Summary Chapter 1 - 4 Chemistry: The Molecular Nature of Matter, 6 th edition

1. Summary Chapter 1 - 4 Chemistry: The Molecular Nature of Matter, 6th 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 vsextensive 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. Equations & Conversions to Memorize Unit Conversions For metric units (m, kg, s, K, mole): mega (M) 106 kilo (k) 103 centi (c) 10-2 milli (m) 10-3 micro (μ) 10-6 nano (n) 10-9 Pico (p) 10-12 NA = 6.022 x 1023 particles/mole Time conversions: dhrms 1 mL = 1 cm3 T(kelvin) = T(°C) + 273.15 °F = 1.8°C +32 Equations Density = mass / Volume d = m/V dH2O = 1 g/mL = 1 g/cm3 Molecular Mass (MM) = Molecular Weight (MW)= mass/moles MM = m/n MMmolecule = ΣMMatoms in molecules

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 Elements on the Periodic Table element symbol atomic weight (amu) = weighted average of atomic weight of isotopes • Emperical Formula: AxByEx: CH3O • Molecular Formula: An × xBn× y Ex: C2H6O2 • Structural Formula: AwAxBzEx: HOCH2CH2OH Drawing Molecules: Methane CH4

5. Chemical Equations aA(physical state) + bB(state)  cC(state) + dD(state) (physical state) = solid (s), liquid (l), gas (g), aqueous (aq) A, B = reactants C, D = reactants a, b, c, d = coefficients to indicate molar ratios of reactants and products 10 molecules of C4H10 2 molecules of C4H10 13 molecules of O2 8 molecules of CO2 Balancing Chemical Equations: Unbalanced equation: C4H10+ O2 CO2+ H2O Balanced equation: 2C4H10 + 13O2 8CO2 + 10H2O

6. Significant Figure Rules Scientific convention: All digits in measurement up to and including first estimated digit are significant. • All non-zero numbers are significant. • Zeros between non-zero numbers are significant. • Trailing zeros always count as significant if number has decimal point • Final zeros on number without decimal point are NOT significant • Final zeros to right of decimal point are significant • Leading zeros, to left of first nonzero digit, are never counted as significant Rounding Guidelines: • If digit to be dropped is greater than 5, last remaining digit is rounded up. • If number to be dropped is less than 5, last remaining digit stays the same. • If number to be dropped is exactly 5, then if digit to left of 5 is • Even, it remains the same. b. Odd, it rounds up. Multiplication/Division & Addition/Subtraction: • Multiplication/Division: the number of significant figures in answer = number of significant figures in least precise measurement • 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 Fe2+ O2- 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. Al3+ O2– Al2O3

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 Mole Ratios Within Molecules: AxBy 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

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 AxBy • Molecular formula will be An × xBn× 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 A mass B mass C mass D x ÷ MM x ÷ MM x ÷ MM x ÷ MM moles A moles B moles C moles D a:b c:d b:c a:c a:d • Limiting Reactant: • Least # moles after equalize with mole ratio • Use limiting reactant to determine amount of product

14. Stoichiometry