Topic 2.2: Electrons Honors Chemistry Mrs. Peters 1

2.2: Electron Configuration Essential Idea: The electron configuration of an atom can be deduced from its atomic number. Nature of Science: Developments in scientific research follow improvements in apparatus – the use of electricity and magnetism in Thomson’s cathode rays. (1.8) Theories being superseded – quantum mechanics is among the most current models of the atom (1.9) Use theories to explain natural phenomena – line spectra explained by the Bohr model of the atom (2.2) 2

2.2: Electron Configuration Understandings: Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level. The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies. The main energy level or shell is given an integer number, n, and can hold a maximum number of electrons, 2n 2. 3

2.2: Electron Configuration Understandings: (Continued) A more detailed model of the atom describes the division of the main energy level into s, p, d, and f sub-levels of successively higher energies. Sub-levels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron. Each orbital has a defined energy state for a given electronic configuration and chemical environment and can hold two electrons of opposite spin. 4

2.2: Electron Configuration Applications and Skills: Description of the relationship between colour, wavelength, frequency, and energy across the electromagnetic spectrum. Distinction between a continuous spectrum and a line spectrum. 5

2.2: Electron Configuration Applications and Skills: (Continued) Description of the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels. Recognition of the shape of an s atomic orbital and the p x, p y, and p z atomic orbitals. Application of the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z=36. 6

Review of Topic 2.1 Let’s Review! List how many protons, neutrons and electrons the following elements have: o Li o C o O o Mg o P o Ar o Ca 7

Review of Topic 2.1 Atomic Structure Review: Protons and Neutrons are located in the nucleus. Electrons are found in the electron cloud outside the nucleus. This unit will focus on the electron cloud and where to find electrons. 8

Bohr Model Bohr Model for the Atom: Nucleus in the center with protons and neutrons Electrons in layers or levels around nucleus 9 Driver.layer.com

Bohr Model Bohr Model for the Atom: Electrons are arranged in energy levels (layers) Shows the number of electrons in each energy level. Electron orbits are circular paths 10 Chemistry.tutorcircle.com

Bohr Model Bohr Model for the Atom: Which element is this the electron arrangement for? How do you know? 11 Chemistry.tutorcircle.com

Bohr Model Bohr Model for the Atom: Useful for explaining and predicting chemical properties Based on the fundamental idea that electrons exist in definite, discrete energy levels Electrons can move from one energy level to another 12

Bohr Model Bohr Model for the Atom: Limitations of this model: Assumes all orbits are fixed Assumes all energy levels are circular Suggests incorrect scale for atom 13

2.3.4 Deduce the electron arrangement for atoms and ions. Models of Electron Arrangement J. J. Thomsen ( ) Negatively charged electrons stuck into a lump of positively charged material. Plum Pudding Model Neils Bohr ( ) Electrons arranged in circular paths, or orbits, around the nucleus. Planetary Model Erwin Schrodinger ( ) Used mathematics (quantum theory) to describe the location and energy of an electron. Quantum Mechanical Model 14

Quantum Mechanical Model Quantum Mechanical Model: Sophisticated mathematical theory that incorporates wave-like nature of electrons Based on 2 key ideas: o Schrodinger’s Equation o Heisenberg’s uncertainty principle 15

16 Deduce the electron arrangement for atoms & ions. The Quantum-Mechanical Model Electrons are not found at certain distances from the nucleus but are located in a region in space that is described by a set of 4 quantum numbers. The exact location and path of the electron can’t be determined. It estimates the probability of finding an electron within a certain volume of space surrounding the nucleus. Electron positions can be represented by a fuzzy cloud surrounding the nucleus (electron cloud).

Quantum Mechanical Model Heisenberg’s Uncertainty Principal It is impossible to determine accurately both the momentum and the position of a particle simultaneously. It is not possible to state precisely the location of an electron and its exact momentum, we can calculate the probability of finding an electron in a given region of space 17

Quantum Mechanical Model Schrodinger’s Equation o Formulated in 1926 by Austrian physicist Erwin Schrodinger o Equation integrates the dual wave-like and particle nature of the electron o Describe atomic orbitals: a region in space where there is a high probability of finding an electron. 18

U3. Quantum Mechanical Model The Quantum Mechanical Model 4 Quantum Numbers: n: energy level (called the principal quantum number) l: sublevels m l : orbital m s : spin 19

U3. Quantum Mechanical Model The Quantum Mechanical Model Each energy level can hold a maximum of electrons based on 2n 2 Ex: if n is 3, there can be up to 2(3) 2 electrons =18 e- Argon has 18e- and is at the end of energy level 3 20

21 Deduce the electron arrangement for atoms and ions. Principle Level / Energy Level / Shell = n Describes how far the electron is from the nucleus (Size) These levels are assigned values in order of increasing energy: (n=1,2,3,4,…). The larger the number, the more energy the electrons have and are farther from the nucleus.

22 Deduce the electron arrangement for atoms and ions. 2 nd Quantum Number Energy sublevel = l Describes the shape of the space the electron can be found in (Shape) There are four sublevels we will use and they are given letters to describe them s: spherical shape; 1s, 2s, 3s, 4s, 5s … p: dumbbell-shaped; 2p, 3p, 4p, 5p … d: clover-leaf shape; 3d, 4d, 5d … f: 4f, 5f, …

23 Deduce the electron arrangement for atoms and ions. 3rd quantum number Orbitals = m l Describes how many different arrangements in space the sublevels can have. (Orientation) Every s sublevel has 1 position Every p sublevel has 3 positions Every d sublevel has 5 positions Every f sublevel has 7 positions

24 Deduce the electron arrangement for atoms and ions. 4th quantum number: Spin = m s Any orbital can only have a maximum of two electrons, each having opposite spins. Every s sublevel can have a maximum of 2 e - s Every p sublevel can have a maximum of 6 e - s Every d sublevel can have a maximum of 10 e - s Every f sublevel can have a maximum of 14 e - s

Quantum Mechanical Model Sublevels: There are 4 types: s, p, d, and f Each type has a characteristic shape, specific number of orbitals and associated energy. Each orbital holds a maximum of 2 electrons 25

Quantum Mechanical Model Sublevels: s: spherical shape, 1 orbital, holds 2 e- 26

Quantum Mechanical Model Sublevels: P: dumbell shaped, 3 orbitals, holds 6 e- Draw the three orbital shapes 27

Quantum Mechanical Model Sublevels: d: double dumbell shaped, 5 orbitals, holds 10 e- f: funky shaped, 7 orbitals, holds 14 e- 28

Quantum Mechanical Model Identify which part of the Hog Hilton analogy represents each of the terms on the left. Differentiate between: Electron Energy level Orbital Spin Sublevel 29

Order of Energy Levels for Electrons This order must be followed every time! Each level must be filled before moving to the next level 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p 30

31 Order of orbitals (filling) in multi-electron atom 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p 7.7

Order of Electrons 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

A & S 5: Electron Configuration Three principles (rules) must be followed when representing electron configurations: 1. Aufbau Principle: electrons fill the lowest energy orbital first 2. Pauli Exclusion Principle: any orbital can hold a maximum of 2 electrons and those electrons have opposite spin. 3. Hund’s rule of maximum multiplicity: when filling orbitals of equal energy, electrons fill all orbitals singly before occupying in pairs 33

Diagramming Electron Arrangement There are three methods for diagramming electron arrangement Electron Configuration Orbital Filling Diagram Electron / Lewis Dot Diagram – we will learn this later 34

A&S 5. Electron Configuration Orbital Filling Diagrams 1.Draw a box or line for each orbital and sublevel. (s = 1 line, p = 3 lines, d= 5 lines) 2.Place arrows to denote electrons. Maximum of 2 electrons per box. The first arrow is pointing up, the second arrow is pointing down to represent opposite spins. 3.Within a sublevel, each orbital must get an electron before the second electron is added ( ie: each p sublevel gets one before doubling up) 35

A&S 5. Electron Configuration Orbital Filling Draw orbital filling diagrams for the following atoms. H Be O 36

A&S 5. Electron Configuration Orbital Filling Draw orbital filling diagrams for the following atoms. H__ 1s Be __ __ 1s 2s O __ __ ___ ___ ___ 1s 2s 2p x 2p y 2p z 37

Orbital Filling Draw a box or circle or line for each orbital. Place arrows to denote electrons. Maximum of 2 electrons per box. The first arrow is pointing up, the second arrow is pointing down to represent opposite spins. Within a sublevel, each orbital must get an electron before the second electron is added Draw orbital filling diagrams for the following atoms. H__ 1s Be __ __ 1s 2s O __ __ __ __ __ 1s 2s 2p x 2p y 2p z 38

Orbital Filling Draw a box or circle or line for each orbital. Place arrows to denote electrons. Maximum of 2 electrons per box. The first arrow is pointing up, the second arrow is pointing down to represent opposite spins. Within a sublevel, each orbital must get an electron before the second electron is added Draw orbital filling diagrams for the following atoms. Al Ca Zr 39

Orbital Filling Draw a box or circle or line for each orbital. Place arrows to denote electrons. Maximum of 2 electrons per box. The first arrow is pointing up, the second arrow is pointing down to represent opposite spins. Within a sublevel, each orbital must get an electron before the second electron is added Draw orbital filling diagrams for the following atoms. Al __ __ __ __ __ __ __ __ __ 1s 2s 2p x 2p y 2p z 3s 3p x 3p y 3p z Ca __ __ __ __ __ __ __ __ __ __ 1s 2s 2p x 2p y 2p z 3s 3p x 3p y 3p z 4s 40

Orbital Filling Draw a box or circle or line for each orbital. Place arrows to denote electrons. Maximum of 2 electrons per box. The first arrow is pointing up, the second arrow is pointing down to represent opposite spins. Within a sublevel, each orbital must get an electron before the second electron is added Draw orbital filling diagrams for the following atom. Zr __ __ __ __ __ __ __ __ __ 1s 2s 2p x 2p y 2p z 3s 3p x 3p y 3p z __ __ __ __ __ __ __ __ __ 4s 3d 4p x 4p y 4p z __ __ __ __ __ __ 5s 4d 41

Electron Configuration Start with 1s and follow the order until all the electrons in the atom have a place. Draw electron configurations for the following elements: H1s 1 Be1s 2 2s 2 O1s 2 2s 2 2p 4 Al Ca Zr 42

Electron Configuration 2p 4 Energy Level Sublevel Number of electrons in the sublevel 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 … etc.

Electron Configuration Start with 1s and follow the order until all the electrons in the atom have a place. Draw electron configurations for the following elements: H1s 1 Be1s 2 2s 2 O1s 2 2s 2 2p 4 Al1s 2 2s 2 2p 6 3s 2 3p 1 Ca1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Zr1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 2 44

Electron Configuration The superscripts should add up to the number of electrons in the species. Draw electron configurations for the following elements: H1s 1 = 1 electron Be1s 2 2s 2 = 4 electrons O1s 2 2s 2 2p 4 = 8 electrons Al1s 2 2s 2 2p 6 3s 2 3p 1 = 13 electrons Ca1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 = 20 electrons Zr1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 2 = 40 electrons 45

Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental) Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental) s orbitals p orbitals d orbitals f orbitals

49 ns 1 ns 2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 5 ns 2 np 6 d1d1 d5d5 d 10 4f 5f Ground State Electron Configurations of the Elements

What two blocks will always be the highest occupied level?

Review of Topic 2.1 Let’s Review Ions! What is an ion? What are the two types of ions? List how many protons, neutrons and electrons do the following ions have: o Li +1 o O -2 o Mg +2 o P -3 o Al +3 52

Review of Topic 2.1 List how many protons, neutrons and electrons do the following ions have: o Li +1 p = 3, n = 4, e = 2 o O -2 p = 8, n = 8, e = 10 o Mg +2 p = 12, n = 12, e = 10 o P -3 p = 15, n = 16, e = 18 o Al +3 p = 13, n = 14, e = 10 53

A & S 5. Ion Configuration Cations lose electrons, Anions gain electrons For orbital filling : add or subtract the number of electrons in the charge, draw in electrons For Electron Configuration : add or subtract the number of electrons in the charge, fill in sublevels and orbitals 54

A & S 5. Ion Configuration Practice: Compare Mg and Mg +2 Number of electrons: Electron configuration: Compare O and O -2 Number of electrons: Electron configuration: 55

A & S 5. Condensed Configuration This is a short cut! Full electron configurations become lengthy and cumbersome with increasing atomic number Condensed Configurations are more convenient 56

A & S 5. Condensed Configuration Core Electrons: electrons that are in the inner energy levels Valence Electrons: electrons that are in the outer energy level Condensed = [nearest noble gas] + valence electrons 57

A & S 5. Condensed Configuration Condensed = [nearest noble gas core] + valence electrons Ex: Oxygen: has 8 e-, 2 in core and 6 valence, nearest noble gas is He (which has 2 e-) [He] 2s 2 2p 4 58

A & S 5. Condensed Configuration Condensed = [nearest noble gas core] + valence electrons Ex: Cobalt: 59

A & S 5. Condensed Configuration Condensed = [nearest noble gas core] + valence electrons Practice: Write the condensed configuration Cl: Zn: Mn: 60

exceptions o orbitals “like” to be empty, half filled, or full therefore, an electron leaves the 4s (leaving it half full) and goes to the 3d in order to make it full Cr we would predict: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 but it is actually: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Cu we would predict: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 but it is actually: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 Weird electron configuration video (3:24)

Electron Dot Diagram Using the symbol for the element, place dots around the symbol corresponding to the outer energy level s & p electrons (valence electrons). Will have from one to eight dots in the dot diagram. Draw electron dot diagrams for the following atoms H Be O Al Ca Zr H Be O 62

Electron Dot Diagram Using the symbol for the element, place dots around the symbol corresponding to the outer energy level s & p electrons. Will have from one to eight dots in the dot diagram. Draw electron dot diagrams for the following atoms Al Ca Zr 63

2.3.4 Deduce the electron arrangement for atoms and ions. Write electron configuration, orbital filling diagrams, and electron dot diagrams. Kr Tb 64

2.3.4 Deduce the electron arrangement for atoms and ions. Write electron configuration, orbital filling diagrams, and electron dot diagrams. Kr 65

2.3.4 Deduce the electron arrangement for atoms and ions. Write electron configuration, orbital filling diagrams, and electron dot diagrams. Tb 66

A & S 1: Wavelength, Frequency and Energy Light consists of electromagnetic waves that can travel through space and matter All electromagnetic waves travel in a vacuum at 3.0 x 10 8 m/s (c) Three components of electromagnetic waves Amplitude (y): Height from the origin to the crest Wavelength (λ) : Distance between the crests Frequency (ν): Number of wave cycles to pass a given point per unit time Related by the equation c = λν Draw and label the wave diagram. 67

The electromagnetic spectrum is an arrangement of all of the types of electromagnetic radiation in increasing order of wavelength or decreasing frequency The higher the frequency, the shorter the wavelength and the higher the energy E = hf E = energy (Joules) h = Planck’s constant (6.63 x J s) F = frequency (s -1 ) Radiowaves: long wavelength, low energy Ultraviolet waves: short wavelength, high energy 68

A & S 1: Wavelength, Frequency and Energy Create a table, showing the relationship between wavelength (λ) frequency (v) energy (E) distance of electron fall color in the visible spectrum wavelengths of each color on opposite ends of the visible spectrum 69

Know the following about the EM spectrum: visible, infrared, and UV regions describe the variation in: o wavelength o frequency o energy o colors for visible light The electromagnetic spectrum Know what is in the red boxes

The Electromagnetic Spectrum High frequency Short wavelength High energy lower frequency longer wavelength lower energy

2.3.1Describe the electromagnetic spectrum 74

2.3.2Distinguish between a continuous spectrum and a line spectrum Emission Spectrum Emission showing all wavelengths and frequencies produced by excited electrons 75

2.3.2Distinguish between a continuous spectrum and a line spectrum Continuous Spectrum Emission showing a continuous range of wavelengths and frequencies; all the colors together without any space between them Line Spectrum Emission of specific elements showing a series of discrete lines; individual lines of color with space between each line On your paper, draw the difference between a continuous and line spectrum 76

2.3.2Distinguish between a continuous spectrum and a line spectrum 77

2.3.2Distinguish between a continuous spectrum and a line spectrum 78 We use an instrument called a spectroscope to detect the emission spectrum for a given source of light.

2.3.3Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels Why do we see different colors of light in line spectra? When electrons of a gaseous atom get excited, they are raised to a higher energy level. The extra energy is released as light when they drop back down to lower energy levels. The energy is provided by thermal or electrical energy Each element has its own unique line spectrum 79

More on emission line spectrum Give off energy when falls back down to normal energy level

emission line spectrum o energy is applied to a specific element this “excites” the element and the light is viewed through a spectroscope o a continuous spectrum is NOT observed, but a series of very bright lines of specific colors with black spaces in-between instead o unique for every element and are used to identify atoms (much like fingerprints are used to identify people)

1.an electron in the atom gains (absorbs) energy from heating 2.electron jumps up an energy level. 3.electron is now unstable (unwelcome) in this level and is “kicked out” 4.when the electron loses the energy and come back to the original level, light is emitted this is a repetitive slide- just couldn’t bear to delete it

The Atomic Emission Spectrum of Hydrogen the emission spectrum of hydrogen is the simplest emission spectrum because there is only one electron o if had more than one electron, they would influence the other’s position o it is not uniform, but concentrated into bright lines, indicating the existence of only certain allowed electron energy levels

The Atomic Emission Spectrum of Hydrogen Emission spectrum of hydrogen: UJ48s UJ48s THANK YOU, Bill Nye!!! Atomic Emission Animation (silent): SPxI SPxI BBC Spectroscopy of Stars: NvpI NvpI

2.3.3Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels You MUST know the hydrogen line series. Draw this in your notes. Indicate the colors and the wavelengths, notice the space between colors. 87

2.3.3Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels Red = 656 nm Blue-green = 486 nm Blue = 434 nm Violet = 410 nm 88

More about energy levels Energy levels of atoms are NOT evenly spaced like the rungs of a ladder o the higher the energy level, the smaller the difference in energy between successive energy levels becomes o the energy difference between levels becomes less as the level number increases o this means that the lines of a spectrum will converge (get closer together with increasing energy)

90

convergence up here (levels are close together)

2.3.3Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels 93

2.3.3Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels When electrons absorb energy they can move out to higher energy levels (the excited state). When they fall back to the lower energy level (the ground state) they release the energy as light. Depending on how far they fall, different colors of light are given off. Red = short fall (low E) Violet = long fall (high E) The Balmer Series of lines (visible light) is formed when the electrons fall back to the second energy level (n=2) 94

2.3.3Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels The line spectrum that we see from visible light is called the Balmer series. Similar sets of lines can be seen in ultraviolet (Lyman Series) and Infrared (Paschen Series) Electrons move in orbits around the nucleus of the atom. Each orbit has a fixed amount of potential energy. The farther from the nucleus the orbit is, the more potential energy it has. 95

IB-- this is referred to as convergence of the spectral lines

This is referred to as convergence of the spectral lines.

2.3.3Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels Lyman Series: Ultraviolet Transition of electrons from outer levels to n=1. Balmer Series: Visible Transition of electrons from outer levels to n=2. Paschen Series: Infrared Transition of electrons from outer levels to n=3 Spectral lines converge at increased values of n due to closer spacing of energy levels 98

2.3.4 Deduce the electron arrangement for atoms and ions. Let’s Review! List how many protons, neutrons and electrons the following elements have: o Li: p = 3, n = 4, e = 3 o C: p = 6, n = 6, e = 6 o O: p = 8, n = 8, e = 8 o Mg: p = 12, n = 12, e = 12 o P: p = 15, n = 16, e = 15 o Ar: p = 18, n = 22, e = 18 o Ca: p = 20, n = 20, e = 20 99

2.3.4 Deduce the electron arrangement for atoms and ions. Electrons are found in energy levels (n) or shells outside the nucleus of the atom. Energy levels are labeled from 1-7, 1 is closest to the nucleus and 7 is farthest away Energy level 1 holds a max of 2 electrons Energy level 2 holds a max of 8 electrons Energy level 3 holds a max of 18, we will only work with the first 8 spaces Energy level 4 holds a max of 32, we will only work with the first 2 spaces 100

2.3.4 Deduce the electron arrangement for atoms and ions. Always fill electrons from energy level 1 out 1 has 2 spaces 2 has 8 spaces 3 has 8 spaces 4 has 8 spaces Ex: Lithium has 3 electrons 2 electrons in level 1, 1 electron in level 2 Electrons always travel in pairs, but want to be alone before they are put in pairs 101

2.3.4 Deduce the electron arrangement for atoms and ions. Electrons travel in pairs, but want to be alone before they are put in pairs If there are 4 electrons fill one on each side If there are 7 electrons, fill one on each side first (4) then add the second to each side to get 7, there will be one electron that is alone 102

2.3.4 Deduce the electron arrangement for atoms and ions. Electron arrangement: how to write out the number of electrons in each level Ex: Carbon 6 electrons 2 in level 1 4 in level 2 Electron arrangement= 2.4 Ex: Sodium has 11 electrons 2 in level 1, 8 in level 2, 1 in level 3 Electron arrangement =

2.3.4 Deduce the electron arrangement for atoms and ions. Electron arrangement Practice Be = 2.2 O = 2.6 Mg = Cl = Ca = Electron arrangement Practice Be (4 electrons) O (8 electrons) Mg(12 electrons) Cl (17 electrons) Ca (20 electrons) 104

2.3.4 Deduce the electron arrangement for atoms and ions. Let’s Review Ions! List how many protons, neutrons and electrons do the following ions have: o Li +1 o O -2 o Mg +2 o P -3 o Al

2.3.4 Deduce the electron arrangement for atoms and ions. Let’s Review Ions! List how many protons, neutrons and electrons do the following ions have: o Li +1 p = 3, n = 4, e = 2 o O -2 p = 8, n = 8, e = 10 o Mg +2 p = 12, n = 12, e = 10 o P -3 p = 15, n = 16, e = 18 o Al +3 p = 13, n = 14, e =

2.3.4 Deduce the electron arrangement for atoms and ions. Electron Arrangement for Ions: o Li +1 e = 2 EA = 2. o O -2 e = 10EA = 2.8 o Mg +2 e = 10 EA = 2.8 o P -3 e = 18 EA = o Al +3 e = 10 EA =