Chapter 4 notes
Chapter 4 notes CH 101
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Popular in Chemistry
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This 8 page Class Notes was uploaded by Lauren Dutch on Wednesday September 14, 2016. The Class Notes belongs to CH 101 at University of Alabama - Tuscaloosa taught by Dave Nikles in Fall 2016. Since its upload, it has received 57 views. For similar materials see General Chemistry in Chemistry at University of Alabama - Tuscaloosa.
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Date Created: 09/14/16
Ch4 Periodic Properties of the Elements I. Low density atoms result in low density metals A. Density increases as you move down a column 1. Mass of each successive element increases more than volume; increase in mass is greater than increase in volume, resulting in increased density B. Periodic properties are properties predicted by an element’s position on the periodic table C. Structure determines properties II. Finding patterns using the periodic law and the periodic table A. Modern periodic table is credited to Dmitri Mendeleev 1. Periodic law: When the elements are arranged in order of increasing mass, certain sets of properties recur periodically B. Henry Moseley listed elements according to atomic number C. Main group elements have properties that can be predicted based on position in periodic table D. Transition elements and inner transition elements have properties that are less predictable based on position in periodic table E. Family/group is a column in the main group region that have similar properties F. The quantum mechanical theory determines why elements have certain properties while the periodic table just displays those properties III. Electrons occupy orbitals according to electron configuration A. Electron configuration shows how electrons occupy each specific orbital in an atom B. Electrons occupy the lowest energy orbital first C. Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers 1. Atoms with the same n, l, and m numLers must have different spin quantum numbers (-1/2 or +1/2) a. There can only be two electrons in each orbital and they must have opposite spin D. Coulomb’s Law states that the potential energy (F) of two charged particles depends on their charges (q an1 q ) an2 their separation (r) 1. A negative energy indicates an attraction between the two particles 2. A positive energy indicates a repulsion between the two particles 3. Energy decreases as particles get farther apart E. Shielding is when one electron is repulsed by other electrons so that it does not receive the full effects of the nuclear charge F. Penetration is when the outer electron goes into the region occupied by the inner electrons where it experiences a greater nuclear charge and has a lower energy G. Filling the orbitals 1. Aufbau principle: electrons occupy the lowest energy orbitals available when the atom is in its ground state 2. Hund’s rule: when filling degenerate orbitals, fill them singly first, with parallel spins 3. Remember orbitals cannot hold more than two electrons 4. S sublevel has only one orbital and can only hold 2 electrons 5. P sublevel has 3 orbitals and can hold 6 electrons 6. D sublevel has 5 orbitals and can hold 10 electrons 7. F sublevel has 7 orbitals and can hold 14 electrons IV. Electron configurations, valence electrons, and the periodic table A. As you move down a column, the number of electrons in the outermost energy level stays the same B. Valence electrons are the most important in chemical bonding; valence electrons are found in the outermost energy levels C. Core electrons are the electrons found in complete principal energy levels D. The periodic table can be seen as 4 different orbital blocks 1. Group number equals the number of valence electrons, e.g. chlorine in group 7A has 7 valence electrons 2. Row number is equal to the highest principal level, e.g. chlorine is in row 3 so its highest n level is n=3 V. Electron configuration of an element relates to its properties A. Noble gases are inert because they have 8 valence electrons (except helium has 2) 1. Full quantum level creates low energy level 2. Systems with high energy want to have low energy so all elements aim to attain a noble gas configuration B. Metals are on the lower left side and middle of periodic table 1. Conductivity 2. Malleability 3. Ductility 4. Shiny 5. Lose electrons when they undergo chemical changes C. Nonmetals lie on the upper right side of the periodic table 1. Poor conductivity 2. Gain electrons when they undergo chemical changes D. Metalloids 1. On the zigzag diagonal line diving metals and nonmetals 2. Semiconductors have intermediate conductivity E. Families of elements 1 1. Alkali metals all have outer electron configuration of ns a. Lose an electron to form ions with a +1 charge 2. Alkaline earth metals all have outer electron configurations of ns2 a. Lose two electrons to from ions with a +2 charge 2 5 3. Halogens all have an outer electron configuration of ns np a. Gain an electron to form ions with a -1 charge VI. Periodic trends in the size of atoms and effective nuclear charge A. Atomic radius can be determined in two ways 1. Van der Waals radius/non bonding radius 2. Covalent radius/bonding radius B. Atomic radius trend: increase down a column, decrease across a period C. Effective nuclear charge is the average charge experienced by an electron ZEFF= Z – S Effective nuclear charge = actual nuclear charge – charge screened by electrons 1. Core electrons shield electron in the outermost principal energy level but valence electrons do not shield each other 2. Effective nuclear charge increases across a period because outermost electrons are more attracted to nucleus and radius is smaller VII. Ions: electron configurations, magnetic properties, ionic radii, and ionization energy A. For anions, add an extra electron to the electron configuration B. For cations, subtract an extra electron from the electron configuration 1. When removing electrons, remove from the HIGHEST N LEVEL first, even if this doesn’t correspond with the written order of orbitals C. Atoms with unpaired electrons is paramagnetic D. Atoms with all electrons paired are diamagnetic E. Ionic radii 1. Cations are much smaller than their neutral atoms 2. Anions are much larger than their neutral atoms F. Ionization energy (IE) is the energy required to remove an electron from the atom or ion in the gaseous state 1. Always positive 2. First IE is the energy required to remove the first electron 3. Second IE is the energy required to remove the second electron; third IE removes the third; etc. 4. Trends in ionization energy: decreases down a column, increases across a period VIII. Electron affinities and metallic character A. Electron affinity (EA) is the energy change associated with the gaining of electron by an atom in the gaseous state 1. Usually negative, exothermic 2. For main group elements, EA becomes more negative across a period 3. The only group with a trend is 1A- EA becomes more positive as you go down the column B. Metallic character 1. Metallic character decreases across a period 2. Metallic character increases down a column
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