Unit 5 Electrons in Atoms Chemistry I Mr. Patel SWHS

Topic Outline Continue Learning Major Ions Atomic Models (5.1) Electron Configurations (5.2) Light and Quantum Mechanics (5.3) Lewis Dot Structures (7.1)

Atomic Models Democritus’s Model Thomson’s Plum Pudding Model Rutherford’s Model Electrons travel in orbit around nucleus Could NOT explain chemical properties of elements Need a model for electrons

Bohr Model Niels Bohr – electrons is found in a specific orbit around nucleus Each orbit has a specific energy = energy level The further away from the nucleus, the higher the energy

Bohr Model

Bohr Model An electron can move between levels Can not be between levels Think of a ladder An electron must gain or lose energy to change levels A quantum of energy – energy to move to another level

Bohr Model More energy between levels when closer to nucleus Less energy between levels when farther Energy levels get closer together

Bohr Model Ground state – lowest energy state for an electron Excited state – any higher energy state

Electron Excitation http://www.youtube.com/watch?v=4jyfi28i928&feature=relmfu

Bohr Model Each ring on a Bohr Model is labeled as “n” n must be a whole number n=1, n=2, n=3, etc. (period number) Each ring (n) can hold a specific number of electrons n=1 2 electrons n=2 8 electrons n=3 18 electrons n=4 32 electrons

Drawing Bohr (Rutherford) Diagrams http://www.youtube.com/watch?v=sKAzHE7A7rQ&feature=relmfu

Ex: Draw the Bohr Diagram for Hydrogen.

Ex: Draw the Bohr Diagram for Neon.

Ex: Draw the PEL Diagram for Calcium.

Ex: Draw the PEL Diagram for Argon.

Bohr Model Correct: Electrons have energy levels and can move Incorrect: Electrons move in orbits Matter has a Wave-Particle Duality

Dual Nature of Electrons Electrons as Particles Electrons as Waves Photoelectric Effect Young's Double Slit Experiment

Modern Theory Rutherford and Bohr based models on behavior of large objects Small objects behave differently – quantum mechanics Schrödinger Equation solutions  quantum mechanical model of the atom

Schrödinger Equation

The Cat – A Thought Experiment Schrodinger Cat 1 Schrodinger Cat 2

Quantum Mechanical Model Determines the allowed energies of the electrons The probability of where an electron is – electrons housed in electron clouds

Atomic Orbitals Region in space where there is a high probability of finding an electron Principal quantum number (n) – energy level think of the ring labels of the Bohr model Each energy level can be made up of sublevels – orbitals of similar energy but different shapes

1. s orbital Shape: sphere

2. p orbital Shape: Dumbbell

3. d orbital Shape: clover (mostly)

4. f orbital Shape: multiple clover

Atomic Orbitals http://www.youtube.com/watch?v=K-jNgq16jEY

Electron Configurations Electrons found in orbitals Electron configuration – ways in which various electrons are arranged in orbitals 4 orbitals: s (2 electrons), p (6 electrons) d (10 electrons), f (14 electrons)

Three Rules to find Elec. Config Aufbau Principle Electrons occupy orbitals of lower energy first For same n, low to high energy: s, p, d, f

Three Rules to find Elec. Config Pauli Exclusion Principle Each atomic orbital can have at most 2 electron Each electron in an orbital must have opposite spins 2 spins: spin up or spin down How we draw: 1 electron in s orbital: ____ 2 electrons in s orbital: ____ We use arrow with “half head”

Three Rules to find Elec. Config Hund’s Rule Electrons occupy orbitals to maximize spin For same n, place electrons spin up first then pair them with spin down 1 electron in p orbital ____ ____ ____ 2 electrons in p orbital ____ ____ ____ 3 electrons in p orbital ____ ____ ____ 4 electrons in p orbital ____ ____ ____ 5 electrons in p orbital ____ ____ ____ 6 electrons in p orbital ____ ____ ____

Orbital Blocks on PT s-block: Groups 1A and 2A (exception: He) p-block: Groups 3A-8A (exception: He) d-block: transition metals f-block: inner transition metals Remember, the period number is n = principal energy level

Orbital Blocks on PT

How to write electron configuration Ex: What is the electron configuration for O? O = oxygen, atomic number 8 = 8 electrons Draw spaces: ____ ____ ____ ____ ____ 1s 2s 2p Fill spaces according to rules: Write: 1s22s22p4

How to write electron configuration Ex: What is the electron configuration for C? C = carbon, atomic number 6 = 6 electrons Draw spaces: ____ ____ ____ ____ ____ 1s 2s 2p Fill spaces according to rules: Write: 1s22s22p2

3 ways to write electron configurations Using boxes and arrows ____ ____ ____ ____ ____ 1s 2s 2p Long EC: Cl: 1s22s22p63s23p5 Short EC: Cl: [Ne] 3s23p5 Put last noble gas in brackets and write electrons from there

Writing EC This is much easier than it looks. Simply, start at hydrogen and walk to the desired element counting all the elements you pass

Ex. Write EC (all three ways) for Boron.

Ex. Write EC (all three ways) for Mg.

Ex. Write EC (all three ways) for V.

Ex. Write EC (long and short) for Fr.

A Look Back… So far we have covered (and mastered): Evolution of the Atomic Model Democritus, Thomson, Rutherford, Bohr, QM Bohr Model and Bohr Diagram Quantum Mechanical Model and Orbitals Rules of Electron Configuration Writing Electron Configurations

Electrons Chemical reactions are the breaking and forming of bonds There are two types of bonds: covalent and ionic (and metallic) = next unit Bonding involves the movement of electrons

Valence Electrons Valence Electron: electrons in the highest occupied level These are the electrons that participate in bonding!!!

Valence Electrons You do not have to draw a Bohr Model every time you need to determine the VE’s The valence electrons (valency) for an atom is the same as the group number Note: In general, transition metals have two valence electrons.

Determine the valence electrons for: Ca Be O Si H Ne Ar 2 6 4 1 10

N B Lewis Dot Structures Show bonding electrons These structures show only valence electrons. How to draw: Write Symbol for element Determine group number Place that many (group number) dots around symbol N B

Lewis Dot Structures 1 2 3 4 5 6 7 8

Lewis Structures (Future)

Draw the Lewis Dot Structure for: Cs Al Ge Br

Practice Time!!!

The Octet Rule Remember that Noble Gases were very stable They all have 8 valence electrons (2 for He) FULL outer shell of electrons Every element will try to become like a noble gas The Octet Rule – atoms will try to have a full outer shell (= 8 electrons) when bonding

Cations Metals tend to lose electrons to have a full outer shell Cation – positively charge ion Results from metals losing electrons Naming: element name + ion Ex: Na = Sodium but Na1+ = sodium ion

Anions Nonmetals/Metalloids tend to gain electrons to have a full outer shell Anion– negatively charged ion Results from nonmetals gaining electrons Naming: element name with –ide ending + ion Ex: Br = Bromine but Br1- = bromide ion

Cations/Anions To determine the charge of an element’s ion, look at the group/column that it is in Group 1: 1+ Group 2: 2+ Group 3: 3+ Group 4: 0 Group 5: 3- Group 6: 2- Group 7: 1-

Lewis Dot Structures for Ions Draw the normal Lewis Dot structure for the neutral element Add electrons if gained or remove electrons if lost Place the appropriate charge 3- N N Nitride ion: Nitrogen:

Draw the Lewis Dot Structure for Phosphorus. Will this element for a cation or anion? What charge will it have? What will be the name of the ion? What noble gas is it similar to? Draw the Lewis Dot Structure for the ion.

Draw the Lewis Dot Structure for Barium. Will this element for a cation or anion? What charge will it have? What will be the name of the ion? What noble gas is it similar to? Draw the Lewis Dot Structure for the ion.

Light – A Wave Newton tried to prove light to be a particle However, experimental data showed that light was actually behaving as a wave The study of light led to the quantum mechanical model of the atom

Properties of Waves Wavelength (λ) Frequency (ν) Amplitude “lambda” Distance between crests Frequency (ν) “nu” Cycles per second Hertz (Hz) Amplitude Height from zero to the crest Crest Trough

Properties of Waves The Wave Equation: c = λ∙ν c = speed of light = 3 x 108 m/s λ = wavelength – must be in meters ν = frequency – in Hertz Energy of Light: E = h∙ν E = energy – in Joules h = Planck’s constant = 6.626 x 10-34 J∙s

Electromagnetic Spectrum When light passes through a prism, it is separated into different frequencies You need to know: Name and order of each regions Order based on wavelength Order based on frequency Order based on energy Details of the Visible Region

Electromagnetic Spectrum EM Spectrum Song Long Wavelength Low frequency Low Energy Short Wavelength High frequency High Energy

Review: Properties of Waves The Wave Equation: c = λ∙ν c = speed of light = 3 x 108 m/s λ = wavelength – must be in meters ν = frequency – in Hertz Energy of Light: E = h∙ν E = energy – in Joules h = Planck’s constant = 6.626 x 10-34 J∙s

Equations: c = λ∙ν E = h∙ν c = speed of light = 3 x 108 m/s h = Planck’s constant = 6.626 x 10-34 J∙s Ex: What is the frequency of a wave with a wavelength of 3.68 x 10-9 m? The energy?

Equations: c = λ∙ν E = h∙ν c = speed of light = 3 x 108 m/s h = Planck’s constant = 6.626 x 10-34 J∙s Ex: What is the frequency of a wave with a wavelength of 700 nm? The Energy?

Practice!!!

Atomic Spectra When atoms absorb energy, they move into higher energy levels These electrons then return back to a lower level and release energy as light Each atom releases light in a special way

Atomic Spectra Atomic Emission Spectrum – the frequencies of light released by an element split into separate discrete lines (unlike light)

Hydrogen Spectrum Balmer Series Visible Region Lyman Series UV Region

Light – A Particle Light deserves a quantum mechanical treatment Light also behaves as a particle and a wave (Particle-Wave Duality) Light particles called photons – packets or quanta of light (E = h∙ν)

deBroglie Relation Louis deBroglie determined that all matter that is moving can be considered as waves Large object have such a small wavelength that it can not be observed His math showed that as mass decreases, the wave function becomes more important (λ=h/mv)

Heisenberg Uncertainty principle Heisenberg Uncertainty Principle – it is impossible to know the exact velocity and position for a particle at the same time Heisenberg Uncertainty principle