Electrons In Atoms.

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Presentation transcript:

Electrons In Atoms

Historical View of Light Early 1900s, light thought to behave as a wave Later, light was discovered to have particle-like characteristics Today, light is thought to have both wave and particle properties

Wave Description of Light Electromagnetic Radiation – form of energy that exhibits wavelike behaviors as it travels through space Travels at the speed of light in air 3.0 E 8 m/s Ex: visible light, microwaves, x-rays, radio, etc. Measures wavelength and frequency

Wavelength λ λ λ Symbol – λ (lambda) Definition – distance between equivalent points on a adjacent waves λ λ λ

Frequency = 1/s = s-1 Symbol –  (sometimes f) Definition – number of waves that pass a given point in a one second Unit: Hertz (Hz) Ex: 300 Hz = 300 1/s = 300 s-1 = 1/s = s-1

Electromagnetic Waves All EM waves travels at the speed of light (c) c = speed of light (3 x 108 m/s) λ = wavelength v = frequency (sometimes f) c = λ = 3 x 108 m/s

Wavelength and Frequency Relationship Inversely related: as one increases, other decreases short λ high frequency long λ low frequency

Electromagnetic Spectrum Definition – shows all forms of electromagnetic radiation arranged according to wavelength/frequency

Ex problem #1 = λ = = c 4.76 x 10-8 m  c = λ 3 x108 m/s What is the wavelength of a radiation with a frequency of 6.30 x 1015 Hz? c = λ c 3 x108 m/s = 4.76 x 10-8 m λ = =  6.30 x 1015 s-1

Ex problem #2 = = c  = λ 3.108 x 1015 Hz c = λ 3 x108 m/s What is the frequency of a radiation, which has a wavelength of 9.653 x 10-8 m? c = λ c 3 x108 m/s =  = = 9.653 x 10-8 m λ 3.108 x 1015 Hz

Particle Nature of Light Photoelectric Effect – emission of electrons from metal’s surface when light of specific frequency shines on surface Radiation emitted from the object is in small, specific amounts called quanta Quantum – minimum amount of energy that can be gained or lost by an atom light e- METAL

Ephoton= h E= Energy (measured in Joules) Photon – particle of light(electron magnetic radiation) having zero mass and carrying a quantum of energy Photon Energy Energy of photon is directly proportional to the frequency of radiation Ephoton= h E= Energy (measured in Joules) h = Planck’s constant = 6.626 x 10-34 J·s  = frequency (s-1)

Ex problem E = h = (6.626 x 10-34 J·s) (6.582 x 1014 s-1) #2) How much energy does a photon of light have if its frequency is 6.582 x 1014 Hz? E = h = (6.626 x 10-34 J·s) (6.582 x 1014 s-1) = 4.361 x 10-19 J

Hydrogen Line-emission Spectrum When atoms in the gaseous state are heated, their energy increases. Ground state- state of lowest energy Excited state- higher potential energy than ground state

Hydrogen Line-emission Spectrum When light shines through a prism, it is separated into a series of specific frequencies and wavelengths of visible light. The bands of light are part of hydrogen’s line emission spectrum

Modern View of Light Wave Theory- waves are forms of energy Particle Theory- particles are pieces of matter Modern view- combines both Einstein’s Theory of Relativity combines both matter and energy into one formula containing the speed of light E= mc^2

Electromagnetic Spectrum

Bohr Model of an Atom Ground State – lowest energy state of an atom Excited State – state when atom gains energy **pay attention to the electrons** Bohr Model – shows electron orbit and energy level of an electron

Bohr Model E1 E2 E3 E3 > E2 > E1 E1 = lowest energy level The electron of the hydrogen atom can circle the nucleus in paths called orbits In an orbit- electron has a definite fixed energy Lowest energy state- closest to nucleus Total energy of electron increases as it moves farther from nucleus E1 E2 E3 E3 > E2 > E1 E1 = lowest energy level

Ground State to Excited State 4 6 in ground state, no energy radiated in excited state, electrons jump to higher energy level electron go from high E level to low E level photon emitted 1) 5 4 2) 3 3 Energy of atom 2 2 3) 1 4) 1

Quantum Theory Describes mathematically the wave properties of electrons and is based on: Heisenberg Uncertainty Principle: impossible to determine position and velocity of a particle at the same time Schrödinger Wave Equation: equation that is used to describe electrons as waves

Quantum Numbers Definition- numbers that specify the properties of atomic orbitals and their electrons 1st quantum number: Main energy level or distance from nucleus 2nd quantum number: Orbital shape 3rd quantum number: Orbital orientation

Principal Quantum Number (n) Definition – indicates the energy level (shells) surrounding the nucleus - use periodic table to tell (look at rows)

Principal Quantum Number (n) n = 1,2,3,…..(values of n are positive) n=1 (closest to nucleus) As n increases, the distance of the energy levels from the nucleus increases 2n2 = number of electrons in each level

Principal Quantum Number

Angular Momentum Quantum Number (l) Definition – indicates shape of orbital that tells the path of the electrons In orbitals- different shapes occupy different regions called sublevels or subshells 4 sublevels: s, p, d, f S= lowest energy

s orbital Shape: electrons travel in a sphere

s orbital 3s 1s 2s The greater the energy level, the bigger the orbital

p orbital Shape: dumbbell or figure 8 shaped

D orbital Shape: double dumbbell

Types of Orbitals

Magnetic Quantum number (m) Orientation of an orbital about the nucleus S= 1 P= 3 D= 5 F=7

Spin Quantum Number (s) + ½ = clockwise turn -1/2= counterclockwise turn

First level S orbital= 2 electrons

Second level S orbital= P orbital= 1 orbital, 2 electrons 3 orbitals, 6 electrons

Third level S orbital= P orbital= D orbital= 1 orbital, 2 electrons 3 orbitals, 6 electrons D orbital= 5 orbitals= 10 electrons

Fourth level S orbital= P orbital= D orbital= F orbital= 1 orbital, 2 electrons P orbital= 3 orbitals, 6 electrons D orbital= 5 orbitals= 10 electrons F orbital= 7 orbitals, 14 electrons

Electron Configuration Definition – arrangement of electrons in an atom Atoms of different elements have different numbers of electrons Electrons will assume lowest energy Where are certain electrons located?

Rules Governing Electron Configurations 1) Pauli Exclusion Principle – no 2 electrons in the same atom can have the same set of 4 quantum numbers

Rules Governing Electron Configurations 2) Hund’s Rule – if orbitals have equal energy, one e- will go in each orbital before doubling up all electrons in singly occupied orbitals must have same spin 1 2 3 5 6 4

Rules Governing Electron Configurations 3) Aufbau Principle – electrons occupy lowest energy orbital available - fill up level 1 first, then level 2, etc.

Orbital filling table

Electron Configurations Orbital notation: An unoccupied orbital is represented by a line. An orbital with one electron is represented by a An orbital containing 2 electrons is represented as Sodium

Orbital Notation 7 n = 2 1s 2s 2p Nitrogen How many electrons? What energy level is nitrogen on? 7 n = 2 1s 2s 2p

Orbital Notation 14 n = 3 1s 2s 2p 3s 3p Silicon How many electrons? What energy level is silicon on? 14 n = 3 1s 2s 2p 3s 3p

Orbital Diagram 29 n = 4 1s 2s 2p 3s 3p 4s 3d Copper How many electrons? What energy level is copper on? 29 n = 4 1s 2s 2p 3s 3p 4s 3d

Electron Configuration Notation Definition: the number of electrons in a sublevel is represented by adding superscripts to the sublevel designation sodium

Electron Configuration Notation 2 2 4 Oxygen (8 e-) Sulfur (16 e-) Vanadium (23 e-) Zirconium (40 e-) 1s 2s 2p 2 2 6 2 4 1s 2s 2p 3s 3p 2 2 6 2 6 2 3 1s 2s 2p 3s 3p 4s 3d 2 2 6 2 6 2 10 6 2 2 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d

Noble Gas Notation Shortcut for electronic notation with a noble gas Noble gases are used because they have a filled outer shell of electrons (octet)

Noble Gas Notation [He] 2s 2p [Ne] 3s 3p [Ar] 4s 3d [Kr] 5s 4d 2 4 2 4 Rule: start from previous noble gas, then write the configuration Oxygen Sulfur Vanadium Zirconium 2 4 [He] 2s 2p 2 4 [Ne] 3s 3p 2 3 [Ar] 4s 3d 2 2 [Kr] 5s 4d

Valence Electrons 1s 2s 2p 5 valence e- 2 valence e- 1s 2s 2p 3s 1s 2s definition – electrons in outer most energy level - located in highest s & p orbitals (max 8) N: Mg: Se: 2 2 3 1s 2s 2p 5 valence e- 2 2 6 2 2 valence e- 1s 2s 2p 3s 2 2 6 2 6 2 10 4 1s 2s 2p 3s 3p 4s 3d 4p 6 valence e-

Electron Dot Structure Definition – shows number of valence e- by diagram, which are in the highest principal quantum number Nitrogen (5 v.e.) Magnesium (2 v.e.) Selenium (6 v.e.) N Mg Se

Irregular confirmations of Cr and Cu Chromium “promotes” a 4s electron to half fill its 3d sublevel Copper “promotes” a 4s electron to FILL its 3d sublevel