1. Unit 1 MR. GATES AP CHEMISTRY

2. Scientific Method  Identify a problem/question.  Making Observations: these could be qualitative (don’t involve numbers) or they could be quantitative (do involve numbers).  Quantitative observations involve a number and a unit.  Formulating Hypotheses: a possible explanation for an observation  Performing Experiments: experiments are used to test a hypothesis by producing new information that reveals if the hypothesis is valid.  Analyze results and repeat

3. Scientific method cont.  Scientists use experimental results to test scientific models.  When experimental results are not consistent with predictions of a scientific model, the model must be revised or replaced with a new model that is able to predict/explain the new experimental results.  If the model isn’t working, change is needed  A robust scientific model is one that can be used to explain/predict numerous results over a wide range of experimental circumstances.  Better models prove to be true under many circumstances and changing variables  Can withstand the test of many experiments

4. Pureblood vs. Mudblood  A pure substance/sample contains particles (or units) of one specific atom or molecule  A mixture contains particles (or units) of more than one specific atom or molecule  Homogeneous mixtures are of a uniform composition throughout.  These are called solutions  Heterogeneous mixtures are NOT of a uniform composition throughout.

5. Chemical Changes  Chemical changes are ones in which a given substance becomes a new substance or substances with different properties and different composition.  Indications of a chemical change:  Color change  Change in temperature  Release of a gas  Formation of a precipitate

6. Physical Changes  A physical change is one in which a substance changes its state, but remains the same substance with the same composition.  Examples:  Freezing,  Boiling,  Melting,  Sublimation,  Etc.

7. Chemical and Physical Changes cont.  The distinction between chemical and physical processes relates to the nature of the change in molecular interactions.  Processes that involve the breaking and/or formation of chemical bonds are classified as chemical processes (changes).  Processes that involve only changes in weak intermolecular interactions, such as phase changes, are classified as physical processes (changes).  A gray area exists between these two extremes.  Ex. The dissolution of a salt in water involves breaking of ionic bonds and the formation of interactions between ions and solvent.  The magnitude of these interactions can be comparable to covalent bonds strengths, and so plausible arguments can be made for classifying dissolution of a salt as either a physical of chemical process.

8. Law of Definite Proportions  Compounds have a constant composition  Because the molecules of a particular compound are always composed of the identical combination of atoms in a specific ratio, the ratio of the masses of the constituent elements in any pure sample of that compound is always the same.

9. Law of Multiple Proportions  When two elements for a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers.  Pairs of elements that form more than one type of molecule are nonetheless limited by their atomic nature to combine in whole number ratios.  This discrete nature can be confirmed by calculating the difference in mass percent ratios.

10. Ionic Formulas and Covalent Formulas  The subscripts in a chemical formula represent the number of atoms of each type in a molecule.  Whenever there is a transition metal, the written name should have roman numerals in parentheses after the metal. The numeral refers to the charge of the metal.  Ex. Iron (III) oxide  To write the formula of an ionic compound with a transition metal, the net charge of each side must be calculated and balanced.  Iron (III) has a +3 charge, and oxygen has a -2 charge so to balance them, we need two iron atoms for every three oxygen atoms  Fe 2 O 3

11. Naming Acids  No oxygen: start with hydro- and end with –ic  HCl is called hydrochloric acid  With Oxygen: -ate becomes –ic, and –ite becomes –ous  SO 4 -2 sulfate ion is named sulfuric acid as H 2 SO 4  SO 3 -2 sulfite ion is named sulfurous acid as H 2 SO 3

12. Atomic Mass  The average mass of any large number of atoms of a given element is always the same for a given element.  The average mass of an element is the atomic mass  Expressing the mass of an individual atom or molecule in atomic mass unit (amu) is useful because the average mass in amu of one particle (atom or molecule) of a substance will always be numerically equal to the molar mass of that substance in grams.

13. Mass Spectrometer  The most accurate way of measuring masses of particles.  An applied electric field accelerates ions into a magnetic field. Since each moving in also creates its own magnetic field, the two interact with each other causing the ion to stray from its path. Using the variation the ion took from its path, we can calculate the mass of the ion.

14. #Avogadro’s#  The molar mass of a substance is the mass in grams of a mole of particles of that substance.  in 12 grams of carbon, there are 6.022x10 23 atoms of carbon…  And thus Avogadro's number was born.  Avogadro’ number provides the connection between the number of moles in pure sample of a substance and the number of constituent particles (or units) of that substance.

15. And they all lived connected ever after  Thus for any sample of pure substance, there is a specific numerical relationship between the molar mass of the substance, the mass of the sample, and the number of particles (or units) present.

16. Percent Composition and Empirical Formula  Because compounds are composed of atoms with known masses, there is a correspondence between the mass percent of the elements in a compound and the relative number of atoms of each element.  An empirical formula is the lowest whole number ratio of atoms in a compound. Two molecules of the same elements wit the same identical mas percent of their constituent atoms will have identical empirical formulas.

17. Chemical Equations  Various types of representations can be used to show that matter is conserved during chemical and physical processes.  Symbolic representations  Particulate drawings  Pg. 97 shows good examples of these.

18. Balancing Chemical Equations  Atoms and molecules interact with one another on the atomic level. Balanced chemical equations give the number of particles that react and the number of particles produced. Because of this, expressing the amount of a substance in terms of the number of particles, or moles of particles is essential to understanding chemical processes.  The coefficients in a balanced chemical formula represent the number of atoms of each type in a molecule.

19. Balancing Equations  The concept of conservation of atoms plays an important role in the interpretation and analysis of many chemical processes on the macroscopic scale.  Chemical equations represent chemical changes, and therefore must contain equal numbers of atoms of every element on each side to be “balanced.”

20. Balancing Chemical Equations  Reactants:  1 Mg  1 N  1 P  4 H  4 O  Products:  2 Mg  1 N  2 P  5 H  8 O

21. Balancing Chemical Equations  Reactants:  2 Mg  2 N  2 P  8 H  8 O  Products:  2 Mg  1 N  2 P  5 H  8 O

22. Balancing Chemical Equations  Reactants:  2 Mg  2 N  2 P  8 H  8 O  Products:  2 Mg  2 N  2 P  8 H  8 O

24. Stoichiometry

28. Stoich. and the limits there of  Coefficients of balanced chemical equations contain information regarding the proportionality of the amounts of substances involved in the reaction. These values can be used in chemical calculations that apply the mole concept; the most important place for this type of quantitative exercise is the laboratory.  Calculate amount of product expected to be produced in a laboratory experiment.  Identify limiting and excess reactant; calculate percent and theoretical yield for given laboratory experiment.

29. Limiting Reagent, Theoretical Yield, and Excess  Ex: What is the limiting reactant if 8.0 mol of HF reacts with 4.5 mol of SiO 2 ? SiO 2 + 4 HF  SiF 4 + 2 H 2 O  a) Limiting Reactant?  b) Amount of Products being produced (theoretical yield)?  c) How much of the Excess Reactant is left over?

30. Limiting Reactant a: What is the limiting reactant if 8.0 mol of HF reacts with 4.5 mol of SiO 2 ? SiO 2 + 4 HF  SiF 4 + 2 H 2 O  Limiting Reactant? 8.0 mol HF 2 mol H 2 O = 4.0 mol H 2 O 4 mol HF 4.5 mol SiO 2 2 mol H 2 O = 9.0 mol H 2 O 1 mol SiO 2

31. Theoretical Yield b: What is the limiting reactant if 8.0 mol of HF reacts with 4.5 mol of SiO 2 ? SiO 2 + 4 HF  SiF 4 + 2 H 2 O  Amount of Products Produced (Theoretical Yield)? 8.0 mol HF 2 mol H 2 O = 4.0 mol H 2 O 4 mol HF 8.0 mol HF 1 mol SiF 4 = 2.0 mol SiF 4 4 mol HF

32. Excess c: What is the limiting reactant if 8.0 mol of HF reacts with 4.5 mol of SiO 2 ? SiO 2 + 4 HF  SiF 4 + 2 H 2 O  How much of the excess reactant is left over? 8.0 mol HF 1 mol SiO 2 = 2.0 mol SiO 2 4 mol HF 4.5 mol SiO 2 - 2.0 mol SiO 2 = 2.5 mol SiO 2

33. Percent Yield