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chem notes week 10

by: Andrea Scota

chem notes week 10 CHE 106 - M001

Andrea Scota
GPA 3.7

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chem notes for week 10
General Chemistry Lecture I
R. Doyle
Class Notes
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This 5 page Class Notes was uploaded by Andrea Scota on Sunday November 8, 2015. The Class Notes belongs to CHE 106 - M001 at Syracuse University taught by R. Doyle in Fall 2015. Since its upload, it has received 47 views. For similar materials see General Chemistry Lecture I in Chemistry at Syracuse University.


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Date Created: 11/08/15
Pink- mentioned in class Chem Notes Week 10 TEXTBOOK CHAPTER 7 (continued sections 7.5-7.8) Periodic Properties of the Elements Electron affinity (7.5)  All ionization energies for atoms are positive. Energy must be absorbed to remove an electron  It is more favorable the more negative it is, the greater the electron affinity  Electron affinity: the energy change upon adding an electron to an atom in the gas phase (forming anion), measures the attraction of the added electron  Exothermic  Most atoms, energy is released when an electron is added (mostly negative)  Ionization energy measures the energy change when an atom loses an electron; whereas electron affinity measures the energy change when an atom gains an electron  The greater the attraction between he atom and the added electron, the more negative the number  The more negative the electron affinity the more stable the negative ion  In general, electron affinities become more negative as we proceed from left to right across the table  Halogens have the most negative electron affinity  Electron affinities of noble gases are positive because the new electron would have to occupy a new, higher energy subshell Metals, Nonmetals, and Metalloids (7.6)  Elements can be broadly grouped as metals, nonmetals or metalloids  S/P block are main group elements  Metals occupy the left side and middle  Nonmetals appear in the upper right section  Metalloids occupy a narrow metal band  Metallic character the more an element exhibits physical/chemical properties of metals, the greater this is o Generally, increases as we go down a group and decreases as we move right across the periodic table  METALS: o Have characteristic luster o Good conductors of heat/electricity o When metals react with nonmetals, the metals are oxidized to cations and ionic substances are usually formed o Most metal oxides are basic; react with acids to form salts and water (function as bases)  NONMETALS: o Lack metallic luster o Poor conductors of heat/electricity o Several are gases at room temperature o Compounds composed entirely of nonmetals are usually molecular o Usually form anions in their reactions with metals o Nonmetal oxides are acidic; they react with bases to form salts and water o They tend to gain electrons in reactions to get to noble gas configurations; they have very negative electron affinities  METALLOIDS: o Have properties that are intermediate between those of metals and nonmetals Trends for Group 1A and 2A Metals (7.7)  Periodic properties of the elements can help us understand the propertied of groups of representative elements  The alkali metals (group 1A) are soft metals with low densities a low melting points o All have characteristic properties like silvery and metallic luster, and high thermals/electrical conductivity  Lox density/low melting points vary in fairly regular way with increasing atomic number o Visual trends as we move down; increase in atomic radius and decrease in ionization energy  Have lowest ionization energy, meaning they are very reactive (towards nonmetals), easily losing outer s shell electron to form 1+ ions  Only exist in nature as compounds  React with hydrogen in nature to form hydrides, with sulfur to form sulfides  React vigorously with water and are very exothermic  The alkaline earth metals (group 2A) are harder and denser and have higher melting point  Very reactive towards nonmetals, but not as reactive as alkali metals  Readily lose their 2 s electrons forming 2+ ions  React with hydrogen to form ionic substances that contain the hydride ion H – Trends for Selected Nonmetals (7.8)  HYDROGEN: o Ionization energy if 1312 kJ/mol o Nonmetal that occurs as colorless, diatomic gas H2 (g) o Forms molecular compounds with other nonmetals in which its electron is shared with it o Water, methane, other halogens o Has ability to gain electron from metal with a low ionization energy  Oxygen and sulfur are the most important elements in group 6A  OXYGEN: o Usually found as diatomic molecule, O2 o Ozone: (O3) is an important allotrope of oxygen o Has string tendency to gain electrons from other elements, thus, oxidizing them o In combinations with metals, oxygen is usually found as oxide ion O 22and super peroxide ion O a2e sometimes found  HALOGENS: o Exist as diatomic molecules o Have the most negative electron affinities of the elements o Their chemistry is dominated by a tendency to form 1- ions, especially in reactions with metals  NOBLE GASES: o Group 8A o Exist as monatomic gases o They are very unreactive because they have completely filled s/p subshells o Only heaviest noble gases are known to form compounds and the do so only with very active nonmetals (like fluorine) TEXTBOOK CHAPTER 8 (sections 8.1-8.3) Basic Concepts of Chemical Bonding  Whenever two atoms or ions are strongly held together, we say there is a chemical bond between them, three groups: o Ionic bonds: due to electrostatic attractions between oppositely charged ions (MgO) o Covalent bonds: formed by sharing electrons in atoms o Metallic bonds: formed by electrons that are relatively free to move through metal; when metal atoms font to other metal atoms Lewis Symbols and the Octet Rule (8.1)  The formation of bonds involves interactions of the outermost electrons of atoms, their valence electrons  The valence electrons of an atom can be represented by electron dot symbols, Lewis Symbol o Consists on elements chemical symbol plus a dot for each valence electron  The octet rule: atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons (want to be a noble gas) Ionic Bonding (8.2)  Generally, result from interaction of metals on left side of periodic table with nonmetals on the right side  Results from the transfer of electrons from one atom to another, leading to 3D lattice of charged particles  Electron transfer to form oppositely charged ions occurs when one atom readily gives up an electron has a low ionization energy) and one atom readily gains an electron (has a high electron affinity)  Ionic substances usually very brittle with a high melting point and crystalline  Formation of ionic compounds are very exothermic  If 2 electrons come together, there must be a net decrease in energy; their bonded state must be more stable (lower in energy)  Attraction of electrons between ions draws them together, releasing energy and causing them to form a solid  Lattice energy: the energy required to completely separate one mole of a solid ionic compound into its gaseous ions o Increases with increasing charge on the ions and with decreasing size of ions o Endothermic o Depends on charges, sized, arrangement o E elKQ Q 1 d2(electrostatic potential energy)  K is constant 8.99 x 10 J-m/C 2  Q /Q are the charges in particles in columbs with signs 1 2  Lattice energies follow trends that of ionic radius  Born-Haber cycle is a useful thermochemical cycle in which we use Hess’s Law to calculate the lattice energy as the sum of several steps in the formation of an ionic compound Covalent Bonding (8.3)  Chemical bond formed by sharing a pair of electrons  Lewis structures: represent the electron distribution in molecules, they indicated how many valence electrons are involved in forming bonds and how many remain as nonbonding electron pairs (lone pairs)  The octet rule helps determine how many bonds will be formed between 2 atoms  The sharing of one pair of electrons produces a single bond, the sharing of two or three pairs of electrons between two atoms produces double or triple bonds o Double/triple are examples of multiple bonding between atoms o Bind length decreases as the number of bonds increase


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