Structure of the atom

Transcription

1 Structure of the atom What IS the structure of an atom? What are the properties of atoms? REMEMBER: structure affects function! Important questions: Where are the electrons? What is the energy of an electron? Chapter 3 Part 1 Dual Nature of Light λν=c E = hν Line Spectrum Bohr Model of Hydrogen Atom Chem 110 1

2 Visible (white) light contains a continuous distribution of frequencies of electromagnetic radiation. Spectrum: distribution of ν in emitted radiation Chem 110 2

3 Different types of light produce different spectra DEMONSTRATIONS: How can we explain all of this? Examples Spectrum Type ν s Laser monochromatic one Light bulb continuous all Hg vapor discrete (or line) a few Hg vapor spectrum Chem 110 3

4 Observation of line spectra implies that atoms have discrete (quantized) energy levels. Excited state (E 3 ) Excited state (E 2 ) absorption ΔE = hν emission ΔE = hν Ground state (E 1 ) Chem 110 4

5 There are only 4 emission lines in the visible spectrum of hydrogen H 2 discharge tube Chem 110 5

6 Each element has a unique line spectrum. (Structure affects function) Chem 110 6

7 To explain this phenomenon, we start with the simplest atom (hydrogen) and try to understand it. H atom has 1 proton (+) and 1 electron ( ). Where is the proton? Where is the electron? Bohr Model of H atom (1913) Bohr proposes that the electron is in one of several possible orbits around the proton. Chem 110 7

8 If the electron is in an orbit, what is it s energy? Bohr used the line spectrum to figure out the energy differences between orbits and then deduced the energy of the electron in each orbit. E n 1 = RH 2 n n = 1, 2, 3,.. principal quantum number n = 6 n = 5 n = 4 n = 3 n = 2 R H = Rydberg constant R H = 2.18x10 18 J Energy Chem 110 n = 1 8

9 Energy of an Electron (Bohr Model) Energy is given off when an electron is put into orbital. Coulomb s Law helps Q1Q 2 E d where Q 1 = charge of electron (negative) Q 2 = charge of proton (positive) d = orbit radius (distance between nucleus and electron) Put electron into the orbital: attractive interaction Energy will be negative (means energy is given off) Reverse the process: try to remove the electron Energy will be positive (energy is absorbed) Note: Orbit energy in Bohr Model is negative, so it must correspond to energy needed to put electron into the orbit. Chem 110 9

10 Bohr Model of H atom (1913) Line spectrum is due to electronic transitions Atoms absorb or emit light when e changes its orbit ΔE = E f E i = hν 1 1 hν = ΔE = RH 2 2 ni nf where n i and n f are integers. This predicts the H-atom spectrum EXACTLY! Note: n f > n i ΔE is + (absorbs photon) n f < n i ΔE is (emits photon) Chem

11 Energy levels in Bohr Model n = 6 n = 5 n = 4 n = 3 n = 2 Energy n = 1 Chem

12 Be able to use the Bohr Model to solve problems describing electronic transitions in the Hydrogen atom. If n i = 2 and n f = 1, is energy emitted or absorbed? 1. emitted 2. absorbed Of the following transitions in an H-atom, which one results in the emission of the highest energy photon? 1. n=1 è n = 6 2. n=6 è n = 3 3. n=3 è n = 6 4. n=1 è n = 4 5. n=6 è n = 1 Chem

13 The Bohr model explained some experimental evidence for hydrogen atom, but it failed for other atoms. From Orbits to Orbitals : DeBroglie (1924): if light has dual wave/particle behavior, perhaps matter does also. Wavelength of matter waves: λ = h/mv Electron waves discovered in 1927 (Davidson and Garmer) (Basis for electron microscope) For a baseball and bacteria, λ is too small to observe, but for electrons λ is of atomic size producing profound effects. Electrons in atoms behave as "standing" waves. (Schrödinger equation, 1926) Enter the Quantum World Chem

14 Heisenberg Uncertainty Principle It is NOT possible to simultaneously know the position & velocity (momentum, mv) of a particle with complete certainty Derives from wavelike nature of matter This really becomes important when dealing with subatomic matter All electrons have a velocity, therefore, you cannot specify their exact location Contradicts Bohr s planetary model of the hydrogen atom In other words: It is not appropriate to imagine e moving in nice little orbits around the nucleus Can we say anything about where the e are? Chem

15 Solutions of Schrödinger equation are wavefunctions (Ψ) H Ψ= E Ψ Ψ(x,y,z) = wavefunction (no physical significance) Ψ 2 (x,y,z) = probability of finding one electron in a region of space, also called electron density Think of electrons as clouds of electron density. Orbitals = Ψ 2 (x,y,z) Chem

16 What is an orbital? Tells us WHERE the electron is Tells us the ENERGY of the electron An orbital specifies the probability of finding an electron in a given region of space, (i.e. orbitals have shapes) specifies the energy of the electron is characterized by quantum numbers (3 of them!) Chem

17 In 3-D expect 3 quantum numbers n, and m 1. Principal quantum # (n) Must be an integer Gives information about size 1s 2s Gives information about energy 3s Chem

18 Quantum numbers and nodes # of nodes is equal to n 1 Chem

19 The Quantum Number determines shape 2. Azimuthal quantum # ( ) integer ranging form 0 to n 1 Gives information about SHAPE Use symbols rather than numbers for = Name s p d f Chem

20 Quantum number m determines orientation 3. Magnetic quantum number (m ) ranges from to + Gives information about ORIENTATION Chem

21 Orbitals: Summary Allowed energy states for electrons in atom. Describes spatial distribution of electrons in these energy states. Orbital number shape name of orbitals? s 1 spherical p 3 dumbell d 5 clover leaf f 7?!?!? Quantum Numbers: defines n principal size azimuthal shape m magnetic orientation Chem

22 Make This Your Own Chem

23 Shells, Subshells & Orbitals n 2 = number of states = number of orbitals in the n th shell. Subshells # of orbitals 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f # orbitals in each subshell Chem

24 Shells, Subshells & Orbitals Shell: defined by Subshell: defined by quantum number quantum numbers Example: 3s (n = 3, = 0) 2p (n = 2, = 1) Orbitals of the same subshell have the same energy: they are degenerate Orbital: defined by quantum numbers Example: 2p x (n = 2, = 1) 2p y (n = 2, = 1) 2p z (n = 2, = 1) and m = 1, 0, 1 when = 1 Note: All of these have the SAME energy. Chem

25 No class on Monday: MLK Holiday By Wednesday You should have read Chapter 3 sections Complete Objective 2 in ALEKS (Dues Tues. Jan. 22) After Objective 2 you will be given a progress assessment in ALEKs. Before recitation on Thursday: complete Week 2 homework problems on pages 31 and 32. Chem