1. Summary Chapter 1 - 4 Chemistry: The Molecular Nature of Matter, 6 th edition By Jesperson, Brady, & Hyslop

2. CH 1-4 Concepts to be Familiar With  Classification of matter: pure substances & mixtures  Distinguish the difference between chemical and physical properties & changes  Atomic Theory  Laws of Definite Proportions & Conservation of Mass  Intensive vs extensive properties  Uncertainty in measurements & communicating that uncertainty with significant figures  Accuracy & Precision  Conversions between units (ie, dimensional analysis)  Structure of an atom: protons, neutrons, electrons  Navigate the periodic table: properties shared within a group, trends, metals/metalloids/nonmetals  Stoichiometric ratios within atoms and between different molecules  Difference between empirical, molecular, and structural formulas

3. 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 N A = 6.022 x 10 23 particles/mole Time conversions: d  hr  m  s 1 mL = 1 cm 3 T(kelvin) = T(°C) + 273.15 °F = 1.8°C +32 Equations Density = mass / Volume d = m/V d H2O = 1 g/mL = 1 g/cm 3 Molecular Mass (MM) = Molecular Weight (MW)= mass/moles MM = m/n MM molecule = Σ MM atoms in molecules Equations & Conversions to Memorize

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 Drawing Molecules: Methane CH 4 Emperical Formula: A x B y Ex: CH 3 O Molecular Formula: A n × x B n × y Ex: C 2 H 6 O 2 Structural Formula: A w A x B z Ex: HOCH 2 CH 2 OH

5. aA ( physical state ) + bB ( state )  cC ( state ) + dD ( state ) (physical state) = solid (s), liquid (l), gas (g), aqueous (aq) A, B = reactantsC, D = reactants a, b, c, d = coefficients to indicate molar ratios of reactants and products Chemical Equations 5 2C 4 H 10 + 13O 2  8CO 2 + 10H 2 O 2 molecules of C 4 H 10 13 molecules of O 2 10 molecules of C 4 H 10 8 molecules of CO 2 Balancing Chemical Equations: Unbalanced equation: C 4 H 10 + O 2  CO 2 + H 2 O Balanced equation:

6. Significant Figure Rules Scientific convention: All digits in measurement up to and including first estimated digit are significant. 1.All non-zero numbers are significant. 2.Zeros between non-zero numbers are significant. 3.Trailing zeros always count as significant if number has decimal point 4.Final zeros on number without decimal point are NOT significant 5.Final zeros to right of decimal point are significant 6.Leading zeros, to left of first nonzero digit, are never counted as significant Rounding Guidelines: 1.If digit to be dropped is greater than 5, last remaining digit is rounded up. 2.If number to be dropped is less than 5, last remaining digit stays the same. 3.If number to be dropped is exactly 5, then if digit to left of 5 is a. Even, it remains the same.b. Odd, it rounds up. Multiplication/Division & Addition/Subtraction: 1. Multiplication/Division: the number of significant figures in answer = number of significant figures in least precise measurement 2. Addition/Subtraction: the answer has same number of decimal places as quantity with fewest number of decimal places.

7. Molecules: Ionic Compounds Ions Transfer of one or more electrons from one atom to another Form electrically charged particles Ionic compound Compound composed of ions Formed from metal and nonmetal Properties Conducts electricity in liquid state when ions are free to move, but not as a solid Nomenclature: Cation (charge) Anion-ide ie: Iron (II) Oxide = FeO Fe 2+ O 2- Sodium Chloride = NaCl Na + Cl - “Criss-cross” rule  Make magnitude of charge on one ion into subscript for other  When doing this, make sure that subscripts are reduced to lowest whole number. Al 2 O 3 Al 3+ O 2–

8. Molecules: Ionic Compounds

9. Molecules: Covalent Compounds Molecules  Electrically neutral particle  Consists of two or more atoms  Covalent compounds have nonmetal-nonmetal bonds Chemical bonds  Attractions that hold atoms together in molecules  Arise from sharing electrons between two atoms  Group of atoms that make up molecule behave as single particle Molecular formulas  Describe composition of molecule  Specify number of each type of atom present Nonmetal hydrides  Molecule containing nonmetal + hydrogen  Number of hydrogens that combine with nonmetal = number of spaces from nonmetal to noble gas in periodic table Naming Binary Molecular Compounds –First element in formula Use English name –Second element  Use stem and append suffix –ide  Use Greek number prefixes to specify how many atoms of each element

10. Nomenclature Overview for Reference

11. Mole Ratios Within Molecules: A x B y Mole ratio of A:B = x:y Mole Ratios Between Molecules: aA + bB  cC + dD Mole ratio of A:B:C:D = a:b:c:d Mole Ratios

12. Percent Composition & Empirical Formulas Percent Composition (experimental or theoretical): Calculate percentage by mass of each element in sample (mass element) / (total mass of sample) x 100% Empirical Formula Simplest ratio of atoms of each element in compound Obtained from experimental analysis of compound If empirical formula is A x B y Molecular formula will be A n × x B n × y Molecular Formula Need molecular mass and empirical formula Calculate ratio of molecular mass to mass predicted by empirical formula and round to nearest integer

13. Stoichiometry aA + bB  cC + dD mass Amass Bmass Cmass D moles Amoles Bmoles Cmoles D a:b b:c c:d a:c a:d x ÷ MM Limiting Reactant: Least # moles after equalize with mole ratio Use limiting reactant to determine amount of product

14. Stoichiometry