1. Review for Exam 1

2. CH 1-2 Concepts to know  Classification of matter: pure substances & mixtures  Homogeneous vs Heterogeneous  Distinguish the difference between chemical and physical properties & changes  We represent uncertainty with significant figures  You do not need to memorize Sig Fig rules  Scientific Notation  Conversions within the metric system and non metric units  Temperature conversions  Density & Specific gravity  Familiarity with how compounds will be drawn  Molecular formulas  Structure of an atom: protons, neutrons, electrons  Atomic number, isotope mass number, atomic weight  Navigate the periodic table: properties shared within a group, trends, metals/metalloids/nonmetals  Determine valance electrons & draw electron dot representations  Ionization Energy & Atomic Size

3. Conversions & Equations To Memorize Unit Conversions For metric units (m, kg, s, K, mole): mega (M) 10 6 kilo (k) 10 3 centi (c) 10 -2 milli (m) 10 -3 micro (μ) 10 -6 nano (n) 10 -9 Pico (p) 10 -12 Time conversions: d  hr  m  s 1 mL = 1 cm 3 T(kelvin) = T(°C) + 273.15 Equations Density = mass / Volume d = m/V d H2O = 1 g/mL = 1 g/cm 3 Specific Gravity = density substance / density of water Coefficient: A number between 1 and 10 y x 10 x Exponent: Any positive or negative whole number

4. Elements & Molecules X = Element symbol (ie O = oxygen) A = Isotope Mass Number = # protons + # neutrons Z = Atomic Number = # protons 6 C 12.01 atomic number element symbol atomic weight (amu) = weighted average of atomic weight of isotopes Elements on the Periodic Table Molecular Formula: A x B y Ex: CH 3 O Drawing Molecules: Methane CH 4

5. Properties of Metals, Nonmetals, Metalloids MetalsNonmetalsMetalloids Metallic luster, malleable, ductile, hardness variable Conduct heat and electricity Solids at room temperature with the exception of Hg Chemical reactivity varies greatly: Au, Pt unreactive while Na, K very reactive Brittle, dull Insulators, non- conductors of electricity and heat Chemical reactivity varies Exist mostly as compounds rather then pure elements Many are gases, some are solids at room temp, only Br 2 is a liquid. Properties intermediate between metals and nonmetals Metallic shine but brittle Semiconductors: conduct electricity but not as well as metals: examples are silicon and germanium

6. Valence Electrons Example: Determine the valence electrons of Selenium (Se): 1.Find Se on the periodic table 2.Focus on just the column Se is in 3.Column number indicates number of e - Se Electron Dot Symbols: Represent the valence electrons by drawing them around the element symbol for Selenium.

7. Periodic Trends Ionization Energy INCREASING Size INCREASING

8. CHAPTER 3-4: Concepts to Know  The difference between ionic and covalent bonds  Define cations and anions  Predict cation/anion charge using the octet rule or group number  Familiar with metals with multiple potential charges (do not need to memorize)  Determine ionic compound formulas from the name of a compound or from the elements that compose it.  Criss-cross rule  Naming ionic compounds and covalent molecules  Familiar with polyatomic ions (do not need to memorize but must be able to recognize)  Draw lewis dot structures  Determine molecular geometry  Identify polar bonds  Determine dipole moment of molecules

9. Need to Memorize

10. Ionic vs Covalent Bonding Ionic Bonds result from electrostatic attraction between a cation and anion: metal-nonmetal (with the exception of NH 4 + and H 3 O + cations). Covalent bonds result from the sharing of electrons between two atoms: nonmetal-nonmetal. Li F Ionic Bonds Covalent Bonds

11. Naming HOW TO Name an Ionic Compound Step [1] Determine the charge on the cation. Step [2] Name the cation and the anion  If the cation could be multiple charges indicate the charge with roman numerals or with a –ous / -ic suffix. Step [3] Write the name of the cation first then the name of the anion HOW TO Name a Covalent Molecule Step [1] Name the first nonmetal by its element name and the second using the suffix “-ide.” Step [2] Add prefixes to show the number of atoms of each element.

12. Predicting Cations & Anions the anion charge = 8 – group number the cation charge = the group number

13. Octet Rule only 6 e − on B B F FF 10 e − on P 12 e − on S S O OHHO O P O OHHO OH The octet rule: a main group element is especially stable when it possesses an octet of e − in its outer shell. octet = 8 valence e − Exceptions (need to memorize):

14. Ionic Compound Formulas HOW TO Write a Formula for an Ionic Compound Step [1] Identify which element is the cation and which is the anion. Step [2] Determine how many of each ion type is needed for an overall charge of zero.  When the cation and anion have different charges, use the ion charges to determine the number of ions of each needed. Step [3] To write the formula, place the cation first and then the anion, and omit charges.

15. Lewis Dot Structures Step [1] Step [2] Step [3] Arrange the atoms next to each other that you think are bonded together. Place H and halogens on the periphery, since they can only form one bond. Count the valence electrons. The sum gives the total number of e − that must be used in the Lewis structure. For each atom the number of bonds = 8 – valence electrons. Arrange the electrons around the atoms. Place one bond (two e − ) between every two atoms. Use all remaining electrons to fill octets with lone pairs, beginning with atoms on the periphery. NH 3 NHHH Nitrogen has 5 valence electrons, so it will have 8 – 5 = 3 bonds. Hydrogen will have 2-1 = 1 bond. There are 8 total valance electrons HNH H 1 lone pair: 2 3 bonds: 6 Total e-8 = total valence e-

16. Resonance Structures Resonance structures exist when there are multiple lewis dot structures with different electron arrangements with the same connectivity between atoms. Resonance structures help us understand delocalization (spreading) of charge within a molecule that stabilizes the anion or cation. Other Examples: CO 3 2- and O 3

17. Molecular Shape

18. Periodic Trend: Electronegativity Electronegativity INCREASING

19. Polarity 1.Assess the relative electronegativity of atoms bonded together, if there is a difference it is a polar bond. 2.Indicate polar bonds with δ+ / δ - or 3.If polarity of bonds does not cancel draw the overall dipole moment of the molecule using Electron density is disproportionately distributed over the molecule. Above red indicates partial negative charge, or greater electron density, and blue indicates partial positive charge. Effectively oxygen is hogging the electrons

20. CH 5 Concepts to know 20 Smith. General Organic & Biolocial Chemistry 2nd Ed.  Define a chemical reaction  Correctly write a chemical reaction  Balance reactions by inspection  Calculate molecular mass for any compound or molecule  Apply mole ratios within molecules and between molecules.  Solve stoichiometry problems  Convert between mass and moles  Identify limiting reagent  Calculate percent yield  Identify reduction and oxidation equations and pick out the compound being reduced or oxidized

21. Need to Memorize 21 Smith. General Organic & Biolocial Chemistry 2nd Ed. 6.02 x 10 23 is Avogadro’s number. Oxidation is the loss of electrons from an atom. Reducing agents are oxidized Reduction is the gain of electrons by an atom. Oxidizing agents are reduced.

22. Smith. General Organic & Biolocial Chemistry 2nd Ed. 22 aA ( physical state ) + bB ( state )  cC ( state ) + dD ( state ) Writing and Balancing Equations HOW TO Balance a Chemical Equation Step [1] Write the equation with the correct formulas. The subscripts in a formula can never be changed to balance an equation, because changing a subscript changes the identity of a compound. Step [2] Balance the equation with coefficients one element at a time. Step [3] Check to make sure that the smallest set of whole numbers is used.

23. 23 aA + bB  cC + dD mass Amass Bmass Cmass D moles Amoles Bmoles Cmoles D a:b b:c c:d a:c a:d FW/MM Solve Stoichiometry Problems mol mol g g FW/MM Note: FW/MM means Formula wt. or Molar mass

24. Limiting Reactant  Compare the actual amount of each reactant to the amount required in the balanced equation to determine how many times the “reaction can be run”  Use the amount of the limiting reactant to calculate how much product can be produced

25. Smith. General Organic & Biolocial Chemistry 2nd Ed. 25 Zn + Cu 2+ Zn 2+ + Cu Zn loses 2 e – Cu 2+ gains 2 e − 25 Oxidation half reaction:ZnZn 2+ + 2 e − Each of these processes can be written as an individual half reaction: loss of e − Reduction half reaction:Cu 2+ + 2e − Cu gain of e − Redox Half Reactions