ATOMIC STRUCTURE Electrons in Atoms

In the early 1900’s scientists studying the behavior of atoms observed that certain elements emitted visible light when heated by a flame or if subjected to a high voltage spark Much of our understanding of how electrons behave in atoms comes from the study of light In order to understand this, you must know about the nature of light…

LIGHT Visible light (the light we see with our eyes) is a type of electromagnetic radiation All other electromagnetic radiation is invisible

ELECTROMAGNETIC RADIATION Electromagnetic (EM) radiation is the transmission and emission of energy in the form of electromagnetic waves Visible light, X-rays, microwaves , infrared waves (IR), ultraviolet waves (UV), radio waves, etc.

ELECTROMAGNETIC SPECTRUM The electromagnetic spectrum encompasses all forms of electromagnetic radiations

CHARACTERISTICS OF WAVES A wave can be thought of as a vibrating disturbance by which energy is transmitted Waves can be characterized by their: Wavelength  (lambda): distance between identical points on a successive wave nanometers (nm) Frequency  (nu): the number of waves that pass through a particular point in 1 second hertz (Hz); 1 Hz = 1 cycle/s Amplitude: the vertical distance from the midline to the peak or trough

High Frequency Low Frequency

The speed of a wave is a product of its wavelength and frequency: SPEED OF LIGHT Another important property of waves is their speed dependent on type of wave and medium it’s traveling through The speed of a wave is a product of its wavelength and frequency: c =  Notice  and  are inversely proportional In a vacuum, All electromagnetic waves travel at the speed of light (3.00 x 108 m/s)

PLANCK’S QUANTUM THEORY According to the theory atoms and molecules can emit or absorb energy only in discreet quantities, which Planck called “quantum” quantum (meaning “fixed amount”) is the smallest quantity of energy that can be gained or lost by an atom The energy E of a single quantum of energy is given by E =h where h is called Planck’s constant (6.63 x 10-34 J•s) and  is the frequency of radiation According to Planck’s theory energy is always released or absorbed by matter in whole-numbers of h, 2h, 3h …

THE PHOTOELECTRIC EFFECT Albert Einstein used Planck’s quantum theory to explain the photoelectric effect When light shines on a clean metal surface the surface emits electrons (can be converted to electrical energy) For each metal there is a minimum frequency of light needed to emit electrons red light in incapable of releasing electrons from sodium metal even if intense faint violet light releases electrons easily

DUAL NATURE OF LIGHT Ephoton = h In explaining the photoelectric effect Einstein proposed that light (electromagnetic radiation) has dual nature: Can behave like a wave Can behave like a stream of particles (photons) PHOTON - a particle of electromagnetic radiation A photon has no mass A photon carries a quantum of energy Einstein calculated that a photon’s energy depends on its frequency Ephoton = h

ATOMIC EMISSION SPECTRA Light of a neon sign is produced by passing electricity through a tube of neon gas The atoms in the tube absorb energy and become excited and unstable They become stable by releasing the energy as light

ATOMIC EMISSION SPECTRA If the light emitted from neon is passed through a prism neon’s atomic emission spectrum is produced

ATOMIC EMISSION SPECTRA The atomic emission spectrum of an element is its “fingerprint” It’s the set of frequencies EM waves the element’s atoms emit (give off) Each element’s atomic emission spectrum is unique and can be used to identify the element Hydrogen’s emission spectrum: violet 410 nm blue-violet 434 nm blue-green 486.1 nm red 656.2 nm

Light Bulb (white light) Hydrogen Bulb

THE BOHR MODEL OF THE ATOM Bohr proposed: Electrons move around the nucleus in circular orbits (“rings”) with distinct energy levels smaller orbits have lower energy, larger orbits higher energy In other words, electrons found closer to the nucleus has less energy than electrons found at greater distances from the nucleus Bohr assigned a quantum number (n) to each level

BOHR’S MODEL OF THE ATOM + Energy n = 7 n = 6 n = 5 - n = 4 n = 3 n = 2 n = 1 This model is often called the planetary model

BOHR’S ATOM CONTINUED The lowest energy state of an atom is its ground state When an atom gains energy (through heating for example) it is in an excited state in an excited state the electron absorbs the energy & jumps to higher energy level when it jumps back down to its ground state it releases excess energy in the form of light Even though hydrogen contains only one electron, it can have many excited states

BOHR MODEL CONTINUED Because electrons jump between orbitals that have specific energy levels only certain frequencies of electromagnetic radiation can be given off (only certain colors can be emitted): If an excited electron drops from n=3 to n=2 red light is emitted If an excited electron drops from n=4 to n=2 blue-green light is emitted 5-2: blue light 6-2: violet light This is how Bohr explained hydrogen’s emission spectrum

E = h E = h

Wait! Bohr’s model explained the emission spectrum of Hydrogen, but it did not explain the emissions of any other element! It was eventually found that Bohr was incorrect: Electrons do not travel in circular orbits around the nucleus In 1924, a young French scientist Louis de Broglie (1892-1987) proposed an idea that eventually accounted for the fixed energy levels of Bohr’s model and better explained the behavior of electrons

THE QUANTUM MECHANICAL MODEL OF THE ATOM If waves can act like particles, particles can act as waves! Electrons behave like waves Can’t know electrons position/path around the nucleus: Electrons move about in a cloud around the nucleus in what appears to be a random pattern The Quantum Model only predicts where an electron is likely to be found

HEISENBERG UNCERTAINTY PRINCIPLE According to this principle it is fundamentally impossible to know the exact position and velocity of a particle at the same time locating an electron produces uncertainty in the position and motion of the electron It’s like trying to locate a helium filled balloon in a completely darkened room: If you locate it by touch, you change it’s position-Once you find it, it’s already somewhere else!

SO WHERE CAN ELECTRONS BE FOUND? In the quantum model, the nucleus is not surrounded by orbits, but by atomic orbitals Atomic Orbital: a three-dimensional pocket of space around the nucleus that the electron is most likely to be found An electron has a 90% chance of being found in the atomic orbital That is the best we can do! Electron probability density for hydrogen

Where 90% of the e- density is found for the 1s orbital e- density (1s orbital) falls off rapidly as distance from nucleus increases

ATOMIC ORBITAL ORGANIZATION Principal energy level (n: 1-7) (n) indicates relative size and energy of orbital. As n increases so do energy and size Energy sublevels (s, p, d, f) sublevels labeled according to shape s: spherical; p: dumbbell; d/f: varied Orbitals: Each sublevel has a specific number of orbitals: s: 1orbital p: 3 orbitals d: 5 orbitals f: 7 orbitals Each orbital can hold two electrons!!!

increasing energy

Each orbital can hold two electrons 4p Energy 3d 4s 3p 3s H Hydrogen 1 1.008 He Helium 2 4.003 Li Lithium 3 6.941 C Carbon 6 12.01 B Boron 5 10.81 Be Beryllium 4 9.012 O Oxygen 8 16.00 Ne Neon 10 20.18 N Nitrogen 7 14.01 F Fluorine 9 19.00 2p       2s   1s  

ORDER OF ORBITALS (FILLING) IN MULTI-ELECTRON ATOM

ELECTRON CONFIGURATION An atoms electron configuration is the way an atom’s electrons are distributed among the orbitals of an atom The most state stable electron configuration is an atom’s ground state Ground state: all electrons are in the lowest possible energy state Electron configuration represented by writing symbol for the orbital and a superscript to indicate the number of electrons in the orbital Li: 1s2 2s1

The Pauli Exclusion Principle The two electrons in an orbital must spin in opposite directions   1s  2s  2p       4s 3d 3s 3p paramagnetic unpaired electrons diamagnetic paired electrons

HUND’S RULE Negatively charged electrons repel each other, so: Electrons won’t pair up unless they have to Once there is one electron in every orbital…the pairing will begin! 2s 1s  2p  1. Add an electron: 2s 1s  2p  2. Add an electron: 2s 1s  2p  3. Add an electron: 2s 1s  2p  4.

ELECTRON CONFIGURATION The periodic table can be divided into four distinct blocks based on valence electron configuration electron configuration explain the recurrence of physical and chemical properties

SHORTHAND (NOBLE GAS) NOTATION Shows electron filling starting from previous noble gas: Na: 1s22s22p63s1 Noble gas configuration: [Ne]3s1

Electron Configuration for Cation and Anions What is the electron configuration for a sodium atom? Na: 1s22s22p63s1 What is the electron configuration for a sodium ion (Na+)? Na+: 1s22s22p6 or [Ne]

TRANSITION CATIONS When transition metals form cations electrons are removed from the s-orbital before the d-orbital Mn: [Ar]4s2 3d3 Mn2+: [Ar]3d3 This happens because d-orbitals are more stable than s-orbitals in transition metals The s-orbitals are always in a higher energy level

ISOELECTRONIC What do you notice about the following electron configurations: F- : 1s22s22p6 or [Ne] O2- : 1s22s22p6 or [Ne] N3- : 1s22s22p6 or [Ne] They have the same number of electrons and therefore identical ground-state electron configurations These three ions are isoelectronic

VALENCE ELECTRONS Valence electrons are the electrons found in the outermost orbitals of the atom (highest energy level) They are the electrons involved in bonding and are responsible for the chemical properties of an element Carbon has 4 valence electrons: 1s2 2s2 2p2 How many does magnesium have?

VALENCE ELECTRONS & GROUP NUMBER One of the most important relationships in chemistry: Atoms in the same group have similar chemical properties because they have the same number of valence electrons! For the representative elements (group A elements) group # = number of valence electrons Exceptions: He is in group 8 but only has 2 valence electrons

The octet/duet rule Atoms will gain, lose or share electrons to achieve noble gas configuration, meaning all atoms want a full outer orbital: 2 valence electrons for He 8 valence electrons for all other noble gases

ELECTRON-DOT STRUCTURE Chemists often represent valence electrons in electron-dot structures Electron-Dot Structure consists of the element’s symbol surrounded by dots representing the atom’s valence electrons valence electrons are placed one to each of the four sides first, when each side has one dot, you may begin doubling up S