Chapter 5 Electrons in Atoms

5.1 Light and Quantized Energy Light, a form of electromagnetic radiation has characteristics of both waves and particles.

The Wave Nature of Light Light is a type of electromagnetic radiation –Form of energy that exhibits wavelike behavior as it travels through space

Characteristics of waves –Wavelength (λ): Shortest distance between equivalent points on a wave –Amplitude: - waves height from origin to crest or origin to trough

Frequency (ν): number of waves that pass a given point per second –SI unit = hertz (Hz)

All electromagnetic waves, including light travel at 3.00 x 10 8 m/s (c) c = λν

Electromagnetic spectrum –Visible light is only a very small part of the electromagnetic spectrum –EM spectrum includes all forms of electromagnetic radiation –Differences between types of radiation = frequency & wavelength

What is the wavelength of a microwave that has a frequency of 3.44 x 10 9 Hz? Light reflected from a green leaf has a wavelength of 4.90 x m. What is the frequency of the light?

The Particle Nature of Light Wave model of light cannot explain all the ways light interacts with matter –Metal glows when heated –Iron = grey at room temp, then glows red, then orange, the blue

The quantum concept –Matter can gain or lose energy only in small, specific amounts called quanta Stairs –Quantum = minimum amount of energy that can be gained or lost by an atom

Energy of a quantum = Planck’s constant x frequency –Planck’s constant = x J*s E quantum = hν

The photoelectric effect –Electrons are emitted from a metal’s surface when light at or above a certain frequency shines on the surface

Einstein proposed the dual nature of light (wave-like & particle- like) to explain the photoelectric effect Light = beam of bundles of energy (photons) –Photons = massless particles that carry a quantum of energy

E photon = hv Every object gets its color by reflecting a certain portion of visible light. The color is determined by the energy of the reflected photons. What is the energy of a photon from violet light if it has a frequency of x Hz

The blue color in fireworks has a wavelength of 4.5 x m. How much energy does one photon of this light carry?

Atomic Emission Spectra Set of frequencies of the electromagnetic waves emitted by atoms of the element Found by putting gas into tube and charging it. Light viewed through prism & colors separate –Each line of color = photon with specific energy

Each element’s emission spectrum is unique and can be used to identify the element.

5.2 Quantum Theory and the Atom Electrons behave much like a person climbing a ladder

Bohr’s Model of the Atom Proposed atoms only have certain allowable energy states Ground state – lowest allowable energy state for an atom Excited state – an atom that has gained energy

Electrons move around nucleus in circular orbits (energy levels) Orbits closest to nucleus have lowest energy

Bohr assigned a number, n, to each orbit (quantum number) Orbit closest to nucleus = energy level 1 & n = 1

Quantum theory & emission spectra When electrons are excited the move to higher energy levels Electron drops from higher to lower energy level and emits a photon corresponding to the energy difference between the two levels –Photon = color of light shown

Limits of Bohr Model –Only explains Hydrogen’s emission spectrum –Does not account for chemical behavior of atoms –Electrons do NOT move around nucleus in circular orbits

Quantum Mechanical Model Louis de Broglie – proposed electrons can have wave characteristics Werner Heisenberg – it is impossible to know the exact position of an electron at any time (Heisenberg uncertainty principle)

The act of seeing an electron changes its path!

The only thing that can be known about an electron is the probability for it to occupy a certain region around the nucleus Schrodinger – treated electrons like waves & developed the quantum mechanical model of the atom

Atomic orbital – 3D region around the nucleus which describes an electrons probable location –More dense areas = more likely to find electron

Atomic Orbitals Bohr’s quantum numbers are now called energy levels Principal energy level – the major energy levels of an atom –Lowest energy level n = 1 –n values range 1 – 7

Energy Sublevels Smaller levels within energy levels where electrons can move (s, p, d, f) –n = 1 : 1 sublevel –n = 2 : 2 sublevels –n = 3 : 3 sublevels –n = 4 – 7 : 4 sublevels

Sublevels are made of orbitals Each type of sublevel has a different number and shape of orbitals 2 electrons can fit in each orbital

S sublevel = 1 orbital p sublevel = 3 orbitals d sublevel = 5 orbitals f sublevel = 7 orbitals Each orbital can be filled: (2 e-), half filled: (1 e-), or empty: (0 e-)

Shapes of sublevels S = sphere (1 orbital) p = peanut shell (3 orbitals) d = clover leaf (5 orbitals) f = shapes vary (7 orbitals)

Energy Levels, Sublevels & Orbitals Energy level 1 = closest to nucleus –One sublevel = s –One orbital = 1s Energy level 2 –2 sublevels = s and p –4 orbitals = 2s, 2p x, 2p y, 2p z

Energy level 3 –Three sublevels = s, p, and d –Nine orbitals = 3s, 3p x, 3p y, 3p z, five 3d orbitals Energy level 4 –Four sublevels = s, p, d, and f –16 orbitals Higher energy levels have all four sublevels

Catholic Central as an atom Pretend you are an electron Floor = Classroom = Desks =

5.3 Electron Configurations Electron configuration = the arrangement of electrons in an atom Most stable, lowest energy configuration = ground state electron configuration

Which sublevels are in each energy level? How many orbitals are in each sublevel? How many electrons can each orbital hold?

The Aufbau Principle The lowest energy levels are filled with electrons first

The Pauli Exclusion Principle Only two electrons can occupy an orbital The electrons must have opposite spins

Hund’s Rule When electrons occupy orbitals of equal energy, one electron enters each orbital until all orbitals contain one electron with spins parallel

Draw the orbital diagram for the following atoms Li O Cl

Important questions using a Cl atom Highest full energy level? Highest full sublevel? Highest occupied energy level? Number of unpaired electrons? Number of empty orbitals?

Electron configurations from orbital diagrams Na P Ar

Electron Configurations Beyond Row 3 Energy levels 3 and up overlap 4s is lower energy than 3d so it comes first! Follow your diagonal diagram

What are the electron configurations of the following? Co Kr Ba

Abbreviated electron configurations Elements in column 8A are noble gasses All very stable and unreactive All except He have 8 valence electrons –Ne: 1s 2 2s 2 2p 6 –Kr: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6

Noble gas configurations = base Ex: Ca –Ar: –Ca: –Ca abbreviated:

How to write abbreviated configurations 1.Find the noble gas in the row above the given element 2.Write its symbol in brackets 3.Add electrons into proper orbitals until total number of electrons is reached  Make sure you start after brackets with correct orbital!

Fr Pt As

EXCEPTIONAL ELECTRON CONFIGURATIONS Up to atom number 23 there are no exceptions! After that exceptions may occur in transition metals Partially full orbitals are more stable that empty orbitals Electrons will sometimes move to fill empty orbitals.

EX: Mo [Kr]5s 2 4d 4 [Kr]5s 1 4d 5 Which version of molybdenum is more stable? Would Zr be an exception?

Valence Electrons Electrons in the outermost s & p Determine an atom’s properties

Boron: 1s 2 2s 2 2p 1 –Highest energy level –Valence electrons Scandium: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 –Highest energy level –Valence Electrons

Electron Dot Structures Valence electrons are the only ones that are involved in chemical reactions Dot structuresshow the valence electrons

How to draw a dot structure 1.Figure out how many valence electrons element has 2.Write the element symbol 3.Add dots to top, right, bottom, and left of symbol one at a time until all valence electrons are used 4.Remember: each orbital can only hold 2 dots

ElementElectron Configuration# of VEDot Diagram Phosphorus Bromine Oxygen