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AP Chemistry

AP Chemistry . Chapter 1: Matter and Measurement. The States of Matter. Solid: Particles are highly ordered and can not flow Liquid: Particles are loosely ordered and are able to flow. Gas: Atoms are highly disordered and separate from one another. Plasma: Ionized gas.

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AP Chemistry

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  1. AP Chemistry Chapter 1: Matter and Measurement

  2. The States of Matter • Solid: Particles are highly ordered and can not flow • Liquid: Particles are loosely ordered and are able to flow. • Gas: Atoms are highly disordered and separate from one another. • Plasma: Ionized gas. • Bose-Einstein Condensate (BEC): More about this later

  3. Bose-Einstein Condensate

  4. Physical vs. Chemical Properties • Physical Properties: Can be measured with out changing the composition of the substance. • Chemical Properties: Describe the way a substance may change, or react, to form other substances. • Properties of Matter: • Intensive properties: Do not depend on the amount of the substance being studied. • Extensive properties: Do depend on the amount of the substance being studied.

  5. Physical vs. Chemical Changes • Physical Change: A change that occurs without altering the chemical composition of a substance. • Chemical Change (or reaction): A change that transforms a substance into a chemically different substance.

  6. Making Measurements • SI units: In 1960 an international agreement was made stating which metric units would be used to make scientific measurements. • Mass: • Kilogram (kg) • Length: • Meter (m) • Temperature: • Kelvin (K) • Amount of Substance: • Mole (mol) • Electric current: • Ampere (A) • Luminous Intensity: • Candela (cd)

  7. SI Prefixes

  8. Temperature • Common measurements of temperature: • Fahrenheit • Celsius • Kelvin • Converting between Fahrenheit and Celsius: • 0F = 0C (9/5) + 32 • Converting between Celsius and Kelvin: • K = 0C + 237.15

  9. Derived SI units • Some things can not be directly measured. • Examples: • Speed is measured as meters per second. • Volume is measured as cubic length. • Density is measured as grams per unit volume.

  10. Uncertainty in Measurment • There are two kinds of numbers in scientific work, exact and inexact. • Exact numbers have a definite value. • Examples: • There are exactly 12 eggs in a dozen • There are exactly 1000 g in a kg • There are exactly 2.54 cm in an inch • Numbers obtained through measurements are always inexact. • This is due to human error as well as error in the measurement tools we use.

  11. Accuracy Vs. Precision • Accuracy is how close a measurement is to the correct or “true” value. • Precision is how close individual measurements are to one another. • To have good precision a measurement device needs to be able to make reproducible measurements. This is why we calibrate instruments. • The precisions of the measurements we make is often expressed in terms of STANDARD DEVIATION.

  12. Significant Figures • All measuring instruments have a certain degree of precision. • Instruments with more subdivisions have greater precision. • We report measurements we take using the smallest subdivision and one guess. • All of the numbers we know for certain and that one guess are called significant figures.

  13. Sig. Fig. Rules • 1. Zeros between nonzero digits are always significant. • 1005 kg – Has 4 significant figures • 2. Zeros at the beginning of a number are never significant. • 0.02 g – Has one sig. fig. • 0.0025 - Has two sig. figs. • 3. Zeros at the end of a number are significant only if there is a decimal in the number. • 0.0200 g – Has three sig. figs. • 3.0 cm – Has two sig. figs. • 100 cm – Has only one sig. fig.

  14. Sig. Figs. in Calculations • When we use measured quantities to do calculations, the least certain measurement limits the certainty of our calculation. • Therefore the number of significant figures in our answer is determined by the number of sig figs in the least certain number. • Rules: • For addition and subtraction: The answer has the same number of decimal places as the number with the least amount of decimal places. • 20.42 + 1.322 + 83.1 = 104.842, we round to 104.8 • For multiplication and division: The answer has the same number of sig figs as the number with the smallest number of sig figs. • 6.221 x 5.2 = 32.3492, we round to 32

  15. Conversions • The best way to do unit conversions is through dimensional analysis. • Once we know the relationship between two units we can use it to convert. • Example: 1 in. = 2.54 cm • Convert 8.50 inches into centimeters.

  16. Conversions with two or more steps • Convert 8.00 m into inches.

  17. Chapter 2 Atoms, Molecules and Ions

  18. Atomic Theory • The modern atomic theory is credited to and English school teacher by the name of John Dalton. • His theory states that: • Each element is composed of extremely small particles called atoms. • All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements. • The atoms of one element cannot be changed into atoms of another element. • compounds are formed when atoms of more than one element combine.

  19. Subatomic Particles • Expanding on Dalton’s theory we now know that atoms are made of even smaller particles. • Protons • Electrons • Neutrons

  20. Electrons • The discovery of the electron can be credited to one simple idea: • Particles with the same charge repel one another, whereas particles with unlike charges attract one another. • British scientist J.J. Thomson is credited with the “discovery’ of the electron.

  21. The Cathode Ray Tube

  22. The Mass of an Electron • After making his observations using cathode ray tubes Thomson was able to calculate the value of 1.76 x 108 Coulombs per gram. • He determined that this was the ratio of the electrons charge to its mass. • Using the charge of one electron (1.602 x 10-19 C) Thompson then calculated the mass of an electron.

  23. Radioactivity • Radioactivity is the spontaneous emission of radiation from an atom. • There are three types of radiation • α • β • γ • Each type of radiation reacts differently to an electric field.

  24. The Nucleus • Thomson reasoned that since electrons contribute only a very small fraction of the mass of the atom then they were probably only responsible for a small fraction of the atoms size….he was wrong

  25. The Rutherford Experiment

  26. The modern view of the atom • Since the time of Rutherford scientists have learned much about the structure of the atom. • Atoms are composed of protons, electrons and neutrons. • We have already seen that the atoms of one element are all the same, and are different than atoms of other elements. • Atoms of different elements have a characteristic number of protons.

  27. Atomic number: • Mass number: • Isotopes:

  28. Average Atomic Masses • Most elements occur in nature as a mixture of isotopes. • We can determine the average atomic mass of an element by using the masses of its various isotopes and their relative abundances.

  29. The Periodic Table

  30. Periodic Table

  31. Molecules and Molecular Compounds • Even though the atom is the smallest possible sample of an element only the noble gasses are normally found in nature as isolated atoms. • A molecule is an assembly of two or more atoms. • Molecular Formulas: • Empirical Formulas:

  32. Ions and Ionic Compounds • The nucleus of an atom can not be changed by chemical processes, but some atoms can gain or lose electrons. • If electrons are removed or added to an atom we are left with a charged particle called an ion. • Cation • Anion • When two or more ions combine they form an ionic compound.

  33. Polyatomic Ions

  34. Naming Ions • Cations (+): Cations formed from metal atoms have the same name as the metal • Examples: • Na+ - Sodium Ion • Zn 2+ - Zinc Ion • Al3+ - Aluminum Ion • If a metal can form different ions the charge is indicated by a roman numeral • Examples: • Fe2+ - Iron (II) ion • Fe3+ - Iron (III) ion

  35. Cations formed from non-metal atoms have names that end in –ium. • Examples: • NH4+ - Ammonium ion • H3O+ - Hydronium ion • Anions that are formed from a single atom end in –ide • H- - Hydride • O2- - Oxide • N3- - Nitride • A few simple polyatomic anions also end in –ide • OH- - Hydroxide • CN- - Cyanide • O22- - Peroxide

  36. Polyatomic anions containing oxygen end in -ate or –ite • NO3- - Nitrate • NO2- - Nitrite • SO42- - Sulfate • SO32- - Sulfite

  37. Naming Ionic Compounds • Names of ionic compounds consist of the cation name followed the anion name. • CaCl2 – Calcium Chloride • Al(NO3)3 – Aluminum Nitrate • Cu(ClO4)2 – Copper (II) Perchlorate

  38. Names and Formulas of Acids • Acids containing anions whose names end in –ide are named by changing the –ide to –ic, and adding the prefix hydro. • HCl – Hydrochloric acid • H2S – hydrosulfuric acid • Acids containing anions whose names end in – ate or – ite are named by changing –ate into –ic, and - ite into –ous. • HClO3 – Chloric acid • HClO2 – Chlorous acid

  39. Names and Formulas of Binary Molecular Compounds • The name of the element farther to the left in the periodic table is usually written first. • If both elements are in the same group in the periodic table, the one having the higher atomic number is named first. • The name of the second element is given and –ide ending. • Greek prefixes are used to indicate the number of atoms of each element.

  40. Some Simple Organic Compounds. • Hydrocarbons • Alkanes: • Alcohols:

  41. Chapter 3 Stoichiometry

  42. Chemical Equations • H2 + O2 H2O • Types of chemical reactions: • Combination (or synthesis): • Mg(s) + O2(g) MgO(s) • Decomposition: • CaCO3(s) CaO(s) + CO2(g) • Combustion: • C3H8(g) + O2(g)  CO2(g) + H2O(g)

  43. Formula Weights (aka molar mass) • The formula weight of a substance is the sum of the atomic masses of each atom in it’s chemical formula. • Percent composition is simply the percent that each element in a compound contributes to the total formula weight. • Calculate the percent composition of C12H22O11 • 42.1% - C • 6.4% - H • 51.5% - O

  44. Avogadro’s Number and the Mole • The mole is just a counting number. • Just like there are 12 things in a dozen there are…. • 6.02 x 1023 things in a mole. • 1 mole of Carbon atoms = 6.02 x 1023 carbon atoms • 1 mole of H2O molecules = 6.02 x 1023 H2O molecules • 1 mole of NO3- ions = 6.02 x 1023 NO3- ions

  45. Empirical Formula • To find the empirical formula of a compound we need to know its percent composition. • Mercury and chlorine combine to for a compound that is 73.9% mercury, and 26.1% chlorine.

  46. Molecular Formula • If we know a substances molar mass we can determine if it’s empirical formula is the same as its molecular formula. • The empirical formula of a hydrocarbon was found to be C3H4. The molar mass of this compound was found to be 121 g/mol. Is the empirical formula also the molecular formula? If not what is the molecular formula.

  47. Stoichiometry • 2H2(g) + O2(g)  2H2O(l) • How many moles of H2O can be produced from 1.57 moles of O2

  48. 2 C4H10(l) + 13 O2(g)  8 CO2(g) + 10 H2O(l) • Calculate the mass of CO2 produced when 1.00g of C4H10 is burned.

  49. Limiting Reactants • Sometimes one reactant runs out before the reaction is complete. • 2 H2(g) + O2(g)  2H2O(l) • If 10 moles of H2 and 7 moles of O2 react

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