1 Keep until June 2011! Unit 2.2: Electrons

2 Vocabulary: Electron: negatively-charged subatomic particleOrbital: (atomic orbital) a region in an atom where there is a high probability of finding electrons. Quantum (L. “how much”) the amount of energy required for an electron to move from one energy level to another. Valence shell: the outermost energy level in an atom that is occupied by electrons; furthest from the nucleus Valence electrons: electrons found in the valence shell; can participate in bonding Octet rule: (“octo-” = 8) valence shell is most stable with 8 electrons

3 I. More Atomic Models (5.1) Rutherford’s model: electrons orbit around the nucleus similar to planets orbiting a star, However, orbits are not at a fixed distance from nucleus Does not explain chemical properties of atoms Why do metals give off characteristic color when burned? Why do objects glow red, then yellow, then white as temperature increases?

4 Niels Bohr ( ): proposed that electrons are found in a specific spherical path (orbit) around the nucleus: Each orbit has a fixed energy – the electrons in those orbits have a fixed energy level Like the rungs of a ladder: lower rungs have less potential energy than higher rungs Requires more energy to move farther from the nucleus

5 Energy levels are not equally spacedA quantum (L. “how much”) is the amount of energy required for an electron to move from one energy level to another a. Cannot be “between” energy levels just as you cannot “hover” between rungs on a ladder Energy levels are not equally spaced Higher energy levels are closer together, so it takes less energy to move between levels (If electron receives enough energy, it can escape the atom altogether.)

6 Energy levels: labeled by “principal quantum numbers” (n)Erwin Schrodinger ( ) develops the Quantum Mechanical Model; uses mathematical expressions to define location of electrons around the nucleus Energy levels: labeled by “principal quantum numbers” (n) n = 1, 2, 3… up to 7; corresponds to the energy level (rung of ladder) being filled Each energy level represents a period on the Periodic Table

7 Each principal energy level (n) has one or more atomic orbitals (sublevels of energy)Atomic orbital: a region around the nucleus of an atom where there is a high probability of finding an electron; like a fuzzy “electron cloud” n =1, so 1 sublevel of energy: an “s” orbital n =2, so 2 sublevels of energy: an “s” and a “p” orbital n =3, so 3 sublevels of energy: “s”, “p” and “d” orbitals n =4, so 4 sublevels of energy: “s”, “p”, “d”, and “f” orbitals

8 Orbitals have different shapess-orbital Spherical shape Holds 2 electrons p-orbital Dumbbell shape Has 3 sub-orbitals (px; py; pz), each holds 2 e-’s p-orbital holds 6 e-’s total d-orbital Has 5 sub-orbitals, each holds 2 e-’s d-orbital holds 10 e-’s total f-orbital: 7 sub-orbitals f-orbital holds 14 e-’s total

9 Valence shell: (7.1) The valence shell is the outermost energy level in an atom, furthest from nucleus e-’s in this shell are called valence electrons Valence e-’s can be donated or “stolen” to form bonds (more on that later…) The “group” that an element is in corresponds to its number of valence e-’s. Group # (Roman numerals) = # of valence e-’s (generally) Or group # -10 = # of valence e-’s In general, the number of valence electrons of a representative element is equal to the group number.

10 Octet rule: “octo-” = 8 Valence shell is most stable with 8 e-’s“s” (2 e-’s) and “p” (6 e-’s) orbitals are both full Noble gases satisfy the octet rule, are stable, generally nonreactive/inert If only 1 or 2 e-’s in valence shell, will tend to give up those e-’s to become stable. If 6 or 7 e-’s in valence shell, will tend to steal e-’s from other atoms to become stable.

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12 II. Electron Configurations (5.2)Three rules that predict the order in which electrons fill orbitals: The aufbau principle (Ger. “building up”): electrons occupy lowest energy levels first Uses aufbau diagram (also called orbital diagrams) to show filling of orbitals Within any energy level, the “s” orbital has the lowest energy so it will fill first But sometimes, principal energy levels can overlap: 4s orbital has less energy than 3d orbital even though 4th energy level is a greater quantum number

13 The Pauli Exclusion Principle: each orbital can only hold a maximum of two electrons, each of opposite “spin” “Spin” is clockwise or counterclockwise Denoted by vertical arrows (↑↓) (but not quite…)

14 Hund’s Rule: one electron enters each sub-orbital until all sub-orbitals contain one electron with the same spin direction. There are exceptions to orbital-filling, but up to atomic # 23, elements obey aforementioned rules. (now it’s correct!)

15 Longhand method for electron configuration:For any element, write the energy level (principal quantum number, n) and the letter designations of the orbitals that are filled, using superscripts to denote the number of electrons in that orbital – huh?! Ex: Oxygen is atomic number 8, in period 2 of the Periodic Table: The first energy level has only an “s” orbital that can hold 2 electrons: 1s2 The second energy level is being filled (period 2) The “s” orbital fills first b/c lower energy than “p”: 2s2 Then, the “p” orbitals fill until all the remaining 4 electrons have been placed: 2p4 So the final electron configuration for oxygen is 1s22s22p4 Note that the superscripts all add up to the atomic number!

16 Shorthand method for writing electron configuration:The outermost energy levels of the noble gases (last group on the Periodic Table) are completely filled. For any element, write [in brackets] the noble gas that precedes it, followed by the remaining energy levels and orbitals. Ex: Write the shorthand configuration for Magnesium: Magnesium is atomic number 12, meaning it has 12 electrons in its energy levels. It is in Period 3 on the periodic table, so the third energy level is being filled. The noble gas that precedes Magnesium is Neon. The longhand configuration would have been: 1s22s22p63s2, but because the electron configuration for Neon is 1s22s22p6, you can substitute [Ne] when writing the shorthand configuration. So the shorthand configuration is [Ne]3s2

17 III. Atomic Spectra When an electron absorbs energy, it moves to an excited state in a higher energy level, further from the nucleus As it returns to its ground state (closer to the nucleus), it emits (gives off) that absorbed energy in the form of light.

18 Amount of energy absorbed/emitted corresponds to the frequency (wavelength) of light given off:Less energy (ex: only moves 1 energy level) means a longer wavelength, lower frequency, so “light” will be nearer the red end of the electromagnetic (EM) spectrum. As more energy is emitted (ex: moves down 2 or more energy levels) means a shorter wavelength, higher frequency, so “light” will be closer to the violet end of the EM spectrum.

19 Each element has an atomic emission spectrum that is unique from all others.Emission spectra are seen as discrete bands of light.

20 Three series of emission spectra:Lyman series: describes electron emissions in the ultraviolet (UV) spectrum Electrons are moving from higher energy levels down to n = 1.

21 Balmer series: describes electron emissions in the visible spectrumElectrons are moving from higher energy levels down to n = 2 Paschen series: describes electron emissions in the infrared (IR) spectrum Electrons are moving from higher energy levels down to n = 3. Keep until June 2011!