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Chapter 14: Periodic Trends …and naming ions (chapter 6) The Modern Periodic Table Review: Energy Levels • Principle quantum numbers (n) = energy level – Lower the number, lower the energy Sublevels – s, p, d, f, g Put it all together… A new way to write electron configurations A new way to write electron configurations Use the Noble gas abbreviation, write the electron configurations for: • Na: • Cl: • W: • Sn: Organizing the Table • Noble Gases – – Filled outermost s and p orbitals (s2p6) • Representative Elements- (s and p block) – Partially filled outermost s and p orbitals • Transition Metals (d block) – Outermost s and nearby d orbitals contain electrons • Inner Transition Metals – Outermost s and nearby f contain electrons Representative Elements • • • • • Sometimes called “group A elements” 1A = alkali metals 2A = alkaline earth metals 7A = halogens 8A = noble gases • The group number equals the number of electrons in its outermost energy level (this will be more important later on…) Metals and Nonmetals • From here on out you need to identify whether an element is a metal or nonmetal (FAST and accurate). Other things you need to know… (Ch. 6) • When atoms gain or lose electrons they form IONS (positively or negatively charged atoms). • Charge determined by difference between p+ and e- • Positive ions = cations (lose valence e-) – Metal atoms lose valence e- • Negative ions = anions (gain valence e-) – Nonmetal atoms gain valence e- Cations • Formed when a metallic atom LOSES electrons. • Examples: Calcium and Magnesium calcium ion magnesium ion • You can determine how many electrons are lost based on location on the periodic table. These must be memorized. – Group 1A, 2A, Al, Ag, Zn Naming Ions • How to name: name of element + “ion” • Examples: • Cations – Aluminum (Al) – Sodium (Na) : : Aluminum Ion (Al3+) Sodium Ion (Na+) Transition Metals… • Can form different charges (you can’t memorize) • Here is how you know the charge: – They are metals, so are cations. – Roman Numerals (I, II, III, IV, V, VI, VII) indicate how many e- lost. – Copper (II) : two val e- : Cu2+ – Copper (I) : one val e- : Cu+ Anions • Non-metal atoms that gain electrons become anions. • Examples: Bromine and Nitrogen bromide ion nitride ion • You need to MEMORIZE the common charges of the anions to be successful for the rest of the school year… Naming Anions • Anions change the ending of the element – unlike cations • Stem-ide ion • Examples: • Chlorine = chloride ion • Oxygen = oxide ion • Anions – Chlorine (Cl) – Oxygen (O) : : chloride ion (Cl-) oxide ion (O2-) Ion Size • Cations are smaller in size than the neutral element. • Anions are larger in size than the neutral element. Polyatomic ions • Polyatomic = many atoms • Ions = charged • You will get a list of 10 polyatomic ions. You must memorize the name and formula and be able to recall them at any time after the first test (i.e. I won’t feel guilty if there is a pop quiz). Now on to trends… • There is no secret to success this chapter other than to memorize the following trends… Types of Trends • Periodic Trends: Trends across a period (row) of the periodic table. • Group Trends: Trends down a group (family) of the periodic table Nuclear Charge (+ in nucleus) • Periodic: Nuclear charge increases as you go left to right across a period. • Group: Nuclear charge increases as you go down a group. Atomic Radii and Size • Periodic: Atomic Radii and size decrease as you go L to R across the table. – Same principle energy level (n) – Add p+ and e-, increase nuclear charge, pulls in orbitals closer to the nucleus • Group: Atomic Radii and size increase as you go down a group – Electrons being added to outer orbital (increasing principle energy level) Ionization Energy • Ionization Energy (IE): The energy required to remove an electron from a gaseous atom. – Remove 1st electron = 1st IE – Remove 2nd electron = 2nd IE Trends in Ionization Energy • Periodic: IE generally increases as you move L to R across the period. – Harder to remove an electron as you go L to R because of greater attraction to nucleus – Shielding effect - • Group: 1st IE generally decreases as you go down a group. – Atom gets bigger, outermost e- farthest from nucleus, easy to be removed. Electronegativity • Electronegativity: The tendency for atoms to attract electrons when they are chemically combined. – Stronger attraction. Electronegativity Trends • Periodic: Electronegativity increases as you go L to R across the period – Elements want to be like noble gases! • Group: Electronegativity decreases as you go down a group. The most electronegative element is Fluorine. Summary Decreasing Atomic Radius Increasing Electronegativty Increasing Atomic Radius Decreasing Electronegativity Decreasing Ionization Energy Increasing Ionization Energy