Overview of the Basics

Transcription

1 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! 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 2

2 Chapter 1 Scientific Method 1. 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 2. Law or Scientific Law Usually an equation Based on results of many experiments Only states what happens Does not explain why they happen 3. Hypothesis Mental picture that explains observed laws Tentative explanation of data Make predictions Leads to further tests Go to laboratory and perform experiments 4. 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 3 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 diamond = carbon gold sulfur 4

3 Chapter 1 Elements 5 Chapter 1 Classification of Matter 6

4 Chapter 2 Chemical vs Physical Properties Physical properties Can be observed without changing chemical makeup of 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 Chemical properties Chemical change or reaction that substance undergoes Chemicals interact to form entirely different substances 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. 7 Chapter 1 Atomic Theory Developed by John Dalton to explain Law of Conservation of Mass & Law of Definite Proportions 1. Matter consists of tiny particles called atoms. 2. Atoms are indestructible. In chemical reactions, atoms rearrange but do not break apart. 3. In any sample of a pure element, all atoms are identical in mass and other properties. 4. Atoms of different elements differ in mass and other properties. 5. In a given compound, constituent atoms are always present in same fixed numerical ratio. 8

5 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 g of sulfur e.g. Water Mass ratio always: 8 g O to 1 g H 9 Chapter 1 Law of Conservation of Mass Mass S g Mass O g sulfur sulfur dioxide trioxide g g 10

6 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: Fe 2 O 3 : iron and oxygen in 2:3 ratio 11 Chapter 1 Hydrates Crystals that contain water molecules e.g. Plaster: CaSO 4 2H 2 O calcium sulfate dihydrate Water is not tightly held Dehydration Removal of water by heating Remaining solid is anhydrous (without water) Blue = White = CuSO 4 CuSO 4 5H 2 O 12

7 Chapter 1 Depicting Molecules CH 4 methane H H C H H 13 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. CH 4 + 2O 2!" CO 2 + 2H 2 O 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 14

8 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 CH 4 + 2O 2!" CO 2 + 2H 2 O 4H + 4O + C = 4H + 4O + C 15 Chapter 1 Balanced Chemical Equations Ex. 2C 4 H O 2!" 8CO H 2 O Ex. 4 C and 10 H per molecule 2 O per molecule 1 C and 2 O per molecule 2 H and 1 O per molecule 2C 4 H O 2!" 8CO H 2 O 2 molecules of C 4 H molecules of O 2 8 molecules of CO 2 10 molecules of C 4 H 10 16

9 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 17 Chapter 2 Measurements 1. 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 2. Measurements are inexact Measuring involves estimation Always have uncertainty The observer and instrument have inherent physical limitations 18

10 Chapter 2 International System of Units 19 Chapter 2 International System of Units 20

11 Chapter 2 International System of Units 21 Chapter 2 Decimal Multipliers 22

12 Chapter 2 4 Common Lab Measurements 1. Distance (d ) Centimeter (cm) 1 cm = 10 2 m = 0.01 m Millimeter (mm) 1 mm = 10 3 m = m 2. Volume (V) 1 L = 1000 ml 1 ml = 1 cm 3 3. Mass (m) 1 1 g = kg = g Temperature (T) K = 0 C 23 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 = ± 0.01 C 24

13 Chapter 2 Significant Figures Scientific convention: All digits in measurement up to and including first estimated digit are significant. 1. All non-zero numbers are significant. e.g has 4 sig. figs. 2. Zeros between non-zero numbers are significant. e.g. 20,089 or ! 10 4 has 5 sig. figs 3. Trailing zeros always count as significant if number has decimal point e.g or 5.00! 10 2 has 3 sig. figs 25 Chapter 2 Significant Figures 4. Final zeros on number without decimal point are NOT significant e.g. 104,956,000 or ! 10 8 has 6 sig. figs. 5. Final zeros to right of decimal point are significant e.g has 3 sig. figs. 6. Leading zeros, to left of first nonzero digit, are never counted as significant e.g or 1.2! 10 4 has 2 sig. figs. 26

14 Chapter 2 Significant Figures: Rounding 1. If digit to be dropped is greater than 5, last remaining digit is rounded up. e.g is rounded up to If number to be dropped is less than 5, last remaining digit stays the same. e.g is rounded to If number to be dropped is exactly 5, then if digit to left of 5 is a. Even, it remains the same. e.g is rounded to 6.6 b. Odd, it rounds up. e.g is rounded to Chapter 2 Significant Figures: Calculations Multiplication and Division Number of significant figures in answer = number of significant figures in least precise measurement e.g ! 31.4! 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 decimal places 1 decimal place 3 decimal places 1 decimal place 28

15 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 29 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! 10 4 If number estimated from aerial photograph Uncertainty ±100 people 7.50!

16 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 31 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 32

17 Chapter 2 Dimensional Analysis Example: Convert m to mm. Relationship is 1 mm = 1! 10 3 m Can make two conversion factors!3 1 mm 1 " 10 m!3 1 " 10 m 1 mm Since going from m to mm use one on left m 1 mm "!3 = 97 cm 1 " 10 m 33 Chapter 2 Density Ratio of object s mass to its volume density = mass volume d = m V 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/cm 3 34

18 Chapter 3 Discovery of Subatomic Particles in the late 1800 s and early 1900s Discovery of electron mass and charge Millikan Oil Drop expt Discovery of the Nucleus Rutherford Alpha scattering expt Discovery of the Electron 1897 Thomson Cathode ray tube expt Rutherford Nuclear Atom Discovery of Protons 1918 Rutherford Mass spectrometer Discovery of Neutron: 1932 Chadwick 35 Chapter 3 " Three kinds of subatomic particles of principal interest to chemists Properties of Subatomic Particles Nucleus (protons + neutrons) Electrons Particle Mass (g) Electrical Charge Electron # Proton # Neutron # Symbol 0!1e 1 1 H, 1 0 n 1 1p 36

19 Chapter 3 Atomic Notation A X Z Atomic number (Z) = Number of protons that atom has in nucleus Isotopes = Atoms of same element with different masses Same number of protons ( 1 ) Different number of neutrons ( 0 ) 1n 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 p Z = 2, A = 4 1 p 4 2 He 37 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) 38

20 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 12 C selected? Common Most abundant isotope of carbon All atomic masses of all other elements ~ whole numbers Lightest element, H, has mass ~1 u 39 Chapter 3 Calculating Atomic Mass Generally, elements are mixtures of isotopes e.g. Hydrogen Isotope Mass % Abundance 1 H u H u 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 40

21 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 41 Chapter 3 Periodic Table 42

22 Chapter 3 Periodic Table 43 Chapter 3 Periodic Table Groups 1A 2A B B 7A 8A Alkali Metals Alkaline Earth Metals Transition Metals Lanthanide & Actinide Halogens Nobel Gases Very reactive Metals except for H +1 ions React with Oxygen to form compounds that dissolve into alkaline solutions in water Reactive +2 ions Oxygen compounds are strongly alkaline Many are not water soluble Metals Form ions with several different charges (oxidation states) Tend to form +2 and +3 ions Lanthanides Actinides Actinides are radioactive Reactive Form diatomic molecules in elemental state -1 ions Salts with alkali metals Inert Heavier elements have limited reactivity Do not form ions Monoatomic gases 44

23 Chapter 3 Metals, Nonmetals, and Metalloids 45 Chapter 3 Metals, Nonmetals, and Metalloids Metals Nonmetals Metalloids 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 Insulators, nonconductors 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. Metallic shine but brittle Semiconductors: conduct electricity but not as well as metals: examples are silicon and germanium 46

24 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 47 Chapter 3 Ions and Ionic Compounds Metal + Non-metal!" ionic compound 2Na(s) + Cl 2 (g)!" 2NaCl(s) Michael Watson Richard Megna/Fundamental Photographs Na + Cl Na + Richard + Cl! Megna/Fundamental Photographs NaCl(s) e! Anions = Negatively charged ions Cations = Positively charged ions 48

25 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 49 Chapter 3 Ions and Ionic Compounds Negative ( ) charge on anion = number of spaces you have to move to right to get to noble gas 50

26 Chapter 3 Rules for Writing Ionic Formulas 1. Cation given first in formula 2. Subscripts in formula must produce electrically neutral formula unit 3. Subscripts must be smallest whole numbers possible Divide by 2 if all subscripts are even May have to repeat several times 4. Charges on ions not included in finished formula unit of substance If no subscript, then 1 implied 51 Chapter 3 Determining Ionic Formulas 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. Ex. What is the formula of ionic compound formed between aluminum and oxygen ions? Al 3+ O 2 Al 2 O 3 52

27 Chapter 3 Transition Metal and Post-Transition Metal Ions 53 Chapter 3 Polyatomic Ions Example: What is the formula of the ionic compound formed between ammonium and phosphate ions? Ammonium = NH 4 + Phosphate = PO 4 3 (NH 4 ) + (PO 4 ) 3 (NH 4 ) 3 PO 4 54

28 Chapter 3 Nomenclature 55 Chapter 3 Nomenclature: Ionic Compounds Cations: Metal that forms only one positive ion Cation name = English name for metal Na + sodium Ca 2+ calcium Metal that forms more than one positive ion Use Stock System Cation name = English name followed by numerical value of charge written as Roman numeral in parentheses (no spaces) Transition metal Cr 2+ chromium(ii) Cr 3+ chromium(iii) 56

29 Chapter 3 Nomenclature: Ionic Compounds Anions: Monatomic anions named by adding ide suffix to stem name for element Polyatomic ions use names in Table Chapter 3 Nomenclature: Hydrates Ionic compounds Crystals contain water molecules Fixed proportions relative to ionic substance Naming Name ionic compound Give number of water molecules in formula using Greek prefixes mono- = 1 hexa- = 6 di- = 2 hepta- = 7 tri- = 3 octa- = 8 tetra- = 4 nona- = 9 penta- = 5 deca- = 10 58

30 Chapter 3 Molecular Compounds Molecules Electrically neutral particle Consists of two or more atoms 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 59 Chapter 3 Nonmetal Hydrides Nonmetal hydrides Molecule containing nonmetal + hydrogen Number of hydrogens that combine with nonmetal = number of spaces from nonmetal to noble gas in periodic table N O F Ne 60

31 Chapter 3 Organic Compound Formulas Molecular formula Indicates number of each type of atom in molecule e.g. C 2 H 6 for ethane or C 3 H 8 for propane Order of atoms Carbon Hydrogen Other atoms alphabetically e.g. sucrose is C 12 H 22 O 11 Emphasize alcohol write OH group last C 2 H 5 OH Structural formula Indicate how carbon atoms are connected Ethane = CH 3 CH 3 Propane = CH 3 CH 2 CH 3 61 Chapter 3 Nomenclature: Molecular Compounds Goal is a name that translates clearly into molecular formula Naming Binary Molecular Compounds Which two elements present? How many of each? Format: 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 62

32 Chapter 3 Nomenclature: Binary Molecules 1. hydrogen chloride 1 H 1 Cl HCl 2. phosphorous pentachloride 3. triselenium dinitride 1 P5Cl PCl 5 3 Se 2N Se 3 N 2 Mono always omitted on first element Often omitted on second element unless more than one combination of same two elements e.g. Carbon monoxide CO Carbon dioxide CO 2 When prefix ends in vowel similar to start of element name, drop prefix vowel 63 Chapter 3 Nomenclature: Exceptions for Binary Molecules Binary compounds of nonmetals + hydrogen No prefixes to be used Get number of hydrogens for each nonmetal from periodic table Hydrogen sulfide = H 2 S Hydrogen telluride = H 2 Te Molecules with Common Names Some molecules have names that predate IUPAC systematic names Water H 2 O " Sucrose C 12 H 22 O 11 Ammonia NH 3 " Phosphine PH 3 64