Ch 1 Chemical Foundations

1. Ch 1 Chemical Foundations AP Chemistry 2014-2015

2. 1.1 Chemistry: An Overview • Matter is anything that takes up space and exhibits inertia. • Composed of only ~100 types of atoms (ex. water is made of hydrogen and oxygen; running an electric current through it separates it into its constituent elements) • Chemistry is the study of matter and energy—and more importantly, the changes between them

3. 1.2 The Scientific Method • Observations • Measurement = quantitative observation • No number involved = qualitative observation • A hypothesis is a possible explanation for an observation • Tested by carrying out an experiment • Repetition of experiments is key • A theory comes into existence when hypotheses are assembled in an attempt to explain “why” the “what” happened • We use many models to explain natural phenomena • A scientific law is a summary of observed behavior • Law of conservation of mass, law of conservation of energy, etc.

4. 1.3 Units of Measurement • A quantitative observation (measurement) consists of two parts: a number and a unit • Units • English (U.S., some of Africa) • Metric (everyone else, pretty much) • SI (Le Système International)—developed in 1960 to improve communication between scientists around the world; based on/derived from metric system

5. Units Continued • Volume (derived from length) • 1 dm3 = 1 L = 1,000 cm3 = 1,000 mL • 1 cm3 = 1 mL = 1 g of water at 4°C • Mass vs. Weight • Mass (g or kg)—a measure of the resistance of an object to a change in its state of motion (a measure of inertia); the quantity of matter in an object • Weight (Newtons)—the response of mass to gravity • Gravity varies with altitude; higher at low altitude, lower at high altitude • Every object has a gravitational field proportional to its mass • Precision and accuracy • Two types of error • See Exercise 1.2

6. 1.4 Significant figures and calculations • Rules • Nonzero digits are significant • A zero is significant if it is • Terminating and right of the decimal (must be BOTH) • “Sandwiched” between significant figures • Exact or counting numbers have an infinite amount of significant figures as do fundamental constants • See Exercise 1.3

7. Rules for calculating with sig figs • Multiplication and division: the term with the least number of sig figs determines the number of sig figs in the answer • Ex. 4.56 x 1.4 = 6.38  6.4 • Addition and subtraction: the term with the least number of decimal places determines the number of sig figs in the answer • Ex. 12.11 + 18.0 + 1.013 = 31.123  31.1 • pH calculations: the number of sig figs in the least accurate measurements determines the number of decimal plces on the reported pH • Round at the end of all calculations

8. 1.6 Dimensional Analysis • Consider a pin measuring 2.85 cm in length. What is its length in inches? 2.54 cm = 1 inch To convert, multiply your quantity by a conversion factor that “cancels” the undesired unit and puts the desired unit in its place. 2.85 cm x (1 inch/2.54 cm) = 1.12 inches

9. Exercise 1.5 Unit Conversions I

10. Exercise 1.6 Unit Conversions II

11. Exercise 1.7 Unit Conversions III

12. Exercise 1.8 and 1.9

13. 1.7 Temperature • Three scales • Fahrenheit TF = TC x (9°F/5°C) + 32°F • Kelvin TK = TC + 273.15 K • Celsius TC = TK - 273.15 K • See Exercise 1.10 Temperature Conversions

14. 1.8 Density • Density = mass/volume • See Exercise 1.13 Determining Density

15. 1.9 Classification of Matter • Solid, liquid, gas • Fluids = liquids and gases • Vapor: the gas phase of a substance that is normally a solid or liquid at room temperature; for example, we say “water vapor” but we don’t say “oxygen vapor”

16. Classification Continued • Mixtures: can be physically separated • Homogeneous • Heterogeneous • Means of physical separation include filtering, fractional crystallization, distillation, chromatography • Pure substances: compounds and elements • Compounds can be separated into elements by chemical means (ex. electrolysis) • Elements can be broken down into atoms, which can be broken down into the nucleus and electron cloud, which can be broken down into protons, neutrons, and electrons, which can be broken down into quarks and leptons

17. Ch2 Atoms, Molecules, and Ions AP Chemistry 2014-2015

18. 2.1 Contains historical information.   • Read it if you would like.

19. 2.2 Fundamental Chemical Laws • Antoine Lavoisier was the first chemist to insist on quantitative experimentation. He got guillotined, but not for that reason. • The law of conservation of mass: matter is neither created nor destroyed. • The law of definite proportions: a given compound always contains exactly the same proportions of elements by mass. • The law of multiple proportions: when two elements combine to form a series of compounds, the ratios of the masses of the second element that combine with one gram of the first element can always be reduced to small whole numbers. You can see an example of this on the right. The likely formulas for these compounds would be CO and CO2.

20. Exercise 2.1 Illustrating the Law of Multiple Proportions The following data were collected for several compounds of nitrogen and oxygen: Mass of Nitrogen That Combines With 1 g of Oxygen Compound A 1.750 g Compound B 0.8750 g Compound C 0.4375 g Show how these data illustrate the law of multiple proportions.

21. 2.3 Dalton’s Atomic Theory • Dalton’s Theory (partially correct, partially not) • All matter is made of atoms. These indivisible and indestructible objects are the ultimate chemical particles. • All the atoms of a given element are identical, in both weight and chemical properties. However, atoms of different elements have different weights and different chemical properties. • Compounds are formed by the combination of different atoms in the ratio of small whole numbers. • A chemical reaction involves only the combination, separation, or rearrangement of atoms; atoms are neither created nor destroyed in the course of ordinary chemical reactions • Two modifications were made when subatomic particles and isotopes were discovered.

22. Avogadro’s Hypothesis • At the same temperature and pressure, equal volumes of different gases contain the same number of particles.

23. 2.4 Early Experiments to Characterize the Atom • The electron • J.J Thomson found that when high voltage was applied to an evacuated type, a “ray” he called a cathode ray was produced. The ray was produced at the electrode (also called the cathode) and was repelled by the negative pole of an applied electric field. He postulated that the ray was a stream of negative particles (now called electrons). He then measured the deflection of beams of electrons to determine the charge-to-mass ratio. Thomson discovered that he could repeat this deflection and calculations using different metal electrodes, showing that all metals contain electrons and all atoms contain electrons. He also deduced that since atoms were neutral, there must be a positive charge within the atom, giving rise to the “plum pudding” model.

24. Millikan’s oil drop experiment • Next up, Robert Millikan sprayed charged oil drops into a chamber. He halted their fall (due to gravity) by adjusting the voltage across two charged plates. He used the stop-drop voltage and Thomson’s charge-mass ratio to determine the charge on one drop of oil, which was a whole number multiple of the electron charge. • The mass of an electron is 9.11 x 10-31 kg.

25. Radioactivity • Henry Becquerel famously (and accidentally) discovered radiation when he left a uranium ore in a closed drawer with a photographic plate. When he realized that the plate had been exposed, he realized that a form of radiation other than light had penetrated it. The uranium, of course, was the culprit.

26. Radioactivity Continued • Three types of radioactive emission • Alpha (particles): helium nuclei, relatively massive and slow, poorly penetrating, somewhat dangerous • Beta (particles): electrons, relatively light and fast, moderately penetrating, a little more dangerous • Gamma (rays): just energy, most penetrating, most dangerous • These are not the only kinds of radioactive emission. We will discuss more in the spring.

27. The nuclear atom • Rutherford’s famous gold foil experiment proved that a positively-charged and somewhat bulky nucleus could be found in the center of an atom. He also found that atoms are mostly empty space.

28. 2.5 The Modern View of Atomic Structure (an introduction) • Elements • All matter composed of only one type of atom is an element. 92 elements are naturally-occurring; the rest are manmade. • Atoms • The atom is the smallest particle of an element that retains the chemical properties of that element. It consists of a bulky, dense nucleus (protons and neutrons) and electrons shells/clouds (which of course contain electrons).

29. Atoms and Isotopes • We can find a few pieces of information about each element using isotope notation. • Mass number = #protons + #neutrons for specific isotopes of an element • Actual mass is not an integral number! mass defect--causes this and is related to the energy binding the particles of the nucleus together • Atomic number = #protons = #electrons in a neutral atom = identity of the element

30. Exercise 2.2 Writing the Symbols for Atoms Write the symbol for the atom that has an atomic number of 9 and a mass number of 19. How many electrons and how many neutrons does this atom have?

31. Isotopes • Isotopes are atoms that have the same number of protons (and therefore are the same element) but different numbers of neutrons (and therefore different masses). • Most elements have at least two stable isotopes. Exceptions include Al, F, P. • Hydrogen isotopes are important because they have special names. • 0 neutrons = hydrogen • 1 neutron = deuterium • 2 neutrons = tritium

32. 2.6 Molecules and Ions • Electrons are responsible for bonding and chemical reactivity. • Chemical bonds—forces that hold atoms together • Covalent bonds—atoms share electrons and make molecules [independent units]; H2, CO2, H2O, NH3, O2, CH4 to name a few. • Molecule--smallest unit of a compound that retains the chem. characteristics of the compound; characteristics of the constituent elements are lost. • Molecular formula--uses symbols and subscripts to represent the composition of the molecule. (Strictest sense--covalently bonded)

33. Molecules and Ions Continued • Structural formula—bonds are shown by lines [representing shared e- pairs]; do not always indicate shape • Ions--formed when electrons are lost or gained in ordinary chem. reactions; dramatically affect size of atom • Cations--(+) ions; often metals since metals lose electrons to become + charged • Anions--(-) ions; often nonmetals since nonmetals gain electrons to become - charged • Polyatomic ions--units of atoms behaving as one entity--MEMORIZE formula and charge! • Ionic solids—Electrostatic forces hold ions together. Strong ions held close together solids.

34. 2.7 An Introduction to the Periodic Table • Metals—malleable, ductile & have luster; most of the elements are metals—exist as cations in a “sea of electrons” which accounts for their excellent conductive properties; form oxides [tarnish] readily and form POSITIVE ions [cations]. Why must some have such goofy symbols?

35. Periodic Table Continued • Groups or families--vertical columns; have similar physical and chemical properties (based on similar electron configurations!!) • Group A—Representative elements • Group B--transition elements; all metals; have numerous oxidation/valence states • Periods --horizonal rows; progress from metals to metalloids [either side of the black “stair step” line that separates metals from nonmetals] to nonmetals

36. Memorize • ALKALI METALS—1A • ALKALINE EARTH METALS—2A • HALOGENS—7A • NOBLE (RARE) GASES—8A

37. 2.8 Naming Simple Compounds • Binary Ionic Compounds (Type I and Type II) • In general—consist of a metal cation and a nonmetal anion. • The cation is written first. The charges from the cation and anion must cancel; we use subscripts to make this happen. • The names of ionic compounds do not contain prefixes such as mono- or di- unless that is part of the name of a polyatomic ion in the compound. • Monatomic ions end in –ide. Ex. NaF is sodium fluoride.

38. Type I Binary Ionic Compounds • Type I contain non-transition metals, which have only one charge when they are cations. • Group 1A = +1, Group 2A = +2, Aluminum = +3 • Zinc, silver, and cadmium also fit into this category; silver ions always have a +1 charge, while zinc and cadmium ions always have a +2 charge. • Writing the name of a Type I Binary Ionic compound is simple. Ex. MgCl2 is magnesium chloride. The formulas are also simple, but you have to swap-and-drop to get the correct formula. Ex. sodium oxide is Na2O, and calcium nitride is Ca3N2.

39. Type II Binary Ionic Compounds • Type II contain transition metals, as well as a few others such as lead, tin, and mercury. • These ions have variable charges which are reflected in the formula using roman numerals. For example, FeCl3 would be iron (III) chloride and SnO2 would be tin (IV) oxide. Conversely, lead (II) chloride would be PbCl2. • Some of them are real weirdoes. For example, the mercury (II) ion is Hg2+ which makes sense, but the mercury (I) ion is Hg22+.

40. Exercise 2.3 Naming Type I Binary Compounds Name each binary compound: • CsF • AlCl3 • LiH

41. Exercise 2.4 Naming Type II Binary Compounds Give the systematic name of each of the following compounds. • CuCl • HgO • Fe2O3 • MnO2 • PbCl2

42. Exercise 2.5 Naming Binary Compounds Give the systematic name of each of the following compounds. • CoBr2 • CaCl2 • Al2O3 • CrCl3

44. Ionic Compounds With Polyatomic Ions • Same as the other ionic names/formulas we’ve seen, but you need to look out for polyatomic ions. I’ve given you a sheet of them, but you will not be given a list for the AP Exam.

45. Exercise 2.6 Naming Compounds Containing Polyatomic Ions Give the systematic name of each of the following compounds. • Na2SO4 • KH2PO4 • Fe(NO3)3 • Mn(OH)2 • Na2SO3

46. Exercise 2.6 continued • Na2CO3 • NaHCO3 • CsClO4 • NaOCl • Na2SeO4 • KBrO3

47. Binary Covalent Compounds • Consist of two nonmetals bonded together • Use prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca • Don’t forget the –ide ending

48. Exercise 2.7 Naming Type III Binary Compounds Name each of the following compounds. • PCl5 • PCl3 • SF6 • SO3 • SO2 • CO2

50. Acids • Hydrogen is listed first in the formula; the anion is listed second • -ide →hydro [negative ion root]ic ACID • -ate →-ic ACID • -ite → -ous ACID