CHAPTER ONE The Foundations of Chemistry

RECALL!!! Elements: Can not be broken down by chemical reactions Pure Substances Compounds: Can be broken down ONLY by chemical reactions Matter Homogeneous: Constant composition throughout Mixtures Can be separated by Heterogeneous: Variable composition physical processesthroughout 2

3 States of Matter

4 A B C D E

5

6 Chemical and Physical Properties Chemical Properties - chemical changes rusting or oxidation chemical reactions Physical Properties - physical changes changes of state density, color, solubility

7 Question 1. Classify as mixtures (homogeneous or heterogeneous), compounds (ionic or molecular), or elements (monoatomic or molecular). Classification C 2 H 5 OH Cl 2 Cu bronze 5% AgNO 3 solution C 6 H 12 O 6 PbCl 2 precipitate Zn Ba 3 (PO 4 ) M CaCl 2

8 Mixtures, Substances, Compounds, and Elements

Separation of Mixtures  Distillation: It separates 2 or more liquids with different boiling points. For example: ethanol (bp: 78°C) and water (bp: 100°C)  Fractional crystallization: It separates 2 or more solids (by means of precipitation) with different solubility.  Filtration: It separates an insoluble solid from a liquid. The solid must be insoluble in the liquid. For example: sand- water, silver chloride-water.  Chromatography: It separates substances that are soluble in a solvent by means of IMF.  Evaporation: It separates a soluble solute from its solvent by evaporating the solvent. For example: NaCl-H 2 O  Liquid – liquid separation: It separates 2 immiscible liquids using a separatory funnel. For example: oil-water

10 Distillation Separates homogeneous mixture on the basis of differences in boiling point. Ethanol-water

11 Filtration Separates insoluble solid substances from liquids and solutions.

12 Chromatography: Separates substances on the basis of differences in solubility in a solvent and IMF.

13 Measurements in Chemistry QuantityUnitSymbol  lengthmeter m  masskilogram kg  timesecond s  currentampere A  temperatureKelvin K  amt. substancemole mol

14 Measurements in Chemistry Metric Prefixes NameSymbolMultiplier mega M 10 6 kilo k 10 3 deci d centi c milli m micro μ nano n pico p femto f

15 Units of Measurement Common Conversion Factors Length 2.54 cm = 1 inch Volume 1 liter = 1.06 qt Mass 1 lb = 454 g

16 Volume The most commonly used metric units for volume are the liter (L) and the milliliter (mL). 1 dm = 10 cm 1 dm 3 = 1000 cm 3 1 L = 1000 mL therefore 1 mL = 1 cm 3

17 Uncertainty in Measurements Different measuring devices have different uses and different degrees of accuracy.

18 Relationships of the Temperature Scales Kelvin and Celsius Relationship K = °C Fahrenheit and Celsius Relationship °F = 1.8 * °C + 32

19 Use of Numbers Exact numbers 1 dozen = 12 things for example Accuracy how closely measured values agree with the correct value. The experimental value is g, the actual value is g. These 2 masses are accurate. Precision how closely individual measurements agree with each other. The value of the mass of the same beaker in 3 trials are: g, g, g. These values are precise.

20 Percent error Percent error = accepted – experimental x 100 accepted From previous example: % error = – x 100 = %

21 Use of Numbers Significant Figures – Rules All non-zero digits are significant. Leading zeroes are never significant has three significant figures Imbedded zeroes are always significant has five sig fig Trailing zeroes are only significant after the decimal point g has 2 sig fig g has 4 sig fig Use scientific notation to remove doubt x 10 3 has 4 significant figures

22 Use of Numbers Multiplication & Division rule Easier of the two rules Product has the smallest number of significant figures of multipliers

23 Use of Numbers Addition & Subtraction rule More subtle than the multiplication rule Answer contains smallest decimal place of the addends.

24 Use of Numbers On a multi-step question, solve all mathematical steps and ONLY round off the final answer.

25 Using Factor Label Method Example: A concentrated hydrochloric acid solution is 36.31% HCl by mass. The density of the solution is 1.185g/mL. What mass of pure HCl is contained in 175 mL of this solution?

26 Problem 2. Calculate the volume of solution required to prepare M solution of Na 2 SO 4 if only 0.050g of the salt is available. Known: Molar mass of Na 2 SO 4 is g/mol M means mol/L.

Problem 3. Calculate the density of a cobalt(II) chloride solution with a molarity of 3.57 M and a percent mass by mass of 17.46%. 27

Review Nomenclature Compounds: Ionic Covalent Hydrates Acids

Ionic compounds Metal + Nonmetal Groups IA, IIA, Al, Ag, Zn, Cd Only one oxidation number Name of the metal does not change: lithium, calcium Other metals More than one oxidation number Roman numerals to specify charge: iron(II), iron(III), tin(II), tin(IV) Stem of element + ide oxide, sulfide, chloride Binary Two elements Ternary More than two elements

Ionic compounds Metal or NH Polyatomic anion Groups IA, IIA, Al, Ag, Zn, Cd Only one oxidation number Name of the metal does not change: lithium, calcium Other metals More than one oxidation number Roman numerals to specify charge: iron(II), iron(III), tin(II), tin(IV) Binary Two elements Ternary More than two elements

Binary Molecular Compounds non metal + stem of second non metal ending ide

32 Formula-to-Name Acids Acids are molecular compounds that often behave like they are made of ions. All names have acid at end. Binary Acids = Hydro- prefix + stem of the name of the nonmetal + -ic suffix. Oxyacids: If polyatomic ion ends in –ate = Name of polyatomic ion with –ic suffix. If polyatomic ion ends in –ite = Name of polyatomic ion with –ous suffix.

Hydrates Ionic compounds that crystallize with water occluded in their crystal structure. NaCO 3 ·10H 2 O BaCl 2 · 2H 2 O Co(NO 3 ) 2 · 6H 2 0