Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
AP / College Chem Chapter 2 – Atoms, Molecules & Ions - Atomic Theory of Matter: o Democritus (400 B.C.) was a Greek philosopher who came up with the first concept of the atom (atomos – indivisible) o Plato and Aristotle did not believe in atoms (no limit to how small you could divide matter) o John Dalton (1803 – 1807) came up with the modern atomic theory Each element is composed of extremely small indivisible particles called ATOMS All atoms of the same element are identical but atoms of different elements are different from each other Atoms of one element cannot be changed into atoms of a different element by chemical reactions (ATOMS are neither created or destroyed!) Compounds are formed when atoms of a more than one element combine and a given compound always has the same relative number and kind of atoms (law of definite composition) o Explains the law of conservation of matter (the total mass of the reactants must equal the total mass of the products) o Dalton also used his theory to develop the law of multiple proportions (if two elements A and B combine to form more than one compound, the masses of B that combine with a given mass of A are in a ratio of small whole numbers) - Discovery of Atomic Structure: o We know today that atoms are not indivisible like Dalton thought but are composed of subatomic particles (protons, neutrons and electrons) o Particles with the SAME charge repel each other and particles with OPPOSITE charges attract each other o Cathode Rays and Electrons mid 1800s scientist were studying cathode ray tubes (glass tubes filled with a gas that had a cathode (-) and anode (+) end) when high voltage was applied to the electrodes, radiation known as CATHODE RAYS were produced that would cause certain materials to emit light (FLUORESCE) cathode rays discovered by Johann Hittorf and Eugen Goldstein named cathode rays in 1876 experiments showed that cathode rays were deflected by electric or magnetic fields in a way that was consistent with their being a stream of negative electrical charge In 1897, J.J. Thompson describe cathode rays as a stream of negatively charged particles that behaved exactly the same no matter what material was used for the cathode Credited with the discovery of the electron Cathode rays (electrons) begin at the cathode and are accelerated toward the anode which has a hole in the center. This produces a beam of electrons that travel toward a fluorescent screen. The strengths of the electric and magnetic fields are adjusted so their effects cancel each other and allow the beam to travel along a straight path (to calculate the charge to mass ratio Thompson also found the electrical charge to mass ratio of an electron to be 1.76 x 108 Coulombs/gram Robert Millikan (1909) measured the charge of an electron via the Oil-Drop Experiment Small drops of oil fall between electrically charged plates and gain electrons after exposure to X-rays. Millikan varied the voltage between the plates and measured the rate of fall. From this data he calculated the charge of a single electron to be 1.602 x 10-19 Coulombs (C). Combining data from these two experiments, the mass of an electron was calculated 1.602 x 10-19 C Electron mass = ------------------------ = 9.10 x 10-28 g 1.76 x 108 C/g This result means the electron is almost 2000 times smaller in mass than a Hydrogen atom!! o Radioactivity Henri Becquerel (1896) discovered that Uranium spontaneously emits high-energy radiation called RADIOACTIVITY Marie & Pierre Curie worked on isolating the radioactive components of that compound that Becquerel worked on Ernest Rutherford determined 3 different types of radiation (alpha, beta and gamma) and showed that alpha and beta rays consisted of fast moving particles Alpha (He nucleus and has a +2 charge), Beta is fast moving electrons and are (-) charged and gamma have no charge and are similar to X-rays o Nuclear Model of the Atom: J.J. Thompson came up with the Plum Pudding Model of the atom (atom was a mass of (+) charge with (-) electrons dispersed throughout like plum pudding) Ernest Rutherford (1910) studied angles at which alpha particles were deflected or scattered through a thin sheet of gold foil He had his assistant (Marsden) look for greater angles and surprisingly found some particles scattered backwards which was totally inconsistent with Thompson’s model Rutherford came up with the NUCLEAR MODEL in which the protons are located in a small, dense mass called the nucleus and the negative electrons located in the mostly empty space outside of the nucleus James Chadwick (1932) discovered the neutron (neutral particle) - The Modern View of Atomic Structure: o Every atom is made of protons and neutrons in the nucleus of the atom and electrons located outside the nucleus in mostly empty space o The charge of an electron is assigned a value of (-1) and a proton (+1) making them equal in charge but opposite in sign (even though protons are MUCH more massive than electrons are the magnitude of the charges are the same!!) o Every atom has the SAME number of protons and electrons making the atom electrically NEUTRAL o Atomic mass unit (AMU) – the mass unit of the atom o 1 amu = 1.66 x 10-24 grams o 1 proton is 1.0073 amu o 1 neutron is 1.0087 amu o 1 electron is 5.486 x 10-4 amu (so that means that it takes 1,836 electrons to equal the mass of a proton!) o Atomic size is usually measured in ANGSTROMS (Å) which is 1 x 10-10 m o Atoms have diameters anywhere from 1 – 5 Å o Electromagnetic Forces are determined by COULOMB’S LAW F is the electromagnetic force of attraction (-) or repulsion (+) between 2 charged particles. It is proportional to the magnitude of the charges and inversely proportional to the square of the distance (r) between the centers of the charged particles o Atomic Numbers, Mass Numbers & Isotopes Atomic number (Z) – the total number of protons in the nucleus of the atom The atomic number identifies an element Mass number (A) – the total number of protons AND neutrons in the nucleus of the atom A = # protons + # neutrons # neutrons = A – Z The mass number identifies an ISOTOPE (different form of the same element that has a different mass because of a different number of neutrons) There are at least 2 isotopes (sometimes more) for each element - Atomic Weights o Atomic Mass Scale Atoms of different elements have different masses Early scales based on H having a mass of 1, then O having a mass of 16 After discovering isotopes it became clear that an isotope needed to be the standard for atomic mass so C-12 was used and given a mass of 12.0000 amu Average atomic mass is calculated as a weighted average Calculate the average atomic mass of Chlorine based on the following data 35 Cl is 75.78% abundance with a mass of 34.969 amu 37 Cl is 24.22% abundance with a mass of 36.966 amu Average mass = (0.7578)(34.969 amu) + (0.2422)(36.966 amu) = 26.50 amu + 8.953 amu = 35.45 amu Relative abundance and masses of isotopes are determined by a MASS SPECTROMETER Produces a MASS SPECTRUM that shows mass and abundance Used to identify chemical compounds and analyze mixtures of substances - The Periodic Table: o Most significant tool that chemists use for organizing and remembering chemical facts o When elements are arranged in order of increasing atomic number, their chemical and physical properties show a repeating (periodic) pattern o PERIODS – horizontal rows (7) correspond to the 7 main energy levels of electrons in the atom o GROUPS – vertical columns (18) correspond to a family of elements that share similar chemical and physical properties mostly because their outer energy level electron structure is the SAME as all other members of that group or family o Metals – located at the left side and middle of the periodic table o Nonmetals – located at the right side of the periodic table o Metalloids – semiconductors or hybrids that are located along a stepwise dividing line in the periodic table (B, Si, Ge, As, Sb, Te) o Metalloids have some properties of metals and some of nonmetals - Molecules and Molecular Compounds: o Molecules have 2 or more atoms combined together by covalent bonds (sharing pairs of electrons) o Molecules may be elements (DIATOMIC ELEMENTS – H2, O2, N2, Cl2, Br2, I2, F2) o Molecules may also be molecular compounds that are sharing electrons between atoms of different elements (H2O, NH3, CO2, etc) o Molecules are almost always made of ONLY NONMETALS (and sometimes metalloids) o Molecules are always identified by their MOLECULAR FORMULA (shows the exact number of each atom in one molecule) that is NEVER reduced as a formula o Empirical formula – shows a chemical formula with the SMALLEST WHOLE NUMBER RATIO of elements in the formula (may be reduced) o Empirical formulas are NOT typically used for molecules - Ionic Compounds: o Ionic compounds are made from the attraction of (+) and (-) ions which are formed by a TRANSFER of electrons from a metal to a nonmetal atom o Results in FULL charges and is a much stronger bond than a covalent bond o Periodic table is very helpful in determining the CHARGE on a monatomic ion (can figure out from what group the element is in) o Ions in ionic compounds are arranged in 3-D structures called a crystal lattice and therefore only EMPIRICAL formulas are used to show the smallest whole # ratio of ions in the compound o Each Na+ in sodium chloride is surrounded by 6 Cl- and each Cl- is surrounded by 6 Na+ so we actually have a 1:1 ratio NaCl o Although ionic compounds are made up of (+) and (-) ions, the compounds themselves have a total charge of ZERO and are NEUTRAL o That means the (+) and (-) charge has to balance so when figuring out the empirical formula for an ionic compound we CRISS-CROSS the charges to become the subscripts on the opposite ion and then reduce the formula (if needed) o If you have more than one unit of a polyatomic ion it must be placed in PARENTHESIS with the subscript after the parenthesis o No parenthesis are needed if your subscript is ONE on a polyatomic ion - Naming Inorganic Compounds: o Chemical nomenclature is used for naming compounds o System of assigning a unique name to every individual compound o Some compounds are known for their common names (NH3 – ammonia) o Organic molecules have their own system for naming o Inorganic based on the following system and depend on the type of compound that is involved o Names and Formulas of IONIC compounds: Naming Ionic Compounds is easy – they are all made of (+) and (-) ions so to name any ionic compound--Name the (+) ion first and the (-) ion last The trick is knowing HOW to name the IONS!!! CATIONS formed from metal atoms have the same name as the metals plus the word “ION” If a metal can form CATIONS with different charges (most transition metals and post-transition metals) then the (+) charge is indicated by a ROMAN NUMERAL in parentheses following the name of the metal Old system used suffixes (“-IC” for HIGHER charge and “-OUS” for LOWER charge) and the Latin root name of the metal CATIONS formed from nonmetal atoms that end in “-IUM” COMMON CATIONS ANIONS of monatomic ions are named by replacing the ending of the element name with “-IDE” A few POLYATOMIC IONS also end in “-IDE” POLYATOMIC IONS that contain O (OXYANIONS) have the endings “-ATE” (MORE) and “-ITE” (LESS) to indicate how many oxygen atoms the ion contains When there is more than just 2 possibilities of number of O atoms then the prefixes “PER-“ (ABOVE) and “HYPO-“ (UNDER) are used Procedure for Naming Anions Anions derived by adding H+ are named by adding the prefix hydrogen or dihydrogen. The old method uses the prefix “bi-“ Names of Common Anions Names of IONIC COMPOUNDS are named by naming the (+) ion then the (-) ion o Names and Formulas of ACIDS: Acid is a molecule that produces H+ when dissolved in water Generic formula: HA Acids containing anions whose names end in “-IDE” are named by using the prefix “HYDRO-“ then the root of the anion and change the ending to “-IC ACID” Acids containing anions whose names end in “-ATE” or “-ITE” (OXYACIDS) are named by changing ATE IC and ITE OUS and then adding the word ACID o Names and Formulas of BINARY MOLECUES: Name the element that is farther LEFT on the Periodic Table (except when O is combined with a HALOGEN, the halogen is named first) If both elements are in the same group, the LOWER one is named first The name of the second element is given the “-IDE” ending Greek prefixes indicated the number of atoms of each element in ONE MOLECULE. (Mono- is never used at the start of the name!) Prefixes used for Molecules: - Some Simple Organic Compounds: o Organic chemistry—study of compounds containing Carbon o Also contain H, and sometimes O, N, or halogens o Alkanes Hydrocarbons—contain only C and H Alkanes – simplest hydrocarbons containing only C—C single bonds Alkanes are SATURATED hydrocarbons Names end in “-ANE” o Some Derivatives of Alkanes Alcohol—has the functional group –OH in the formula Names end in “-OL Compounds with the same molecular formula but different structural formulas are called ISOMERS - Suggested Problems: 2.1, 2.2, 2,3 2,5, 2,7, 2.9, 2.11, 2.12, 2.13, 2.15, 2.17, 2.19, 2.20, 2.21, 2.22, 2.23, 2.25, 2.27, 2.29, 2.31, 2.35, 2.37, 2.41, 2.42, 2.43, 2.45, 2.47, 2.49, 2.50, 2.51, 2.53, 2.55, 2.57, 2.59, 2.61, 2.63, 2.65, 2.67, 2.69, 2.71, 2.74, 2.77, 2.79