Chapter 4 Notes The Arrangement of Electrons in the Atom

1. Chapter 4 NotesThe Arrangement of Electrons in the Atom

2. 4.1 - Developing a New Atomic Model • Rutherford’s Model (gold foil experiment) – doesn’t explain nuclear forces or why electrons don’t collapse into the nucleus. Neils Bohr • Bohr’s Model – explains the relationship between light and an atom’s electrons. Electrons are in specific energy levels (Hydrogen Spectrum). • Light behaves like a wave.  1900 – Light has properties of waves and particles. • Electromagnetic Radiation (EMR) = energy that exhibits wave-like behavior through space. (e.g. – xrays, UV, IR, visible light, microwaves, radio waves.)

3. The Electromagnetic Spectrum* • All EMR moves at a constant speed: c = 3.00x108 m/s (speed of light) • (in a vacuum vs. reality)

4. Wavelength(l- lambda) - the distance between corresponding points on adjacent waves (m, nm, cm) Frequency (n - nu) - number of waves that pass a given point in a specified time (s-1, 1/s, Hz) C = ln = 3.00x108 m/s l is inversely proportional to n because their product equals a constant (c).

5. The Photoelectric Effect – the emission of electrons from a metal when light shines on it. • The PEE relates to the frequency (n) of the light, NOT the light’s intensity. Therefore, light has to be of a minimum frequency for the PEE to occur. • Max Planck – studied emission of light by hot objects. He proposed that light is not emitted continuously, but in small energy “packets” called quanta. • A quantum of energy is the minimum quantity of energy that can be lost or gained by an electron. E = h n • E is the energy of a quantum of radiation (Joules, J) • nis the frequency in s-1 of the radiation emitted. • h = Planck’s constant = 6.626x10-34 J*s • n.b.: by rearranging the equations, we also get E = hc/l

6. Einstein (1905) – EMR has a dual wave-particle nature. Light is like a stream of particles called “photons” and each photon carries a quantum of energy. • Photon – a particle of electromagnetic radiation having zero mass and carrying a quantum of energy. • Einstein explains the PEE! EMR is absorbed by matter only in whole numbers of photons. To be ejected, an electron on a metal’s surface must be struck by a single photon with at least the minimum energy required to knock the electron loose. (minimum E correlates to a minimum n) Ephoton varies for different metals.

7. Hydrogen Atom Line-Emission Spectrum • Ground State – the lowest energy state of an atom. • Excited State – a state in which an atom has higher potential energy (PE) than in its ground state. • When an atom goes from an excited state to a ground state, it gives off emr (e.g. - neon lights). • Electric current is passed through Hydrogen gas at low pressure  pink light  passed through a prism  4 colors of the visible spectrum. (Hydrogen’s line emission spectrum) • The result: The Lyman Series (UV), Balmer Series (visible) and Paschen Series (IR). • Surprise!!! – NOT a continuous spectrum. NOT a continuous range of frequencies of emr.

8. Quantum Theory • When a H atom falls from an excited to a ground state, it emits a photon of radiation. • Ephoton = hn = the difference between the two energy states = E2-E1. • (1913) Niels Bohr linked the H atom’s electron to photon emission. Electrons can only circle the nucleus in allowed paths or orbits. When the electron is in one of these orbitals, the atom has a definite fixed energy. • Lowest/Ground State = electron is closest to the nucleus. • Higher Energy States/Excited States are farther from the nucleus. (similar to the “rungs of the energy ladder” – higher up ~ higher PE of an electron) • Electrons can also absorb energy packets to move up in energy. • Absorption (energy is added to the atom) vs. Emission (Energy is given off)

9. The H – Emission Spectrum • Lyman Series: E6E1, E5E1, E4E1, E3E1, E2E1 (5 lines) • Balmer Series: E6E2, E5E2, E4E2, E3E2 (4 lines) • Paschen Series: E6E3, E5E3, E4E3 (3 lines) • Problems with Bohr’s Model: • It’s only good for the hydrogen atom. • It doesn’t explain the chemical behavior of atoms.

10. 4.2 - The Quantum Model of the Atom • Why does the Bohr Model work with electrons existing only at certain energy levels? • Can we think of electrons as both waves and particles, like we think of light? (deBroglie – 1924) • deBroglie: “Let’s think of electrons as waves confined to the space around the nucleus.” • E=hn • Electrons can only have specific energies and frequencies. • Electrons can be bent (diffracted) and have interference like waves. • If an electron is both a wave and a particle, WHERE are they in the atom?

11. The Heisenberg Uncertainty Principle: (1927) “It is impossible to locate the position and velocity of an electron or any other particle at the same time.” • Schrodinger Wave Equation: (1926) Schrodinger used electron’s dual wave-particle nature to develop an equation that proved the quantization of electron’s energy levels. The solutions to the Schrodinger Equation are called wave functions. Y2 gives the probability of finding an electron at a given place around the nucleus. • The result: Orbitals are not neat spheres like Bohr thought, but a 3-D region that indicates the probable location of an electron. • The HUP and SWE laid the foundation for quantum theory.

12. The 4 Quantum Numbers • Specify the properties of atomic orbitals and the properties of electrons in the orbitals. • describe the location of electrons (similar to home address) • Principal Q.N. (n) describes the main energy level occupied by the electron. (n = 1, 2, 3…) • Angular Momentum Q.N. (l) describes the sublevels (shapes) of orbitals. (0  l  n-1)

13. 3. Magnetic Q.N. (ml) indicates the orientation of an orbital around the nucleus ( -l  m  +l ) 4. Spin Q.N. (ms = +½ or -½) indicates the two fundamental spin states of an electron in an orbital. A single orbital can only hold 2 electrons which must have opposite spin states. Practice: How many electrons can fit in the third main energy level (n=3)? • l = 0, 1, 2 (s, p, d) • The s orbital can hold 2 electrons (2). • The 3 p orbitals (px, py, pz) can hold 2 each (6) • The 5 d orbitals (dxy, dyz, dxz, dx2-y2, dz2) can hold 2 each (10) Answer: 2 + 6 + 10 = 18 electrons can fit.

14. Practice: Which atom is represented by the quantum numbers (3, 2, 0, -½)? We’re looking at the 3d orbital level. The spin Q.N. (-½) tells us we are filling electrons into empty shells. Then we count over -2, -1, 0. This brings us to Vanadium (V). Try these: • (2, 0, 0, +½) • (4, 1, -1, +½) • (5, 3, 0, -½)

15. 4.3 Electron Configuration – the arrangement of electrons in an atom. (ground state configurations) The Rules: • Aufbau Principle – an electron occupies the lowest energy state that can accept it. (1s2s2p3s3p4s*3d4p) • Pauli Exclusion Principle – no 2 electrons in an atom can have the same 4 quantum numbers. (therefore, if 2 electrons have the same n, l & ml, then they will have opposite spins, ms) • Hund’s Rule – orbitals of equal energy are each occupied by one electron before any orbital is occupied by a 2nd electron, and all electrons in singly occupied orbitals must have the same spin state.

16. 3 Notations can represent electrons within an atom in their ground state: • Orbital Notation • Electron Configuration Notation • Noble Gas Notation Nitrogen has 7 electrons and can be represented 3 different ways: • Orbital Notation (see picture on the board) • Electron Configuration Notation: 1s22s22p3 • Noble Gas Notation: [He]2s22p3

17. Do together: Chlorine & Magnesium Try on your own: Oxygen & Aluminum