Modern Atomic Theory and Electron Configurations

Transcription

1 Chem 101 Modern Atomic Theory and Electron Configurations Lectures 8 and 9

2 Types of Electromagnetic Radiation Electromagnetic radiation is given off by atoms when they have been excited by any form of energy, as shown in flame tests.

3 Electromagnetic radiation (a beam of light) can be pictured in two ways: as a wave and as a stream of individual protons.

4 The wavelength of a wave is the distance between peaks.

5 A photon of red light (relatively long wavelength) carries less energy than does a photon of blue light (relatively short wavelength).

6 Properties of Electromagnetic Waves Velocity = c = speed of light x 10 8 m/s ( use 3.00 x 10 8 m/s ) All types of light energy travel at the same speed. Amplitude = A = measure of the intensity of the wave, i.e. brightness Wavelength = = distance between two consecutive peaks or troughs in a wave Generally measured in nanometers (1 nm = 10-9 m) Frequency = = the number of waves that pass a point in space in one second Generally measured in Hertz (Hz), 1 Hz = 1 wave/sec = 1 sec -1 c = Energy = h ν Energy is equal to Planck s constant times frequency Planck s constant, h, is 6.63 x Joule seconds

7 When salts containing Li +, Cu 2+, and Na + dissolved in methyl alcohol are set on fire, brilliant colors result: Li +, red; Cu 2+, green; and Na +, yellow. Hmco Photo Files

8 Emission of Energy by Atoms/Atomic Spectra Atoms that have gained extra energy release that energy in the form of light.

9 Atomic Spectra Line spectrum: very specific wavelengths of light that atoms give off or gain Each element has its own line spectrum, which can be used to identify that element.

10 When an excited H atom returns to a lower energy level, it emits a photon that contains the energy released by the atom. Each photon emitted by an excited hydrogen atom corresponds to a particular energy change in the hydrogen

11 The colors and wavelengths (in nanometers) of the photons in the visible region that are emitted by excited hydrogen atoms.

12 Atomic Spectra The atom is quantized, i.e. only certain energies are allowed.

13 The Bohr model of the hydrogen atom represented the electron as restricted to certain circular orbits around the nucleus. Energy of electron is related to the distance of electron from the nucleus

14 Bohr s Model Energy of the atom is quantized Atom can only have certain specific energy states called quantum levels or energy levels. When atom gains energy, electron moves to a higher quantum level When atom loses energy, electron moves to a lower energy level Lines in spectrum correspond to the difference in energy between levels

15 (a) The hydrogen 1s orbital. (b) The size of the orbital is defined by a sphere that contains 90% of the total electron probability.

16 Bohr s Model Ground state: minimum energy of an atom Therefore electrons do not crash into the nucleus The ground state of hydrogen corresponds to having its one electron in the n=1 level Excited states: energy levels higher than the ground state

17 Orbitals and Energy Levels Valence shell: the highest-energy occupied ground state orbit Regions in space of high probability for finding the electron. These are called orbitals. Each principal energy level contains one or more sublevels. Sublevels are made up of orbitals. Each type of sublevel has a different shape each and energy. Each sublevel contains one or more orbitals.

18 The 1s orbital.

19 A diagram of principal energy levels 1 and 2 showing the shapes of orbitals that compose the sublevels.

20 The shapes and labels of the five 3d orbitals.

21 Pauli Exclusion Principle No orbital may have more than 2 electrons. Electrons in the same orbital must have opposite spins. s sublevel holds 2 electrons (1 orbital) p sublevel holds 6 electrons (3 orbitals) d sublevel holds 10 electrons (5 orbitals) f sublevel holds 14 electrons (7 orbitals)

22 Orbitals, Sublevels & Electrons For a multiple-electron atom, build-up the energy levels, filling each orbital in succession by energy Degenerate orbitals: orbitals with the same energy e.g. Each p sublevel has 3 degenerate p orbitals

23 Electron Configurations For a set of degenerate orbitals, fill each orbital halfway first before pairing Electron configurations show how many electrons are in each sublevel of an atom describes where electrons are. - 1s 2 2s 1 is the electron configuration for a ground state Li - 1s 2 2s 2 2p 3 is for nitrogen

24 Electron Configurations Valence shell: highest energy level Electrons in the valence shell are called valence electrons. Core electrons: electrons not in the valence shell Often use symbol of previous noble gas in brackets to represent core electrons, giving [He]2s 2 2p 3 for nitrogen or [Ne]3s 2 for magnesium

25 Electron Configuration and the Periodic Table Elements in the same column on the periodic table have: Similar chemical and physical properties Similar valence shell electron configurations same numbers of valence electrons same orbital types different energy levels

26 The electron configurations in the sublevel last occupied for the first eighteen elements.

27 s 1 s 2 d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 p 1 p 2 p 3 p 4 p 5 s 2 p 6 f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14

28 The orbitals being filled for elements in various parts of the periodic table.

29 A box diagram showing the order in which orbitals fill to produce the atoms in the periodic table. Each box can hold two electrons.

30 The periodic table with atomic symbols, atomic numbers, and partial electron configurations.

31 Metallic Character: Form cations Lose electrons in reactions oxidized Oxidation is Loss of electrons - OIL The easier it is for an element to lose electrons, the more metallic character is has.

32 The classification of elements as metals, nonmetals, and metalloids.

33 Metallic Character Reactivity of metals increases to the left on the period and down in the column Follows ease of losing an electron Reactivity of nonmetals (excluding the noble gases) increases to the right on the period and up in the column

34 Trend in Ionization Energy Minimum energy needed to remove a valence electron from an atom Gas state The lower the ionization energy, the easier it is to remove the electron. Metals have low ionization energies Ionization energy decreases down the group. Valence electron farther from nucleus Ionization energy increases across the period. Left to right

35 The Group 1 elements: the farther down a group and element appears, the more likely it is to lose an electron.

36 The Group 2 elements: the farther down a group and element appears, the more likely it is to lose an electron.

37 Ionization energies tend to decrease in going from the top to the bottom of a group.

38 Ionization energies tend to increase from left to right across a given period on the periodic table.

39 Relative atomic sizes for selected atoms. Note that atomic size increases down a group and decreases across a period.

40 Electronegativity Measure of the ability of an atom to attract shared electrons Larger electronegativity means atom attracts shared electrons more strongly 40