 # Chapter 4 Arrangement of Electrons in Atoms

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Chapter 4 Arrangement of Electrons in Atoms

Section 4.1 The Development of a New Atomic Model
Objectives: Discuss the significance of the photoelectric effect and the line-emission spectrum of hydrogen to the development of the atomic model. Describe the Bohr model of the hydrogen atom.

Electromagnetic radiation = energy that exhibits wavelike behavior
All electromagnetic radiation travels at the same speed. Different types are the result of different wavelengths and frequencies. This spectrum shows the different forms of electromagnetic radiation.

Light is a small part of the electromagnetic spectrum.

Absorption/Emission Spectra
When light strikes a metal, the metal ejects e- from the surface and creates an electric current = Photoelectric effect Remember the cathode-ray tube? When the glass tube was filled with a pure gas and an electric current passed through, the gas will gave off light. Different gases give off different colors of light. If this light is passed through a prism, a series of bright lines is seen (emission spectrum). Every element has a distinct emission spectrum.

So, what’s happening to create
the emission spectrum? Carbon Oxygen Iron

Rutherford’s model of the atom provided information about the structure of atoms, it did not explain where the electrons were located in the space surrounding the nucleus.

Max Planck Albert Einstein
In 1900, Max Planck suggested that objects could give off energy in small, specific amounts he called quanta. A quantum (singular of quanta) is the minimum amount of energy that can be lost or gained by an atom. Albert Einstein Einstein proposed that different elements require different frequencies of energy to eject electrons

So scientists agreed… Light is a form of energy Different colors of light have different levels of energy on the electromagnetic spectrum Atoms of different elements had different values for a quantum (the minimum amount of energy they can gain and lose) When atoms of different elements absorb their quantum of energy they can temporarily eject electrons Different colors of light are created by the different levels of energy being absorbed, and then given off, by electrons.

The question still remained:
Why would different elements absorb different amounts of energy and then give off different light colors when energy was applied and they ejected electrons? What was it about their structures that allowed this?

Neils Bohr Model of the H Atom
When an e- is hit by light energy, it absorbs the energy. If the energy is of the correct frequency (quantum), the e- will jump to another energy level (excited state vs ground state). The electron cannot stay in excited state so it falls back to the ground state. It cannot take the energy with it so it releases the energy in the form of LIGHT energy! We see this energy as different colors of light! When the electron returns to ground state, it emits energy in the form of different colors of light – depending on frequency of energy gained.

Section 4.2 The Quantum Model of the Atom
Objectives: Compare and contrast the Bohr model and the quantum model of the atom. List the 4 quantum numbers, and describe their significance. Relate the number of sublevels corresponding to each of an atom’s main energy levels, the number of orbitals per sublevel, and the number of orbitals per main energy level.

Today’s Quantum Model of the Atom
Orbital -- a 3-d region around the nucleus that indicates the probable location of an e- . Quantum theory -- describes mathematically the wave properties of e- and other very small particles Quantum numbers -- numbers or letters that specify the properties of atomic orbitals and the properties of e- in orbitals

4 Quantum Numbers Principal Quantum Number
Angular Momentum Quantum Number Magnetic Quantum Number Spin Quantum Number

Principal quantum number (n) - indicates
the main energy level occupied by the e- n = 1,2,3,4, etc. Representation of Bohr’s proposal of orbitals. Lower numbered levels are closer to nucleus and of lower energy.

Angular Momentum Quantum Number – indicates the orbital and its shape = s, p, d, f
The maximum number of electrons in an energy level is 2n2, thus a shell with n = 2 may hold a maximum of 8 electrons.

Magnetic Quantum Number – orientation of orbital around the nucleus (the axis, or axes, it is on or between); uses x, y, z, xy, yz, etc.

These orbitals can overlap, cause interference with each other, and affect over all energy of each other.

Spin Quantum Number – spin state of an electron in an orbital – can be +½, -½ - often represented with arrows

Review 4 quantum numbers describe each e- .
Principal quantum number - indicates the main energy level occupied by the e- = 1,2,3,4, etc. Angular Momentum Quantum Number – indicates the orbital and its shape = s, p, d, f Magnetic Quantum Number – orientation of orbital around the nucleus (the axis, or axes, it is on or between) = x, y, z, xy, yz, etc. Spin Quantum Number – spin state of an electron in an orbital – can be +½, -½ - often represented with So…to accurately describe the energy (location) of electron, all 4 quantum numbers must be used.

Section 4.3 Electron Configurations
Objectives: List the total number of electrons needed to fully occupy each main energy level State and explain the Aufbau principle, the Pauli exclusion principle, and Hund’s rule. Write orbital notations, electron-configuration notation and noble-gas notation for atoms.

Electron Configuration
Electron configuration -- the arrangement of e- in an atom Electron configurations summarize the locations of each e- in atoms. Like all systems in nature, e- tend to assume arrangements that have the lowest possible energies Ground-state configuration -- lowest energy arrangement of the e- for each element

3 basic rules help govern these ground-state configurations.
Aufbau principle -- an e- occupies the lowest-energy orbital possible Pauli exclusion principle -- no 2 e- in the same atom can have the same set of 4 quantum numbers Hund’s rule -- orbitals of equal energy are each occupied by one e- before any orbital is occupied by a second e- , and all e- in singly occupied orbitals must have the same spin

Types of Electron Configurations
Orbital Notation – includes all 4 quantum numbers (spin is indicated with an arrow) Fluorine – 9 e- _____ _____ _____ _____ _____ 1s s px py pz

Complete Orbital Notations for Elements 1 – 10 on the Periodic Table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon

Types of Electron Configurations
2. Electron Configuration Notation – includes principal and Angular Momentum (combines magnetic and ignores spin) Chromium e- 1s2 2s2 2p6 3s2 3p6 4s2 3d4

Complete Electron Configurations for
Elements 11 – 20 on the Periodic Table Sodium Magnesium Aluminum Si P S Cl Ar K Ca

Types of Electron Configurations
Noble Gas Notation – abbreviates part of the electron configuration by using the Noble Gas Symbol just prior to the element and adds the rest of the electron configuration Magnesium 12e- Mg = [Ne] 3s2

Complete Noble-Gas Notations for
the following elements: S Ca Rb Li O Al

Chapter 4 Study Guide Know the contributions of Plank and Einstein to the Quantum Theory. Be able to define “ground state” and “excited state” and know what happens to cause an atom to moved from ground state to excited state and what happens when it moves back to the ground state. Know the 4 quantum numbers by name and descriptions. Know the 3 rules: Aufbau, Pauli Exclusion, and Hund and know how they are applied when doing configurations. Know the order of energy levels and orbitals from lowest energy to highest energy. Be able to write configurations: orbital notation, electron configuration, and noble gas notations. Be able to use configurations to identify elements. Know the s, p, d, and f blocks on the Periodic Table so you can double check your work.

Class work A quantum of electromagnetic energy is called a(n) .
A bright-light spectrum of an atom is caused by the energy released when electrons . Explain the difference in an atom’s ground state and excited state. Name and describe the information provided by all four quantum numbers. Describe the shapes of the s, p, and d orbitals. How many possible orientations can each of the following sublevels have? p sublevel? d sublevel? f sublevel?

How many electrons can be held in a
p sublevel? d sublevel? f sublevel? How many total electrons can be held in each energy level? 1st – 2nd – 3rd – 4th –

Explain each of the following rules governing
electron configurations. Aufbau principle – Pauli exclusion principle – Hund’s rule – Do the following for an atom of Cesium. Orbital Notation Electron Configuration Noble Gas Notation