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Overview of the Basics CHAPTER 1-3 Review Chemistry: The Molecular Nature of Matter, 6 th edition By Jesperson , Brady

Overview of the Basics CHAPTER 1-3 Review Chemistry: The Molecular Nature of Matter, 6 th edition By Jesperson , Brady, & Hyslop. CHAPTER 1-3 Review. Learning Objectives Scientific Method Matter: definition, elements, compounds, mixtures, changes/properties Atomic Theory

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Overview of the Basics CHAPTER 1-3 Review Chemistry: The Molecular Nature of Matter, 6 th edition By Jesperson , Brady

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  1. Overview of the Basics CHAPTER 1-3 Review Chemistry: The Molecular Nature of Matter, 6th edition By Jesperson, Brady, & Hyslop

  2. CHAPTER 1-3 Review • Learning Objectives • Scientific Method • Matter: definition, elements, compounds, mixtures, changes/properties • Atomic Theory • Law of definite proportions • Law of conservation of mass • Chemical formulas • Chemical equations • Balancing • Measurements: units, conversions, uncertainty • Significant Figures • Density • Subatomic particles • Atomic #, mass #, atomic weights • Periodic Table • Ionic Compounds, hydrates, molecular compounds • Basic nomenclature Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  3. Chapter 1 Scientific Method • Make observations/collect data • Empirical fact • Something we see, hear, taste, feel, or smell • Something we can measure • Organize data so we can see relationships • Law or Scientific Law • Usually an equation • Based on results of many experiments • Only states what happens • Does not explain why they happen • Hypothesis • Mental picture that explains observed laws • Tentative explanation of data • Make predictions • Leads to further tests • Go to laboratory and perform experiments • Theory • Tested explanation of how nature behaves • Devise further tests • Depending on results, may have to modify theory • Can never prove theory is absolutely correct Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  4. Chapter 1 Elements • Substances that can’t be decomposed into simpler materials by chemical reactions • Substances composed of only one type of atom • Simplest forms of matter that we can work with directly • More complex substances composed of elements in various combinations sulfur diamond = carbon gold Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  5. Chapter 1 Elements Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  6. Chapter 1 Classification of Matter Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E http://ridenourmhs.wikispaces.com/ESUnit2

  7. Chapter 2 Chemical vs Physical Properties • Physical properties • Can be observed without changing chemical makeup of substance • Chemical properties • Chemical change or reaction that substance undergoes • Chemicals interact to form entirely differentsubstances with different chemical and physical properties • Describe behavior of matter that leads to formation of new substance • “Reactivity" of substance e.g. Iron rusting • Iron interacts with oxygen to form a new substance. Solids: • Fixed shape and volume • Particles are close together Liquids: • Fixed volume, but take container shape • Particles are close together Gases: • Expand to fill entire container • Particles separated by lots of space Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  8. Chapter 1 Atomic Theory Developed by John Dalton to explain Law of Conservation of Mass & Law of Definite Proportions • Matter consists of tiny particles called atoms. • Atoms are indestructible. • In chemical reactions, atoms rearrange but do not break apart. • In any sample of a pure element, all atoms are identical in mass and other properties. • Atoms of different elements differ in mass and other properties. • In a given compound, constituent atoms are always present in same fixed numericalratio. Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  9. Chapter 1 Law of Definite Proportions • Atoms react as Whole particles. • When two elements form more than one compound, different masses of one element that combine with same mass of other element are always in ratio of small whole numbers. • e.g.Fool’s gold, pyrite, iron(III) sulfide • Mass ratio always • 1.00 g of iron to 0.574 g of sulfur • e.g. Water • Mass ratio always: 8 g O to 1 g H Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  10. Chapter 1 Law of Conservation of Mass sulfur sulfur dioxide trioxide Mass S 32.06 g 32.06 g Mass O 32.00 g 48.00 g Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  11. Chapter 1 Molecules and Chemical Formulas Atoms combine to form more complex substances = Molecules Chemical Formulas: • Specify composition of substance • Chemical symbols represent atoms of elements present • Subscripts: • Given after chemical symbol • Represents relative numbers of each type of atom Example: Fe2O3 : iron and oxygen in 2:3 ratio Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  12. Chapter 1 Hydrates • Crystals that contain water molecules e.g. Plaster: CaSO4∙2H2O calcium sulfate dihydrate • Water is not tightly held • Dehydration • Removal of water by heating • Remaining solid is anhydrous (without water) White = CuSO4 Blue = CuSO4 •5H2O Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  13. Chapter 1 Depicting Molecules CH4 methane Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  14. Chapter 1 Chemical Equations • Use chemical symbols and formulas to represent reactants and products. • Reactants on left hand side • Products on right hand side • Arrow () means “reacts to yield” e.g.CH4 + 2O2 CO2 + 2H2O • Coefficients • Numbers in front of formulas • Indicate how many of each type of molecule reacted or formed • Equation reads “methane and oxygen react to yield carbon dioxide and water” Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  15. Chapter 1 Conservation of Mass in Reactions • Mass can neither be created nor destroyed • This means that there are the same number of each type of atom in reactants and in products of reaction • If number of atoms same, then mass also same CH4 + 2O2CO2 + 2H2O 4H + 4O + C = 4H + 4O + C Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  16. Chapter 1 Balanced Chemical Equations Ex.2C4H10 + 13O2 8CO2 + 10H2O 10 molecules of C4H10 2 molecules of C4H10 4 C and 10 H per molecule 13 molecules of O2 2 H and 1 O per molecule 8 molecules of CO2 2 O per molecule 1 C and 2 O per molecule Ex. 2C4H10 + 13O2 8CO2 + 10H2O Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  17. Chapter 2 Intensive vs Extensive Properties Intensive properties • Independent of sample size • Used to identify substances e.g.Color Density Boiling point Melting point Chemical reactivity Extensive properties • Depend on sample size e.g.volume and mass Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  18. Chapter 2 Measurements • Measurements involve comparison • Always measure relative to reference e.g. Foot, meter, kilogram • Measurement = number + unit e.g. Distance between 2 points = 25 • What unit? inches, feet, yards, miles • Meaningless without units • Measurements are inexact • Measuring involves estimation • Always have uncertainty • The observer and instrument have inherent physical limitations Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  19. Chapter 2 International System of Units Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  20. Chapter 2 International System of Units Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  21. Chapter 2 International System of Units Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  22. Chapter 2 Decimal Multipliers Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  23. Chapter 2 4 Common Lab Measurements • Distance (d ) • Centimeter (cm) • 1 cm = 10–2 m = 0.01 m • Millimeter (mm) • 1 mm = 10–3 m = 0.001 m • Volume (V) • 1 L = 1000 mL • 1 mL = 1 cm3 • Mass (m) • 1 g = 0.1000 kg = g • Temperature (T) • 273.15 K = 0°C Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  24. Chapter 2 Uncertainty in Measurements • Measurements all inexact • Limitations of reading instrument Example: Consider two Celsius thermometers • Left thermometer has markings every 1˚C • T between 24 °C and 25 °C • About 3/10 of way between marks • Can estimate to 0.1 °C = uncertainty • T = 24.3  0.1 °C • Right thermometer has markings every 0.1 °C • T reading between 24.3 °C and 24.4 °C • Can estimate 0.01 °C • T = 24.32  0.01 °C Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  25. Chapter 2 Significant Figures Scientific convention: All digits in measurement up to and including first estimated digit are significant. • All non-zero numbers are significant. e.g.3.456 has 4 sig. figs. • Zeros between non-zero numbers are significant. e.g. 20,089 or 2.0089 × 104 has 5 sig. figs • Trailing zeros always count as significant if number has decimal point e.g.500. or 5.00 × 102 has 3 sig. figs Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  26. Chapter 2 Significant Figures • Final zeros on number without decimal point are NOT significant e.g.104,956,000 or 1.04956 × 108 has 6 sig. figs. • Final zeros to right of decimal point are significant e.g.3.00 has 3 sig. figs. • Leading zeros, to left of first nonzero digit, are never counted as significant e.g.0.00012 or 1.2 × 10–4 has 2 sig. figs. Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  27. Chapter 2 Significant Figures: Rounding • If digit to be dropped is greater than 5, last remaining digit is rounded up. e.g.3.677 is rounded up to 3.68 • If number to be dropped is less than 5, last remaining digit stays the same. e.g.6.632 is rounded to 6.63 • If number to be dropped is exactly 5, then if digit to left of 5 is • Even, it remains the same. e.g.6.65 is rounded to 6.6 • Odd, it rounds up. e.g.6.35 is rounded to 6.4 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  28. Chapter 2 Significant Figures: Calculations Multiplication and Division • Number of significant figures in answer = number of significant figures in leastprecise measurement e.g.10.54 × 31.4 × 16.987 4 sig. figs. × 3 sig. figs. × 5 sig. figs. = 3 sig. figs. Addition and Subtraction • Answer has same number of decimal places as quantity with fewest number of decimal places. 4 decimal places 1 decimal place 3 decimal places 1 decimal place 12.9753 319.5 + 4.398 336.9 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  29. Chapter 2 Significant Figures: Exact Numbers • Numbers that come from definitions • 12 in. = 1 ft • 60 s = 1 min • Numbers that come from direct count • Number of people in small room • Have no uncertainty • Assume they have infinite number of significant figures. • Do not affect number of significant figures in multiplication or division Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  30. Chapter 2 Scientific Notation • Clearest way to present number of significant figures unambiguously • Report number between 1 and 10 followed by correct power of 10 • Indicates only significant digits e.g.75,000 people attend a concert • If a rough estimate • Uncertainty 1000 people • 7.5 × 104 • If number estimated from aerial photograph • Uncertainty 100 people • 7.50 × 104 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  31. Chapter 2 Accuracy & Precision Accuracy • How close measurement is to true or accepted true value • Measuring device must be calibrated with standard reference to give correct value Precision • How well set of repeated measurements of same quantity agree with each other • More significant figures equals more precise measurement Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  32. Chapter 2 Dimensional Analysis • Also called the Factor Label Method • Not all calculations use specific equation • Use units (dimensions) to analyze problem Conversion Factor • Fraction formed from valid equality or equivalence between units • Used to switch from one system of measurement and units to another Given Quantity Conversion Factor Desired Quantity × = Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  33. Chapter 2 Dimensional Analysis Example:Convert 0.097 m to mm. • Relationship is 1 mm = 1 × 10–3 m • Can make two conversion factors • Since going from m to mm use one on left. = 97 cm Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  34. Chapter 2 Density • Ratio of object’s mass to its volume • Intensive property (size independent) • Determined by taking ratio of two extensive properties (size dependent) • Frequently ratio of two size dependent properties leads to size independent property • Density useful to transfer between mass and volume of substance • Density decreases slightly as temperature increases • Units: g/mL or g/cm3 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  35. Discovery of Subatomic Particles in the late 1800’s and early 1900s Chapter 3 Discovery of the Nucleus Rutherford Alpha scattering expt Discovery of electron mass and charge Millikan Oil Drop expt Discovery of Protons 1918 Rutherford Mass spectrometer Discovery of the Electron 1897 Thomson Cathode ray tube expt Discovery of Neutron: 1932 Chadwick Rutherford Nuclear Atom Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  36. Chapter 3 Properties of Subatomic Particles • Three kinds of subatomic particles of principal interest to chemists Nucleus (protons + neutrons) Electrons Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  37. Chapter 3 Atomic Notation Atomic number (Z) = Number of protons that atom has in nucleus Isotopes = Atoms of same element with different masses • Same number of protons ( ) • Different number of neutrons ( ) Isotope Mass number (A) • A = (number of protons)+(number of neutrons) = Z + N • For charge neutrality, number of electrons and protons must be equal Atomic Symbols = Summarize information about subatomic particles • Every isotope defined by two numbers Z and A Ex.What is the atomic symbol for helium? He has 2 e–, 2 n and 2 pZ = 2, A = 4 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  38. Chapter 3 Isotopes • Most elements are mixtures of two or more stable isotopes • Each isotope has slightly different mass • Chemically, isotopes have virtually identical chemical properties • Relative proportions of different isotopes are essentially constant • Isotopes distinguished by mass number (A): e.g. • Three isotopes of hydrogen (H) • Four isotopes of iron (Fe) Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  39. Chapter 3 Carbon-12 Atomic Mass Scale • Need uniform mass scale for atoms Atomic mass units(symbol u) • Based on carbon: • 1 atom of carbon-12 = 12 u (exactly) • 1 u = 1/12 mass 1 atom of carbon-12 (exactly) Why was 12C selected? • Common • Most abundant isotope of carbon • All atomic masses of all other elements ~ whole numbers • Lightest element, H, has mass ~1 u Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  40. Chapter 3 Calculating Atomic Mass • Generally, elements are mixtures of isotopes e.g. Hydrogen Isotope Mass % Abundance 1H 1.007825 u 99.985 2H 2.0140 u 0.015 How do we define atomic mass? • Average of masses of all stable isotopes of given element How do we calculate average atomic mass? • Weighted average • Use isotopic abundances and isotopic masses Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  41. Chapter 3 Periodic Table • Summarizes periodic properties of elements Early Versions of Periodic Tables • Arranged by increasing atomic mass • Mendeleev (Russian) and Meyer (German) in 1869 • Noted repeating (periodic) properties Modern Periodic Table • Arranged by increasing atomic number (Z ): • Rows called periods • Columns called groups or families • Identified by numbers • 1 – 18 standard international • 1A – 8A longer columns and 1B – 8B shorter columns Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  42. Chapter 3 Periodic Table Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  43. Chapter 3 Periodic Table Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  44. Chapter 3 Periodic Table Groups Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  45. Chapter 3 Metals, Nonmetals, and Metalloids Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  46. Chapter 3 Metals, Nonmetals, and Metalloids Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  47. Chapter 3 Ions and 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 • Infinite array of alternating Na+ and Cl– ions Formula unit • Smallest neutral unit of ionic compound • Smallest whole-number ratio of ions Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  48. Chapter 3 Ions and Ionic Compounds Metal + Non-metal  ionic compound 2Na(s) + Cl2(g)  2NaCl(s) Michael Watson Richard Megna/Fundamental Photographs Richard Megna/Fundamental Photographs Anions = Negatively charged ions Cations = Positively charged ions Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  49. Chapter 3 Ions and Ionic Compounds Electrical conductivity requires charge movement Ionic compounds: • Do not conduct electricity in solid state • Do conduct electricity in liquid and aqueous states where ions are free to move Molecular compounds: • Do not conduct electricity in any state • Molecules are comprised of uncharged particles Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  50. Chapter 3 Ions and Ionic Compounds Negative(–) charge on anion = number of spaces you have to move to right to get to noble gas Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

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