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PERIODIC PATTERNS Unit 3 – Periodic Table Lesson Essential Question: What patterns exist on the periodic table? INCREASES METALLIC TREND ATOMIC RADIUS Radius is the distance from the center of the nucleus to the “edge” of the electron cloud. Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is 1 x 10-10 m. ATOMIC RADIUS BROMINE = Br2 Since a cloud’s edge is difficult to define, scientists use define covalent radius, or half the distance between the nuclei of 2 bonded atoms. 2.86 Å 1.43 Å 1.43 Å ATOMIC RADII TRENDS WHY? DOWN A FAMILY OR GROUP INCREASES As you go down a family the number of energy levels increases making the radius larger. ATOMIC RADII TRENDS WHY? ACROSS A PERIOD DECREASES As you go across a period the number of protons increases, (nuclear charge) pulling the electrons in tighter making the radius smaller. IONS - remember Metals Lose electrons becoming positive. Calcium (Ca) Loses 2 electrons becoming Ca+2 and [Ar] Noble gas Configuration. (Octet Rule) Nonmetals Gain electrons becoming negative. Chlorine (Cl) Gains one e- becoming Cl-1 and [Ar] Noble gas configuration. (Octet Rule) IONS – How can I remember? Metals Nonmetals This is Cat-ion - CATION This is Ann ion - ANION He is a “plussy” cat! She is unhappy and negative. IONIC RADII TRENDS WHY? DOWN A FAMILY OR GROUP INCREASES As you go down a family the number of electron shells increases making the radius larger. IONIC RADII TRENDS WHY? ACROSS A PERIOD DECREASES then INCREASE For the metals the nuclear charge is greater than then number of electrons pulling them in tighter making the radius smaller. At the nonmetals the radius gets larger because the ion has gained electrons. METALLIC ATOM AND ION COMPARISON NONMETALLIC ATOM AND ION COMPARISON Why do the Noble Gases not have an ionic Radius? ATOM AND ION COMPARISON Why does Hydrogen not have an ionic Radius? Shielding Effect As more electrons are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus. The effective nuclear charge on those outer electrons is less, and so the outer electrons are less tightly held Example of Shielding Effect Ionization Energy The energy required to remove an electron from an atom. (measured in kilojoules, kJ) Why? • Closer to nucleus (more +) • Electrons less likely to be removed • Requires more energy to form ion • Less shielding INCREASES IONIZATION TREND IONIZATION ENERGY The larger the atom is, the easier its electrons are to remove. (Why?) Ionization energy and atomic radius are inversely proportional. Ionization energy is always endothermic, that is energy is added to the atom to remove the electron. IONIZATION TREND NCREASES Why? • Elements in alkali metals have 1 valence electron so what to remove that electron, they therefore take the least amount of energy to remove an electron INCREASES Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. Why? • Closer to nucleus (more +) so electrons are more attracted INCREASES Electronegativity http://www.thecatalyst.org/electabl.html Electronegativity Why? • Elements in halogens only need 1 more electron to have a full valence shell so are MOST likely to attract electrons. As you move to left elements are more likely to LOSE electrons. INCREASES http://www.thecatalyst.org/electabl.html In Summary…. Electronegativity Electronegativity