Chapter 3 Arrangement of Electrons in the Atom

Spectra When white light passed through a prism= Continuous spectrum

Hydrogen Spectra Niels Bohr studied light emitted from a Hydrogen discharge tube and observed a spectrum consisting of a series of narrow coloured lines.(emission line spectrum) Use spectrometer or hand held spectroscope.

Line Spectra

FlameFlame Tests In this experiment we investigate the colours given off by certain salts

Flame Emission Spectra Photographs of flame tests of burning wooden splints soaked in different salts. Include link to web page methane gas wooden splintstrontium ioncopper ionsodium ion calcium ion

Emission Spectrum of Hydrogen 1 nm = 1 x m = “a billionth of a meter” 410 nm434 nm486 nm656 nm

Emission Spectrum of an Element 1 nm = 1 x m = “a billionth of a meter” 410 nm434 nm486 nm656 nm 1 nm = 1 x m = “a billionth of a meter”

Continuous and Line Spectra 4000 A o light Na H Ca Hg nm Visible spectrum  (nm)

Continuous and Line Spectra 4000 A o A o light Na H Ca Hg

Flame Emission Spectra Photographs of flame tests of burning wooden splints soaked in different salts. Include link to web page methane gas wooden splintstrontium ioncopper ionsodium ion calcium ion

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Bohr’s Theory Why is an emission spectrum produced? -Bohr theory 1.Electrons revolve around the nucleus in fixed paths called orbits. 2.Electron has a fixed amount of energy. 3.As long as the electron is in one energy level it neither loses nor gains energy. 4.When the electron absorbs energy- jumps from a lower to a higher energy level. 5.Energy lost as it falls from a higher energy level to a lower energy level. 6.Definite amount of energy lost.

Each definite amount of energy emitted gives rise to a line in the emission spectrum. Since only definite amounts of energy are emitted, therefore electrons can occupy only definite energy levels. Light emitted h= Plank’s constant f = frequency of light emitted E 2 -E 1 =hf Line spectra applet

Each electron transition results in a line of a particular wavelength. N=1 Lyman series Ultraviolet N=2 Balmer series visible lines N=3 Paschen series Infra Red

Limitations to Bohr Theory Limitations of the Bohr Theory It is mainly successful in explaining one electron system e.g. Hydrogen. The Bohr Theory does not explain the splitting of lines in the emission spectra and therefore the account for the existence of sublevels. Doesn’t take into account the wave nature of electrons The presence of atomic orbitals is also not accounted for.

Atomic Absorption Spectrometry. Atoms can also absorb light. If white light is passed through a gaseous sample of an element it is found that the light that comes out has certain wavelengths missing. We can analyse and correspond these to various elements.

Spectra

Atomic absorption spectrometer

Differences of spectra An emission spectrum consists of coloured lines against a black background. An absorption spectra consists of dark lines against a coloured background One use of Atomic absorption spectrometry is the analysis of water for metals like lead, mercury. Cadmium(heavy metals)

Wave nature of the electron Wave nature of the electron. When the Bohr theory is applied to atoms with more than 1 electron it failed to account for many of the lines in the emission spectrum of these atoms. De Broglei-moving electrons had a wave motion associated with them.

Heisenberg’s Uncertainty principle Heisenberg’s Uncertainty Principle.-It is impossible to measure at the same time both the velocity and position of electrons. This gives rise to the probability of finding electrons in a particular position inside the atom. Gives rise to atomic orbitals.

Quantum Mechanics Heisenberg Uncertainty PrincipleHeisenberg Uncertainty Principle –Impossible to know both the velocity and position of an electron at the same time Microscope Electron  Werner Heisenberg ~1926

Atomic orbital’s An atomic orbital is a region in space where there is a high probability of finding an electron. Erwin Schroedinger.— used maths equations to work out probability of finding an electron in any particular sublevel in an atom.

Quantum Mechanics Schrödinger Wave EquationSchrödinger Wave Equation (1926) quantized –finite # of solutions  quantized energy levels probability –defines probability of finding an electron Courtesy Christy Johannesson Erwin Schrodinger ~1926

Quantum Mechanics Orbital (“electron cloud”) –Region in space where there is 90% probability of finding an electron Courtesy Christy Johannesson Electron Probability vs. Distance Electron Probability (%) Distance from the Nucleus (pm) Orbital 90% probability of finding the electron

Orbitals S-orbital-----spherical in shape P-orbital Dumbell in shape

Relative Sizes 1s and 2s 1s 2s Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 334

Quantum Numbers Principal Quantum Number n 1. Principal Quantum Number ( n ) –Energy level –Size of the orbital –n 2 = # of orbitals in the energy level Courtesy Christy Johannesson 1s1s 2s2s 3s3s

Quantum Numbers pxpx pzpz pypy x y z x y z x y z

p-Orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 335 pxpx pypy pzpz

Quantum Numbers s p d f Angular Momentum Quantum # l 2. Angular Momentum Quantum # ( l ) –Energy sublevel –Shape of the orbital Courtesy Christy Johannesson

Quantum Numbers Orbitals combine to form a spherical shape. 2s 2p z 2p y 2p x Courtesy Christy Johannesson

Quantum Numbers n=# of sublevels per level n 2 =# of orbitals per level Sublevel sets: 1 s, 3 p, 5 d, 7 f Courtesy Christy Johannesson n = 3 n = 2n = 1 Principallevel Sublevel Orbital ssp sp d pxpx pypy pzpz d xy d xz d yz dz2dz2 d x 2 - y 2 pxpx pypy pzpz

Maximum Capacities of Subshells and Principal Shells n n l Subshell designation designation s s p s p d s p d f Orbitals in subshell subshell Subshell capacity capacity Principal shell capacity capacity n 2 Hill, Petrucci, General Chemistry An Integrated Approach  1999, page 320

Quantum Numbers Magnetic Quantum Number m l 3. Magnetic Quantum Number ( m l ) –Orientation of orbital –Specifies the exact orbital within each sublevel Courtesy Christy Johannesson

d-orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 336

Shapes of s, p, and d-Orbitals

Atomic Orbitals

s, p, and d-orbitals A s orbitals: Hold 2 electrons (outer orbitals of Groups 1 and 2) B p orbitals: Each of 3 pairs of lobes holds 2 electrons = 6 electrons (outer orbitals of Groups 13 to 18) C d orbitals: Each of 5 sets of lobes holds 2 electrons = 10 electrons (found in elements with atomic no. of 21 and higher) Kelter, Carr, Scott,, Chemistry: A World of Choices  1999, page 82

Principal Energy Levels 1 and 2

Quantum Numbers Spin Quantum Number 4. Spin Quantum Number ( m s ) –Electron spin  +½ or -½ –An orbital can hold 2 electrons that spin in opposite directions. Courtesy Christy Johannesson

Quantum Numbers 1. Principal #  2. Ang. Mom. #  3. Magnetic #  4. Spin #  energy level sublevel (s,p,d,f) orbital electron Pauli Exclusion PrinciplePauli Exclusion Principle –No two electrons in an atom can have the same 4 quantum numbers. –Each electron has a unique “address”: Courtesy Christy Johannesson

Electron Orbitals: Electron orbitals Equivalent Electron shells (a) 1s orbital (b) 2s and 2p orbitalsc) Neon Ne-10: 1s, 2s and 2p 1999, Addison, Wesley, Longman, Inc.

Feeling overwhelmed? Read Section 4-2! Courtesy Christy Johannesson "Teacher, may I be excused? My brain is full." Chemistry