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Periodic Law

Periodic Law. Chemistry I. Group IA alkali metals Group IIA alkaline earth Metals Group IIIB-IIB transition metals Inner Transition Metals Lanthanide Series - also called the rare earth elements . Actinide Series.

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Periodic Law

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  1. Periodic Law Chemistry I

  2. Group IA alkali metals Group IIA alkaline earth Metals Group IIIB-IIB transition metals Inner Transition Metals Lanthanide Series - also called the rare earth elements. Actinide Series

  3. Group VIIIA noble gases or inert gases Group VIIA - halogens Bright Green - nonmetals Purple - metalloids or semiconductors Light Blue - other metals

  4. The Periodic Law states that “The physical and chemical properties of the elements are periodic functions of their atomic numbers. 1. the properties of the elements go through a pattern of change 2. elements of similar properties occur at certain intervals, provided the elements are arranged in a periodic table in order of increasing atomic number.

  5. Periodic Chart in Block Form http://www.gordonengland.co.uk/elements/blocks.htm

  6. Periodic Trends Atomic Radius – ½ the distance between 2 adjacent nucleii in the crystalline forms of elements and/or in the molecules of gaseous elements. Pattern of radius increase is as follows: Top to Bottom Right to left

  7. Periodic Trends of Ionic Radii • Ionic radii are formed by the loss of electrons forming positive ions (cations) or by the gain of electrons forming negative ions (anions). Metals usually form cations which, due to the loss of electrons are smaller than the atomic radius from which they were derived. Nonmetals usually form anions which, due to the gain of electrons, are larger than the atomic radius from which they were derived. • Radii increase from top to bottom • Radii increase from right to left • See chart on next frame.

  8. Source: http://picsicio.us/image/5e6a3797/

  9. Ionization Energy • The amount of energy needed to remove an electron from an atom. Low ionization energy is characteristic of metals and high ionization energy is characteristic of nonmetals. For any given atom, the ionization potential increases as successive electrons are removed. • Within groups, 1st ionization potential increases with decreasing atomic number – Bottom to Top. • Within periods, 1st ionization potential decreases with increasing atomic number – Left to Right

  10. http://www.grandinetti.org/Teaching/Chem121/Lectures/IonizationEnergy/index.htmlhttp://www.grandinetti.org/Teaching/Chem121/Lectures/IonizationEnergy/index.html

  11. Electron Affinity • The energy change that is required for a neutral atom to acquire an electron. • Negative electron affinities indicate that energy is emitted when an electron is accepted indicating an exothermic reaction. • Positive electron affinities indicate that energy is absorbed when an electronic is accepted indicating an endothermic reaction.

  12. Electron Affinity Patterns • Within a family of light, active elements, electron affinity tends to increase with decreasing atomic number – BOTTOM TO TOP. • Within a family of transition or rare earth metals, electron affinity tends to increase with increasing atomic number – TOP TO BOTTOM. • Within a given period of elements, electron affinity tends to increase with increasing atomic number – LEFT TO RIGHT.

  13. Source: http://www.grandinetti.org/Teaching/Chem121/Lectures/ElectronAffinity/index.html High Affinity

  14. Electronegativity • Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. Fluorine has the highest electronegativity (4 on the Pauling scale). Cesium and Francium are only 0.7. We will discuss electronegativity later when we discuss bonding types. Source: http://www.chemguide.co.uk/atoms/bonding/electroneg.html

  15. Periodic Trends of Valence Electrons • Valence Electrons – found in the outermost energy level of an atoms that can be transferred or shared with other atoms. • Within a given group, the number of valence electrons tend not to vary from one element to the next. (Same # of valence e’s in a group.) • Within a given period, the number of valence e’s increase with increasing at. number – LEFT TO RIGHT. Note “d” & “f” block elements have only 2 valence electrons.

  16. Source: http://www.chemprofessor.com/periodicqm_files/image005.gif

  17. Chemical Activity is dependent on: 1. Size of atom or ion 2. Ionization potential 3. Electron affinity 4. # of valence electrons • Most active metal: Francium (lower left) which has a large atom/ionic radius, low ionization potential, low electron affinity, & has the least number of valence electrons. • Most active nonmetal: Fluorine (upper right) which has a small atomic/ionic radius, high ionization potential, high electron affinity; & greatest number of valence electrons without achieving the stable octet of the noble gases.

  18. Ions – charged particles • Ions form when a neutral atom loses or gains electrons. Atoms become more stable (full outer shell) when they form ions. • Losing yields a positive ion called a cation. • Gaining yields a negative ion called an anion.

  19. A Sodium ion has 11 protons +11 charge 12 neutrons 0 charge 10 electrons -10 charge Total charge +1 A cation!! A Chlorine ion has 17 protons +17 charge 18 neutrons 0 charge 18 electrons -18 charge Total charge - 1 An anion!!

  20. Source: http://www.grandinetti.org/Teaching/Chem121/Lectures/ElectronAffinity/index.html

  21. Source: http://www.grandinetti.org/Teaching/Chem121/Lectures/ElectronAffinity/index.html Affinity

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