Chapter 9: Electrons and the Periodic Table

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1 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 1 Chapter 9: Electrons and the Periodic Table Work on MasteringChemistry assignments What we have learned: Dalton s Indivisible Atom explained the Law of Constant Composition and the Law of Conservation of Mass and led to the Law of Multiple Proportions. J.J. Thomson, through Cathode Ray Tube experiments, discovered electrons are small negatively charged particles inside a divisible atom and came up with the Plum Pudding Model of the atom. Rutherford came up with the gold foil experiment shooting alpha particles through thin gold foil to test the Plum Pudding Model and discovered that some alpha particles were deflected. This led to Rutherford s Nuclear Model of the atom in which a heavy positive nucleus is surrounded by a cloud of electrons. Now we will further our knowledge of the atom by examining the Bohr model and the quantum-mechanical model which describe how electrons behave inside the atom and how those electrons affect the chemical and physical properties of elements. These models explain the observed emission/absorption spectra and the periodic behavior of the elements such as why He is inert. Pictured below are Niels Bohr (left) and Erwin Schrödinger (right)

2 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 2 Quantum Mechanics explains the behavior of the electrons inside atoms, the periodic law and expectations in chemical bonding. Electrons are extremely tiny. Electron behavior determines much of the behavior of atoms Problem: Directly observing electrons is impossible. Observing an electron would change its behavior Light has properties of both waves and particles Wave-like Property of Light Light is a form of Electromagnetic Radiation Electromagnetic radiation is a form of energy that travels through empty space at 3.00 x 10 8 m/s = c = speed of light. (186,000 mi/s) Electromagnetic radiation has a magnetic field component and perpendicular to that an electric field component. c = ; speed of light= (wavelength)(frequency)

3 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 3 Visible light: The components of white light R O Y G B I V separate when passed through a prism Practice 1: c = ; wavelength in meters (, frequency or hertz in 1/s ( ), speed of light (c=3.00 x 10 8 m/s) a) Solve for the frequency of green light with a wavelength of 540nm. b) A radio signal has a frequency of 100.7MHz. Solve for its wavelength. Hz = s -1

4 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 4 Particle-like Property of Light (photons) Light is particle-like: Light has photons of quantized energy. Photoelectric Effect: Emission of an electron from a metal surface caused by shining light (electromagnetic radiation) of certain minimum energy. The current increases with increasing intensity of radiation. This experiment from Albert Einstein led to the idea of photons and E = h hc/ (Planck s constant = h = x J s). The shorter the wavelength the higher the energy. The electromagnetic spectrum is continuous starting with the low energy, long wavelength, low frequency radio waves and increasing in energy through microwaves, IR, VIS (ROYGBIV), UV, Xrays, and gamma rays which are high energy, short wavelength, high frequency). Visible light ranges around 750nm red - 400nm violet

5 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 5 Emission Spectra: the emission of light from excited gas atoms. This observation led to Bohr s Atomic Model to explain it. Mercury (blue) Hydrogen (pink) Emission Spectra are like fingerprints, each element or compound has a unique emission spectrum. This allows scientist to investigate what matter makes up the stars without going to the sun and bringing back a sample to test.

6 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 6 Problems with Rutherford s nuclear atom: Electrons are moving charged particles. According to the known classical physics at that time, moving charged particles give off energy, therefore electrons should constantly be giving off energy; they should glow, lose energy, crash into the nucleus, and the atom should collapse; but it doesn t! Neils Bohr ( ) 1913 Bohr s Model (electrons move around the nucleus in circular orbits): Emission spectra of hydrogen gave experimental evidence of quantized energy states for electrons within an atom. Quantum Theory: Explains the emission and absorption spectra 1. An atom has discrete energy levels (orbits) where e - may exist without emitting or absorbing electromagnetic radiation. 2. An electron may move from one orbit to another. By doing so the electromagnetic radiation is absorbed or emitted. 3. An electron moves in circular orbits about the nucleus and the energy of the electron is quantized.

7 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 7 Hydrogen Series: n final = 1, Lyman series (UV) n final = 2, Balmer series (Visible) n final = 3, Paschen series (IR) Balmer series which gives visible light is the only one to remember Problem: Bohr s theory is limited. It only explained spectra for an item with one electron, the element H,

8 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 8 Bohr s atomic theory only worked for 1 electron systems, to explain further the next theory involves orbitals not orbits Quantum Mechanical Model of the Atom (orbitals): Replaced Bohr model Electrons can be treated as waves or particles (just as in light) Weakness: Heisenberg s Uncertainty Principle. It is impossible to determine both the momentum and position of an electron simultaneously. This means that the more accurately you know the position of a small particle, such as an electron, the less you know about its speed (momentum) and vice-versa Quantum Mechanical Model: Use 90% probability maps (orbitals not orbits) volume of space. 1. Electrons have quantized energy states (orbitals). 2. Electrons absorb or emit electromagnetic radiation when changing energy states. 3. Allowed energy states are described by four quantum numbers which describe size, shape, position, and spin respectively.

9 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 9

10 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 10 Orbitals: The orbitals in an atom are all centered around the same center nucleus. As orbitals are filled, some seem to overlap, creating an electron cloud type appearance. Orbitals are quantized and exist at discrete energy levels. Shapes: within an energy level s-1, p-3, d-5, f-7 Amounts: Maximum 2 electrons in any one orbital, Maximum 2n 2 electrons for any n level n= 1; 2 electrons in 1s 2 n = 2; 8 electrons 2s 2 2p 6 n = 3; 18 e - s 3s 2 3p 6 3d 10 Electron configuration: Aufbau Principle: method that fills electrons in the ground state Ground state electrons fill lowest energy and up Excited state occurs when electrons have jumped to higher states. Hund s Rule of Multiplicity: Electrons fill up singly with the same spin within an energy sublevel before doubling up Pauli s Exculsion Principle: No two electrons in the same atom may have the same four quantum numbers (they cannot occupy the same orbital with the same spin)

11 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 11 Electrons: core [noble gas element] pseudocore valence electrons Isoelectronic series: Atom and ions having the same number of electrons Electron Configurations: Long form: Condensed form: Condensed configurations start with a noble gas in brackets. N [He]2s 2 2p 3 Energy vs Size when d, or f orbitals are involved. Orbital diagram: labels and a drawing of the electronic configuration in order to show each orbital as filled with two electrons, half filled with one electron or empty for each energy sublevel. Paramagnetic: (weakly attracted to a magnetic field) The electron configuration has unpaired electrons. Diamagnetic: (weakly repelled by a magnetic field) All electrons are paired.

12 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 12 Electron Configuration of atoms and ions Understand the general trends. Some exceptions (Cr, Mo s 1 d 5, Cu, Ag, Au s 1 d 10 ), Nb is 5s 1 4d 4, Ce, Th, exceptions are not tested in Introductory chemistry Cations lose valence p,s orbital electrons before d orbital electrons. Sn Sn +2 Sn +4 Ions and their Electron Configurations: Metals: Give up electrons and become positive cations Nonmetals: Accept electrons and become negative anions Main group atoms give up or accept electrons toward the goal of establishing a core (noble gas) electron configuration. Metals will first lose the valence p electrons before the valence s electrons. Transition metals and Inner Transition metals first lose the largest s electrons, then the d electrons and for inner transitions last are the f electrons.

13 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 13 Try this: Predict the ground state short electron configuration for each. Ca Ca +2 F F - Ag Ag +1 Fe Fe +2 Fe +3 Periodic Table: The periodic table gives us information in an organized manner. Patterns and properties can be predicted following groups and periods. Effective Nuclear Charge: Negatively charged electrons are attracted to the positively charged nucleus. The attraction depends on the net nuclear charge acting on an electron and the average distance between the nucleus and the electron. Z eff (effective nuclear charge), is smaller than the total charge of the nucleus. Z eff = Z total - S screening constant S is close to the # of core electrons, (i.e. for Na, 10 core electrons, 1 valence electron. The Z eff = = +1 for the valence electron 3s 1 ) Periodic Trends: Size: Atomic radii generally increase from right to left top to bottom Along a period the Z eff increases left to right pulling in electrons closer to the nucleus and causing the atoms to decrease in size. Vertically, the size of the orbitals (quantum number n) increases from top to bottom.

14 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 14 Size: Ionic radii and isoelectronic series. Cations lose electrons and are smaller than the original atom Anions gain electrons and are larger than the original atom In an isoelectronic series (all have the same number of electrons, same electron configuration) the size increases as the charge of the nuclei decreases. (smallest Sr +2, Rb +1, Kr, Br -1, Se -2 largest) Ionization Energy (Energy required to remove the outermost ground state electron, endothermic): Ionization Energy generally decreases from right to left; top to bottom of periodic table. The small nonmetals require the highest ionization energy; they do not want to lose electrons Large metals have lowest ionization energy, they want to lose electrons and become positively charged cations. Oxidation is a term for an atom losing electrons and increasing its charge (its charge is also known as its oxidation state). Electron Affinity (Energy given away when adding an electron, exothermic): The greatest negative value (most preferred) electron affinity is for fluorine, (small nonmetals, ignore noble gases they do not want to add electrons) decreases from right to left; top to bottom of periodic table.

15 C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e 15 Metallic Character: Metallic character increases from right to left; top to bottom of periodic table. Nonmetal character increases from left to right bottom to top. Electronegativity (Ch 10): The ability of an element to attract electrons within a covalent bond. Electronegativity increases from left to right; bottom to top of periodic table. It does not include the noble gases, so the strongest one is F.

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