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CHAPTER 2 ATOMS, MOLECULES AND IONS 6/2015 ATOMIC THEORY OF MATTER: 1. Democritus 400BC Greek Through logical deduction theorized that matter is made up of small particles that cannot be divided. Called these particles “atomos” His idea was dismissed until the 1700’s 2. John Dalton’s Atomic Theory, 1805 - Based on his reading and observations in the lab Dalton’s Atomic Theory of Matter ( 1803) a. Each element is composed of very small, indivisible particles called atoms (Democritus) b. All atoms of an element are identical, but they differ from the atoms of other elements. c. Atoms are neither created or destroyed in chemical reactions (Law of Conservation of Matter /Mass– Antoine Lavoisier) d. A given compound always contains the same elements in the same proportions, by mass. (Law of Constant Composition , or Law of Definite Proportions, John Proust, 1799) Dalton used his theory to propose the Law of Multiple Proportion: If two elements, (ex A & B) form more than one compound, then the ratio of the second element that reacts with a fixed amount of the first element will always be in ratios of small whole numbers. Ex H2O and H2O2 H2O For every 2 g H : 16 g O H2O2 For every 2 g H : 32 g O The ratio of oxygen in each compound is 1:2 More advanced technology and application of knowledge of electrical charges were used to explore the structure of atoms. Two types of electrical charges , + and – Like charges repel Opposite charges attract a. Cathode Ray Tubes: closed glass tubes with electrons at each end, and containing little air. * when electricity is applied , radiation was produced that travelled from the cathode (- end) to the anode (+ end), causing the gas in the tube to glow. 3. JJ Thomson, 1897 Found he could move a cathode ray in a cathode ray tube using a magnet or electrical field Concluded that the cathode ray was made up of a stream of neg charged particles with a mass smaller than hydrogen, he named them electrons. This was a breakthrough in that it was realized that subatomic particles exist. Thompson published a mass: charge ratio for the electron of 1.76 x 10-28 C/g He developed the Plum Pudding model of atom – the atom was a large, positively charged sphere, with negative charges spread throughout . 4. Robert Millikan, 1909 Determined charge and mass of an electron using an oil drop experiment 1.60 X 10-19 coulomb, 9.11 x 10-28 g 5. Henri Becquerel, 1896 Discovered radioactivity – the spontaneous emission of radiation from the nucleus of an atom. (He observed Uranium) 6. Later , in the early 1900’s Ernst Rutherford was able to separate radiation into 3 types using a focused stream of radiation, charged electrical plates, and a detection screen. Alpha particles ( α ) with a charge of +2 Beta particles (β) with a charge of -1 Gamma radiation (α), no charge; no mass 7. Rutherford and Associates, 1909 Gold foil experiment http://phet.colorado.edu/en/simulation/rutherfordscattering http://www.youtube.com/watch?v=XBqHkraf8iE Evidence most alpha particles went straight through foil Conclusion the atom is mostly empty space some alpha particles bounced nearly straight back the atom contains a dense, positively charged nucleus Rutherford’s model of an atom: Rutherford and his associates concluded that (1) most of an atom’s mass and all of its positive charge is in a small core (the nucleus) and (2) most of the atom is empty space. 8. Rutherford discovered protons in 1919 9. James Chadwick discovered neutrons in 1932. Although subatomic particles smaller than protons, neutrons, and electrons exist, chemical behaviour can be explained by considering these three particles. STRUCTURE OF ATOMS – THE BASICS ATOM: the smallest particle of an element that shows the chemical properties of that element.) Three main sub particles make up an atom Particle Proton Charge +1 Mass 1 amu Location nucleus Neutron 0 1 amu nucleus Electron -1 .0006 amu (~0) orbitals around nucleus Atomic mass units (amu) * a unit of mass equal to 1/12 the mass of a Carbon-12 atom (which has a mass of 12 amu) 1 amu ~ 1.66 x 10-24 g protons & neutrons have masses of ~ 1 amu Atomic number: the # protons in its nucleus The number of protons in the nucleus determines the element. Atomic mass = # protons + # neutrons (the number of particles in the nucleus) Most of the mass of an atom is in the nucleus Single, unreacted atoms are neutral, meaning the number of protons equals the number of electrons, so the opposite charge values cancel out: Ex. Carbon 6 + (protons) 6 – (electrons) 0 no overall charge The number of neutrons in atoms of one element can vary. Isotopes : naturally occurring versions of an element that vary in the numbers of neutrons in their nucleus. Isotopes of an element will vary in mass because some have more or fewer neutrons. Isotopes of an element have the same chemical properties since these are determined by the number of protons and electrons. symbols for isotopes may be written 2 ways: name – mass Ex: Chlorine – 35 symbol - mass Cl-35 atomic mass atomic # Symbol 14 6C Average atomic mass: the “weighted” average of all naturally occurring isotopes of an element periodic table lists average atomic masses avg atomic mass is close to the most common isotope Calculating Average Atomic Mass: Steps to solve the problem: 1. Read problem: What is the average atomic mass of chlorine, if 75% of chlorine atoms have a mass of 35 and 25% have a mass of 37? 2. Organize your data into a table, matching mass and % abundance of each isotope Isotope mass 35 37 % abundance 75% 25% decimal .75 .25 3. Change % to a decimal by moving decimal point two places to the left (put in table) 4. Plug numbers into the equation AAM = (mass isotope1) (abundance of 1) + (mass isotope 2)(abundance of isotope 2) +… etc AAM = (35)(.75) + (37)(.25) = 35.45 amu Practice: Calculate the average atomic mass of Silicon. 92.21% of Si atoms have a mass of 27.977 amu, 4.70% have a mass of 28.976 amu and 3.09% have a mass of 29.974 amu. FOUR BASIC FORCES A basic understanding of the four basic forces is helpful in understanding the structure and reactivity of atoms. 1. Gravitational force: the attraction between two objects. It is directly related to mass and inversely proportional to distance between centers of the two objects. Subatomic particles are too small for gravity to have a significant effect on chemical behaviour. 2. Electromagnetic forces: attractive and repulsive forces that act between charged particles or magnetic objects. Like charges (poles) repel, opposite attract Similar to gravity in that it is directly proportional to size of the charge of magnetic force and inversely proportional to distance. 3. Strong nuclear force: inward force on the nucleus of an atom that works to keep the particles together. (Remember protons are all repelling each other). 4. Weak nuclear force: interacts with quarks, bosons, and fermions. Involved in beta decay. THE PERIODIC TABLE – or Why are the elements arranged like that? 1. Periodic Law: when elements are arranged by increasing atomic number, they show repeating patterns of chemical and physical properties. 2. Groups (aka Chemical family): are elements found in a column of the periodic table that share similar chemical (and often physical) properties. - They have similar chemical properties b/c they have the same number of electrons in their outermost energy level. - Groups may be labelled several ways Ex 1 – 18 or 1A – 8A and 1B – 8B 3. Some groups have special names 1 (1A) Alkali metals 2 (2A) Alkaline earth metals 17 (7A) Halogens 18 (8A) Noble gases 4. Periods: horizontal rows in the periodic table. These rows represent energy levels in the atoms. - An element in the 3rd period has electrons in the 1st, 2nd, and 3rd energy levels. 5. Metals: elements located on the left and middle section of the periodic table are metals - Exception is H, a non-metal Characteristics of metals include luster, ability to conduct heat; ability to conduct electricity. Most are solids at room temperature (exception being Hg,which is a liquid at RT) 6. Non-metals: located on the right side of the table. - Non-metals vary more than metals in their properties and characteristics - Many are gases, they are biologically important (C, H, O )most common element in living things. 7. Metalloids: are located along the stair-step line of the periodic table - they have characteristics between metals and non-metals - B, Si, Ge, As, Sb, Te 8. Information listed for each element includes Atomic number Symbol Name Atomic mass (average) Density Oxidation states Atoms react with each other to form molecules and ionic compounds in order to obtain more favourable, lower energy states. MOLECULES AND MOLECULAR COMPOUNDS Molecule: general name used for two or more non-metal atoms bonded together by sharing of electrons. Bonds formed by the sharing of electrons are called covalent bonds Several elements are found in nature in molecular form , meaning 2 or more of the same type of atom bonded together. Ex: Oxygen gas is O2 Ozone is O3 Diatomic molecules: molecules made up of two of the same type of atom bonded together. There are 7 elements that occur as diatomic molecules, all non-metals H2 N2 O2 F2 Cl2 Br2 I2 Molecular compounds : substances made of more than one type of atom. Most molecular substances are made up of only non-metals. Molecular formula: formula written using chemical symbols followed by subscripts to indicate the actual number of each type of atom in one molecule. element symbols to represent atoms subscripts follow symbols to show how many of those atoms are in the molecule no subscript means one atom C6H12O6 -- this molecule is composed of 6 C, 12 H and 6 O atoms. H2O2 -- One molecule of hydrogen peroxide contains 2 H & 2 O atoms Molecular formulas are always used for molecules. Empirical formulas: Formulas that show the lowest whole number ratio of elements in a compound. The empirical formula for glucose : CH2O The empirical formula for hydrogen peroxide is HO Empirical formulas are usually used for ionic compounds. Structural formulas: show the relative arrangement of atoms in a compound in a diagram in which chemical symbols indicate the element and dashes represent bonds. Used for molecules. Writing Formulas and Names for Molecular Compounds: Some molecules have special names. (memorize ) H2O is water NH3 is ammonia Rules for Binary Molecular Compounds This naming system is for compounds composed of two nonmetallic elements. 1. The first element keeps its name and gets a prefix if it has a subscript in the formula 2. The second element gets the –ide suffix (ending) and ALWAYS gets a prefix Details….details…. 1. Prefixes are used in molecular names to indicate the number of each type of atom in the compound. : Mono = 1 Hexa = 6 Di = 2 Hepta = 7 Tri =3 Octa = 8 Tetra = 4 Nona = 9 Penta = 5 Deca = 10 If the second element’s name begins with an “a” or “o”, drop the last vowel off the prefix name - Ex Carbon monoxide CO Carbon dioxide CO2 Diphosphorus pentoxide www.sciencegeek.net/Chemistry P2O5 IONS AND IONIC COMPOUNDS Metals react with non-metals to form ionic compounds. Metals give away their electrons to the non-metals, resulting in positively charged metal ions and negatively charged nonmetal ions. The oppositely charged ions are attracted to each other to form an ionic compound. Ions: Atoms that have lost or gained electrons by reacting with another atom, and , as a result , have an net (overall) charge. Cation = a positively charged ion caused by losing e Generally metals form cations Anion = a negatively charged ion caused by gaining e Generally non-metals form anions The charge of an ion is equal to the total # of positive (proton) charges plus the total # of negative (electron) charges Ion symbol : written using the chemical symbol followed by its charge written as a superscript Ex: Cl1- Mg2+ or atomic mass atomic # Symbol charge Ex: An unreacted oxygen atom has 8 protons (+8) 8 electrons (-8) Total charge 0 neutral, uncharged But… oxygen reacts with other atoms by gaining two electrons…so 8 protons (+8) 10 electrons (-10) Total charge -2 the oxygen ion has a -2 charge General chemical family trends for ion formation: Alkali metals lose 1 e- and form +1 cations Alkaline earth metals lose 2 e- and form +2 cations Group 3A lose 3 e- and form +3 cations Group 4A can lose or gain 4 e- to form either +4 or -4 charged ions Group 5A gain 3 e- and form -3 anions Group 6A gain 2 e- and form -2 anions Halogens gain 1 e- and form -1 anions Noble gases do not generally form ions – they are unreactive or inert. Polyatomic ion: group of generally non-metal atoms bonded together to form a molecule with an overall charge. Ex: SO42- sulfate - see reference sheet The majority of polyatomic ions are oxoanions – polyatomic anions made up of an element covalently bonded to oxygen (covalent bonds are the result of atoms sharing electrons) You will always have the reference sheet for these, but the following is the rationale behind their names… a. The “most important” form is named by shortening the element name and adding “ate” (_____ ate) Ex ClO3- Chorate b. Other forms are named relative to the ___ate form * if it has 1 less oxygen = _______ ite ClO2- = Chlorite * if it has 2 less oxygens = hypo____ ite ClO- = Hypochlorite * if it has 1 more oxygen = per ____ ate ClO4- = Perchlorate Ionic compounds are electrically neutral, meaning total positive charges on the cations = total charges the anions. This is the key to writing formula correctly. NaCl Na has a 1+ charge, and Cl has a 1- charge Na2S 2 x Na1+ = +2 total One S2- = -2 total ionic compounds form crystalline solids ( the ions are arranged in a repeating pattern called a crystal lattice) ionic compounds have no distinct molecule because of the lattice arrangement, so they are represented by an empirical formula: the simplest, whole number ratio of atoms or ions in a compound WRITING FORMULAS FOR IONIC COMPOUNDS 1. Write the chemical symbol for the cation first, followed by the symbol for the anion 2. Write subscripts after the chemical symbols to indicate the ratio of cations to anions for a neutral compound How to determine the ratio of cations to anions in an ionic compound….. Remember ionic compounds are neutral, meaning total charges must add up to zero! …this is because all electrons being lost MUST be gained by another atom, so total positive charges on cations MUST equal the total negative charge on anions involved. i. Look up charges of ions ii. Figure out how many you will need of each ion, so when you add up the total positive charges from the cations, it equals the total negative charges of the anions (and is then neutral!) Note – do not write a subscript of 1 Ex: Fe2O3 is a solid made up of a ration of 2 Fe atoms for every 3 O atoms. 3. Polyatomic ions must be put in parentheses if they need a subscript added. Ex Fe(NO3)2 indicates 1Fe2+ : 2 NO31- Ex: Write the correct formula for Sodium sulfide Magneisum hydroxide Calcium nitrate Barium oxide Ammonium nitrate Sodium fluoride Aluminum oxide http://www.chemfiles.com/flash/formulas.html NAMING IONIC COMPOUNDS A little background info is necessary to know: Cation names (three types of cations): a. Metals that form only one kind of cation are named with just their element name. Alkali metals +1 Alkaline earth +2 Boron Group +3 Some transition metal as such as Zn and Ag b. Transition Metals form more than type of cation, so their name includes the name of the element followed by their charge in roman numerals in parentheses. Ex Iron2+ and Iron3are named Iron (II) and Iron (III) c. Polyatomic cations – learn to recognize these from your reference sheet - and just write their name. (most polyatomic ions are anions) Look up name on reference sheet NH41+ Ammonium * Note if your ionic compound formula includes 3 or more elements, it includes a polyatomic ion. They are not always surrounded by parentheses. They only need to be surrounded by parentheses when the compound contains more than one of them. NaNO3 Sodium nitrate NH4NO3 Ammonium nitrate Fe(CN)2 Iron (II) cyanide Anion names: * there are two types of anions a. Monatomic non-metal anions are named by shortening element name and adding an “ide” ending F1O2Fluoride Oxide b. Polyatomic anions – look the names up on your reference sheet. Naming Ionic Compounds 1. If the compound you are naming has more than 2 elements in a formula, then you have a polyatomic ion. Ex: KOH is K & OH Potassium hydroxide NaNO3 is Na & NO3 Sodium nitrate NH4F is NH4 & F Ammonium flouride Practice – name the following: NaBr Ca(OH)2 (NH4)2SO4 AlCl3 KNO3 K2S Na2SO3 KCN 2. If the metal is a transition metal, you must put ( ) after the name and show the charge as Roman numerals Ex Lead ( IV ) You must figure out the charge on the metal by using your knowledge that ionic compounds are neutral!!! a. Figure out charge on transition metals by looking up the charge on the anion, then multiply that by the number of anions to get the total negative charge. b. Remember that total charges on cations must be equal (but opposite charge) to total neg charges so that the ionic compound is neutral. c. divide total positive charge by the number of cations present in the formula to find the charge on one of them. Write this value in the () as a Roman numeral. Ex: Write the name for CoCl2: Cobalt ( ) chloride start with the anion charge Cl has a charge of -1 …there are two Cl, so total neg charge is -2. Total neg charges MUST equal the total positive charges, so the charge on the one Cobalt cation must be +2 Cobalt (II) chloride Practice naming the following: SnO2 CuO Cu2O Co(NO3)2 NiSO4 Latin System for naming ionic substances that involve a metal that forms two types of ions The metal is named using a latin root followed by a suffix that indicates the higher or lower of the two possible charges. Latin roots include: Plumb = Lead Stann = Tin Ferr = Iron Cupr = Copper ous suffix – used for the lower of the two possible charges ic suffix – used for the higher of the two possible charges 2+ Ex: Fe = Ferrous and Fe3+ = Ferric Practice naming the following using the Latin system FeCl2 FeO Fe2O3 ACIDS: a molecular substance that produces hydrogen ions (H+) when dissolved in water. 1. Acids are covalently bonded 2. Acids dissociate (break apart) when dissolved in water to produce H+ and an anion Ex: HCl dissociates to form H+ and ClFormulas for acids can be written by adding one H for every negative charge on the anion. (Like ionic compounds they are electrically neutral) Ex: SO42- Acid formula: H2SO4 Sulfate will join with 2 H to produce Sulfuric acid Ex: Cl1Acid formula: HCl Chloride will join with 1 H to form Hydrochloric acid Acids are named depending on the anion they form when dissolved in water. Name of anion ________ ide Hypo____ite ________ite ________ate Per_____ate Name of the acid Hydro________ ic acid Hypo____ous acid _______ous acid _______ic acid Per_____ic acid So, how do you go about naming an acid: 1. Recognize the compound is an acid from the H in the beginning of the formula. 2. Look at the second half of the formula and identify (look up the name) of the anion. 3. Find the anion name in the list above, and name the acid accordingly. Ex: HCl HClO H3PO3 H3PO4 HClO4 chloride anion hypochlorite ion phosphite ion phosphate ion perchlorate ion Hydrochloric acid Hypochlorous acid Phosphorous acid Phosphoric acid Perchloric acid http://www.kentchemistry.com/links/naming/acids.htm Writing Formulas for Acids: - work backwards from the name and identify the anion, and the charge on the anion. - Formula will have 1 H for every neg charge on anion Ex: Hydrosulfuric acid the name tells you it is an acid containing a S2anion Formula will be H2S Naming pure molecular compounds that contain H: HCl in the pure state (not dissolved in water) is named using Hydrogen + anion name with ide ending Hydrogen chloride It is Hydrochloric acid only when it is dissolved in water. This is usually designated by an (aq) subscript. Acid Salt – formed when a metal is substituted for some of the Hydrogens in an acid. The remaining H become part of a polyatomic ion. Ex: Hydrosulfuric acid H2S Sodium replaces one of the H’s = NaHS Acid Salt Naming Rule #1 metal name + hydrogen (if one H) + polyatomic or anion name dihydrogen (if 2 H’s) NaHS = Sodium hydrogensulfide NaH2PO4 = Sodium dihydrogen phosphate Alternate naming of acid salts: when only one of the H is substituted for by a metal, the prefix “bi” is inserted between the metal name anion NaHSO4 = Sodium bisulfate SIMPLE ORGANIC COMPOUNDS 1. Compounds of C and H, often containing O, N and other non-metals 2. The carbon atoms in organic compounds are usually bonded to each other in the center of the molecule forming a “carbon backbone”. 3. Hydrocarbons are usually mostly non-polar (they may have polar ends or side chains), and generally have weak attractions among neighboring particles, causing low melting and boiling points. 4. FYI : Common organic substances used in a lab include benzene, toluene, hexane, propanol, methanol, ethanol. They are generally used as solvents. Functional group: specific groupings of atoms that give compounds similar properties. Hydrocarbons: compounds containing only C and H Alkanes: a sub category of hydrocarbon compounds in which each C is bonded to 4 other atoms , meaning 4 single bonds around the Carbons. - They have an “ane” ending to name 3. Alcohols: derivative of alkanes in which an H is replaced by an OH (hydroxide) group. (aka and OH group is attached to a carbon ) - Named by removing “e” from the alkane name and adding “ol”. Methanol: Ethanol 1-Propanol the “1” indicates the hyroxyl group is on the first or #1 carbon. 2-Propanol the “2” indicates the hyroxyl (OH) group is on the #2 carbon Isomers: compounds with the same molecular formula but different configurations of atoms, such as the 1-propanol and 2-propanol.