pz

3dxy 3dxz 3dyz 3dx2y2 3dz2

4dxy 4dxz 4dyz 4dx2y2 4dz2

4fz3 4fxz2 4yz2 4fy(3x2-y2)

4fx(x2-3y2) 4fz(x2-y2) 4fxyz

• 1s2 2s2 2p6 3s2 3p5• OR• [Ne]3s2 3p5