Chemistry AI Chapter 6 A. Organizing elements 1. Dobereiner (1829) He arranged the elements into triads, groups of three elements, based on similar chemical properties. Ex. Cl 2, Br 2, I 2

2. D. Mendeleev (1869) 1. Father of the modern periodic table 2. He organized the 60 known elements by increasing atomic mass. 3. Mendeleev left spaces for the unknown elements. He even predicted the properties of these unknown elements. (Ga and Ge) 4. Mendeleev realized that some elements could not be organized by atomic mass, because their chemical properties placed them elsewhere. Ex. Te and I are reversed

3. H. Moseley He determined the atomic number for each element. The periodic table is now organized by atomic number NOT atomic mass. B. Periodic Table 1. Periodic Law The elements are arranged according to similarities in their properties. When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties

a. Period b. Group/ Family A row on the periodic table. There are 7 that correspond to the 7 energy levels of an atom. These are the columns on the periodic table Elements in the same Group have similar chemical properties C. Lay out of the Periodic Table 1. Metals 1. Found to the left of the stairs 2. good conductors 3. high luster (shiny) 4. malleable- able to be hammered into sheets 5. ductile- able to be drawn into wire 6. give away valence electrons resulting in a positively charged cation (outer shell electrons)

2. Nonmetals 1. found to the right of the stairs 2. dull 3. poor conductors 4. brittle 5. may gain electrons resulting in a negatively charged anion or share valence electrons (no charge) 6. many are gases at room temperature 3. Metalloids1. found on the stairs 2. have properties of both metals and nonmetals, depending on the condition Ex. Silicon is a very poor conductor but if a small amount of boron is added to it, it becomes a very good conductor

D. Classifying Elements 1. Group 1 (1A) 2. Group 2 (2A) 3. Group 17 (7A) 4. Group 18 (8A) Alkali Metals- very reactive Alkaline Earth Metals- less reactive Halogens-combine with metals to make salts Noble gases. Mostly inert because of a full outer energy level (8 valence electrons; except He) 5. Representative Elements (Labeled A) (Tall stacks) They display a wide range of physical and chemical properties. The group number tells you how many electrons are in its outer shell (valence electrons)

6. Group 3-12 (Labeled B) (Short stacks) Transitions Metals Valence electrons are variable Often form colored pigments in compounds 7. Bottom Group Inner Transition Metals

E. Periodic Trends 1. Atomic Size a. shielding effect Atomic size increases from top to bottom within a group, mostly due to the shielding effect. The increase in the number of occupied orbitals shields the outer electrons from the attraction of the nucleus Atomic sizes decreases from left to right in a period, shielding does not affect the atom because all have the same orbitals occupied. The pull of the nucleus due to increased protons causes the atom to decrease. Remains constant from left to right. Increases from top to bottom.

2.Ionization Energy a. Ions b.Group c.Period Energy required to remove an electron from an atom An atom or group of atoms with a charge. Atoms become charged by gaining or losing electrons. Decreases as you go down the table Increases as you go across the table

3. Electro- negativity The ability of an atom of an element to attract electrons when in a compound. Electronegativity decreases from top to bottom within a group and tend to increase from left to right across the period. Why are noble gases not in the table?

Chemistry AI Chapter 7 and 8 A. Ions (7.1) 1. Valence Electrons Atoms or groups of atoms with a charge; they are formed when atoms gain or lose electrons a. Electrons in the highest occupied energy level of an atom b. They determine the chemical properties of an atom because they are usually the only electrons that participate in a chemical bond c. Can be determined by the group number of the representative elements (Except He) 2. Octet RuleEach representative element wants 8 valence electrons to achieve the electron configuration of the nearest noble gas. (Except Li, Be, B want 2 valence electrons like He due to their small size)

a. Cationsa. Positive oxidation number b. Metals tend to lose their valence electrons producing cations Ex. Sodium cation Na+ Magnesium cation Mg2+ c. Oxidation number for transition and inner metals can not be predicted based on the group number due to the complicated electron orbitals, they may have multiple oxidation states (although always positive) b. Anions a. negative oxidation number b. nonmetals may gain valence electrons creating anions ex. chloride anion Cl- oxide anion O2- (name ends in –ide)

B. Ionic Bonds (7.2) 1. Ionic Compounds (page 194) a. ionic bond Made of a metal (cation) and a nonmetal (anion) Although composed of charge particles, ionic compounds are electrically neutral. a. Electrostatic attraction between cation and anion (like a magnet) b. Formed by metal transferring electrons to the nonmetal

b. Chemical formula (Page 195) Shows the kinds and numbers of atoms in the smallest representative unit of a substance Ex. NaCl, MgCl 2 In ionic compounds the ratio of ions is reduced to the lowest possible whole number (i.e. formula unit) Ex. Mg 4 Cl 8, Mg 200 Cl 400 NO MgCl 2 Yes 1. Subscript 2. Superscript The small number in a chemical formula written down low that indicates the number of a particular atom The small number written up high that indicates the charge of an ion.

Ca F Al O + + Ca + Al N +.

d. properties of ionic compounds most ionic compounds are crystalline solids at room temperature the ions are arranged in repeating three- dimensional patterns called crystals ions are locked in place. This makes the solid very brittle generally have high melting points and boiling points ex. NaCl melts at 800 o C conduct electric currents when dissolved

C. Covalent Bonds (8.1) 1. Molecular Compounds a. Covalent Bonds Made of two nonmetals 1. Sharing electrons; shared electron spends time around both atoms; no ions involved 2. electrons are always shared in pairs with each atom donating one electron

b. Molecular formula c. properties of molecules 2. Diatomic molecules Tells you what is in one molecule Ex. H2O molecules are individual units and are not strongly held together compared to ions molecules can easily break away from each other they can slide past each other molecular substances are usually liquidsor gases at room temperature (low melting and boiling points) a molecule that consist of two of the same atoms Br 2, I 2, N 2, Cl 2, H 2, O 2, F 2