chem notes week 10
chem notes week 10 CHE 106 - M001
Popular in General Chemistry Lecture I
CHE 106 - M001
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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