# Chapter 5 : Electrons in Atoms. Problems with Rutherford’s Model Chlorine # 17 Reactive Potassium # 19 Very reactive Argon # 18 Not reactive.

## Presentation on theme: "Chapter 5 : Electrons in Atoms. Problems with Rutherford’s Model Chlorine # 17 Reactive Potassium # 19 Very reactive Argon # 18 Not reactive."— Presentation transcript:

Chapter 5 : Electrons in Atoms

Problems with Rutherford’s Model Chlorine # 17 Reactive Potassium # 19 Very reactive Argon # 18 Not reactive

The Quest for a Better Model Electromagnetic radiation behaves like a wave.

Characteristics of a Wave Wavelength = λ Frequency = v (number of waves that pass a point per second) 1 Hertz (Hz) = 1 wave per second (SI Unit for frequency)

Speed and Frequency of Light c = λv c = speed of light (3.0 x 10 8 m/s) ↑ wavelength ↓ frequency ↓ wave length ↑frequency

What is the relationship between energy and frequency?

Problems

Light: Particle or Wave? Wave model doesn’t address: Why heated objects emit only certain frequencies of light at a given temperature? Why some metals emit electrons when a colored light of a specific frequency shines on them?

Iron Dark gray = room temp Red = hot temp Blue = extremely hotter temp ↑ temp, ↑ kinetic energy, emit different colors of light Wave model could not explain this

Max Planck - 1900 Matter gains or loses energy only in small, specific amounts called quanta quantum is the minimum amount of energy that can be gained or lost by an atom E quantum = hv h – Planck’s constant – 6.626 x 10 -34 J·s J = joule, SI Unit for energy

Photoelectric Effect – The problem with wave theory. Only certain frequencies of light could emit an electron from a plate of Ag. Accumulation of low frequencies couldn’t

Einstein and the Dual Nature of EMR (1900) EMR acts as a wave of individual particles (photon) E photon = hv

Calculating the energy in a Photon E photon = hv E = (6.626 x 10-34 J·s) x (7.23 x 1014 s -1 ) E = 4.79 x 10 -19 J

Atomic Emission Spectra The frequencies of the EMR emitted by atoms of the element. Unique to each element

Flame Test Demo

Niels Bohr - 1913 Worked in Rutherford’s lab Proposed a quantum model of the atom Explain why emission spectra were discontinuous Predicted frequencies of light in Hydrogen’s atomic emission spectra

Bohr’s Explanation Ground state – lowest energy state of an atom Excited state – when an atom gains energy Electrons move in circular orbits –Smaller orbit – lower energy state, “energy level” –Larger orbit – higher energy state, “energy level”

An explanation for the Emission Spectra Ground state Excited States Atoms absorb energy and are excited. As the electron returns to the ground state they give off energy “photon” equal to the difference in energy levels.

Problem: Bohr’s Model Only explains Hydrogen Louis de Broglie (1924) – proposed that the energy levels are based on the wave like nature of electrons

Heisenberg Uncertainty Principle It is impossible to know the velocity and position of a particle at any given time Photon and electron are about the same mass.

Erwin Schrodinger - 1926 Developed the quantum mechanical model of the atom –Assigns electrons to energy levels like Bohr –Does not predict the path of the electron –It predicts the probability of finding an electron An electron’s “atomic orbital”

Each dot is a picture of an electron during a given amount of time. Where does the electron spend most of the time? Boundary represents the location of an electron 90% of the time.

Principle Energy Levels 7 energy levels Lowest energy is 1 – greatest energy 7 Each level consists of sublevels The second energy level is larger and the electrons are farther from the nucleus.

Types of Sublevels Same energy

Putting it together: Principal quantum number “energy level” Sublevels (Types of orbitals and total number) 1S – 1 2S – 1, P – 3 3S – 1, P – 3, D – 5 4S – 1, P – 3, D – 5, F - 7

Ground State Electron Configurations Most stable – lowest energy 3 principals to follow

Aufbau Principle – each electron must occupy the lowest energy state 1.Orbitals in an energy sublevel have equal energy 2.The energy sublevels in a principle energy level have different energies. 3.The sublevels increase in energy from s,p,d,f 4.Principal energy levels can overlap

Aufbau Diagram Equal energies – 2 p Sublevels have different energy levels Energy levels overlap

Each electron spins Electrons must spin in opposite directions 2 electrons per orbital Pauli Exclusion Principle Written as

Hund’s Rule Electrons must occupy each orbital before additional electrons can be added. Pauli Exclusion Principle

Representing Electron Configurations

Electron Configurations Sub level diagram – indicates the order that orbitals are filled What are the orbital diagrams and electron configuration notation for Al and Cl?

Electron Configuration Shorthand Substitute noble gases from preceding energy levels in the notation Li – [He] 2s 1 C – [He] 2s 2 2p 2

Valence Electrons Electrons in the outer most energy levels S [Ne] 3s 2 3p 4 Sulfur has 6 valence electrons How many valence electrons do Al, Ne, and Cl have?

Electron Dot Structures Valence electrons are used in reactions and are represented by an electron dot structure.

Writing Electron Dot Structures Fill the valence electrons 1 at a time in any particular order. Ca C O * * * ** * * ** * * * What are the electron dot diagrams for K, Ar and F?

Modern Atomic Theory Any electron in an atom can be described by 4 quantum numbers Principal Quantum Number Azimuthal Quantum Number Magnetic Quantum Number Spin Quantum Number

Principal Quantum Number (n) Related to the size and energy of principal energy level. The farther away from the nucleus the more energy the electron has 1 < 2 < 3 < 4 < 5 < 6 etc….

Azimuthal Quantum Number (Angular Momentum) = l Refers to the subshells in each principal energy level (n) S = 0 P = 1 D = 2 F = 3 n l 10 20 1 30 1 2 40 1 2 3

Magnetic Quantum Number (m l ) Specifies the orbital within a energy level where an electron is likely to be found n l Orbital designation mlml 101s0 202s0 12p-1,0,+1 303s0 13p-1,0,+1 23d-2,-1,0-,1,2 404s0 14p-1,0,+1 24d-2,-1,0-,1,2 34f-3,-2,-1,0,1,2,3

Spin Quantum Number (m s ) + ½ or – ½ Electrons in the same orbitals must have opposite spins (Pauli Exclusion Principle)

n l Orbital designation mlml msms 101s 0 + ½, - ½ 202s 0 + ½, - ½ 12p -1,0,+1 + ½, - ½ 303s 0 + ½, - ½ 13p -1,0,+1 + ½, - ½ 23d -2,-1,0-,1,2 + ½, - ½ 404s 0 + ½, - ½ 14p -1,0,+1 + ½, - ½ 24d -2,-1,0-,1,2 + ½, - ½ 34f -3,-2,-1,0,1,2,3 + ½, - ½

n l mlml msms 21 + ½ or What are the quantum numbers for A? B? AB

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