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Back Chapter 2: Atoms Molecules and Ions The investigation and understanding of the atom is what chemistry is all about! Anyone who says that they can contemplate quantum mechanics without becoming dizzy has not understood the concept in the least. -Niels Bohr ©Bires, 2002 Slide 1 Back Atomic History • The Greek philosopher Democritus (400BC) coined the term atomon which means “that which cannot be divided.” • Idea of an indivisible thing that made up all matter was refined by colorblind chemist John Dalton in 1803. Among his interests, Dalton was very interested in a scientific explanation for his colorblindness and the behavior of gases. • In his A New System of Chemical Philosophy, Dalton published five principles of matter. ©Bires, 2002 Slide 2 Dalton’s Top Five Back • All matter is made of indestructible and indivisible atoms. (atoms are hard, unbreakable, and the smallest thing there is) • Atoms of a given element have identical physical and chemical properties. (all atoms of X will behave the same anywhere) These two are usually combined • Different atoms have different properties. (X behaves differently than Y) • Atoms combine in whole-number ratios to form compounds. (two H’s and one O = Water (H2O) • Atoms cannot be divided, created or destroyed, (just rearranged) in chemical reactions. ©Bires, 2002 Slide 3 Back The Laws: • Constant Composition: Ratios of atoms in a compound is constant for that compound. • Conservation of Mass: Mass is not created or destroyed in a chemical reaction. • Multiple Proportions: Since atoms bond in small, whole number ratios to form compounds, their masses are small whole number ratios. ©Bires, 2002 Slide 4 Back The CRT (cathode ray tube) Electrostatics! • A new invention, the cathode ray tube, (c1850s) suggested the presence of charges – areas of positive and negative…charge. • This suggested that atoms must be divisible, and Dalton’s theory had to be modified. • In 1897, English Physicist J. J. Thompson proposed that the atom is a sphere of positive charge with small areas of negative charge. • This theory became known as the “plum pudding” model after an English “dessert” of purple bread and raisins. ©Bires, 2002 Slide 5 Millikan’s Oil Back • In 1897, Thompson used electrostatics experiments to determine the charge-to-mass 1 C = 1 Amp•1 second ratio of 1.76 x 108 C/g. • Millikan’s classic oil-drop experiment allowed the charge of a single electron to be determined: 1.60 x 10-19 C. • Using these two numbers, we can calculate the 1.60 x1019 C mass of an electron: m 9.10 x1028 g electron • • ©Bires, 2002 1.76 x108 C / g The mass of an electron is about 1/2000 of the mass of a proton! Slide 6 Back Ernest Rutherford • While studying radioactive elements, New Zealander Physicist Ernest Rutherford found that radioactive alpha particles deflected when fired at a very thin gold foil. • This was known as the gold foil experiment, and it suggested that the atom was not a hard sphere as thought, but was mostly space, with a small concentration of mass. • This concentration of mass became known as the nucleus. • Link to experiment… ©Bires, 2002 Slide 7 Back The Bohr Model • A Danish physicist, Niels Bohr, (a student of Rutherford) rebuilt the model of the atom placing the electrons in energy levels (the “solar system” model). • Bohr was one of the founders of quantum physics – a discipline that states that energy can be given off in small packets or quanta of specific size. • Energy levels closer to the nucleus were lower in energy than those farther away. • When a specific amount of energy was added to an atom, an electron could jump into a higher energy level. ©Bires, 2002 Slide 8 Back Adding the Neutrons • Although theorized by Rutherford, British physicist, James Chadwick, proved the existence of massive, neutral particles. • These particles came to be called neutrons… • and their discovery in 1932 opened the door for more in depth investigations into radioactive materials and gave the WWII Allies the ability to enrich and purify fissionable uranium, a necessity in the production of nuclear weapons. ©Bires, 2002 Slide 9 Back The Modern Model Democritus and Dalton’s atom electron Thompson’s electrons neutron Rutherford’s space and nucleus proton Bohr’s energy levels (not to scale) ©Bires, 2002 Chadwick’s neutrons Most atoms ~ 1-5 Å (Angstroms) = (1x10-10m) Slide 10 Back Elements 8 O • We currently know of about 110 elements, 92 of which are naturally occurring. OXYGEN • We illustrate an element with an 15.9994 atomic symbol. • We add the number of protons, the atomic number… • And the atomic mass, the total mass of the protons plus the neutrons. Notice that the atomic mass is not a round number, ©Bires, 2002 even though protons and neutrons each have a mass of 1. This is due to natural abundance. Slide 11 Back Natural Abundance - Isotopes • There is not just one type of each atom, there are several. When a nuclide has more or less neutrons than another nuclide of the same element, we call them isotopes. • For instance, the element carbon has 6 protons, but it could have 5, 6, 7, or 8 neutrons, to form 12C Carbon-11, Carbon-12, Carbon-13, and 13C Carbon-14. Each has a different mass. 14C • In nature, there is a mix of different natural isotopes. We use this mix to calculate average atomic mass… ©Bires, 2002 Slide 12 Back Calculating Average Atomic Mass • To find average atomic mass, we multiply the relative abundance of an isotope by the mass of the isotope. We then add each of the products for each isotope. • Example: The isotopes of element Bob are found below: • Bob-18, 25% 0.25x18 0.60x19 0.15x20 • Bob-19, 60% • Bob-20, 15% • What is the average atomic mass of naturally occurring Bob? 1amu = 1.66x10-24 grams 18.90amu ©Bires, 2002 Slide 13 Back Isotopes, Ions, and Allotropes (Oh my) • Isotopes are atoms of the same element with different numbers of neutrons. • Ions are atoms of the same element with different numbers of electrons. • Ions are easy to create; adding or removing electrons can be done with electric current. • Allotropes are forms of the same element, bonded in different structures. • Diamond and pencil graphite are allotropes. They are both pure carbon, but in different structures. ©Bires, 2002 Slide 14 Back Isotopes and allotropes ©Bires, 2002 Slide 15 Balancing Chemical Equations Back General Rules: Compounds and Diatomics Never change subscripts (H2SO4) Balance groups first (H2SO4) Balance water, oxygen, and hydrogen last Diatomic Atoms Hydrogen = H2 Fluorine = F2 Bromine = Br2 Nitrogen = N2 Chlorine = Cl2 Iodine = I2 • Recall that compounds are collections of atoms. • Compounds can be made by: Ionic Bonding: Giving and taking electrons (Metal-nonmetal bond) Covalent Bonding: (Molecules) Sharing electrons (nonmetal-nonmetal bond) • There are gases that are bound to themselves as elements: • Hydrogen = H2 Nitrogen = N2 Oxygen = O2 • Fluorine = F2 Chlorine = Cl2 Bromine = Br2 • Iodine = I Oxygen = O2 ©Bires, 2002 Slide 16 Back Formulas • We can represent molecules with molecular and empirical formulas. • Molecular Formula = shows actual number of atoms in a molecule. C6 H12O6 H 2O2 • Empirical Formula = shows ratio of atoms in a molecule. CH2O HO • Compounds are represented with empirical formulas. ©Bires, 2002 Slide 17 Back Ions • Recall that an ion is an atom that has gained or lost one or more electrons. • Recall that the octet rule predicts that Cl atoms try to achieve zero or eight electrons in their outer (valence) shell. Cl- Na Na+ • When an atom bonds with another atom, it seeks to gain electrons or lose them. For instance: Positive ions are • Cl has 7 and will gain one electron called “cations” • Na has 1 and will lose one electron Negative ions are called “anions” ©Bires, 2002 Slide 18 Back Ionic attraction • Just like opposite poles of a magnet attract, oppositely charged ions attract. • Here, our positively charged sodium ion and negatively charged chlorine ion are attracted. • They remain ions, but stick together in a lattice (3D grid pattern) as other ions join them. ©Bires, 2002 Cl Na+ Cl- Na+ Na+ Cl- + Na Cl Slide 19 Back Formula Units • In our example, the Na+ and Clcombined to form NaCl. • Here, NaCl is one formula unit – Cl- Na+ the smallest unit that has the correct formula for our compound. • NaCl is also our compound’s empirical formula – the smallest ratio of atoms that make up our compound. • Although there may be several Na+ ions and Cl- ions in this lattice, the empirical unit is still just NaCl. ©Bires, 2002 Slide 20 • • • • Inorganic Nomenclature General Rules: Cations are listed first Anions/polyatomics last Binary compounds get the “–ide” ending ex: (sodium chloride) • Acids: Monatomic anion : “hydro-” + name + ic “acid” (ex: hydrochloric acid) Polyatomic anion “-ite” = name + “ous acid” ex: (sulfite = sulfurous acid) • Polyatomic anion “-ate” = name + “-ic acid” (ex: sulfate = sulfuric acid) • Compounds: Use Greek prefixes to denote number of atoms of each ©Bires, 2002 Back Naming Compounds Flow Chart Does the Formula begin with H NO YES Does it begin with a metal which has more than one oxidation number? (e.g.; Fe, Ni, Cu, Sn, Hg) NO YES It is an Acid (must be aqueous) Does the acid contain a polyatomic ion? Name the first element followed by its oxidation number (Roman numeral) YES NO Ending of polyatomic ion Does the formula contain a polyatomic ion? NO -ite YES Are both elements nonmetals? NO -ate Name the polyatomic ion, replacing the –ate ending with –ic. Add the word acid. YES Name the first element, Then the second element With the –ide ending. Name the first element using the proper prefix. (never mono-). Name the second element with the proper prefix (including mono- and –ide ending. 1 mono2 di3 tr4 tetra5 penta6 hexa7 hepta8 octa9 nano10 decaName the first element, then the polyatomic ion. If two elements are present, name both, then the polyatomic ion. (e.g.; NaHCO3 Is sodium hydrogen carbonate) Name the polyatomic ion, replacing the –ite ending with –ous. Add the word acid. Write the prefix hydro, then the name of the second element with –ic ending. Add the word acid. PRACTICE! PRACTICE! PRACTICE! PRACTICE! Slide 21 Back Molecular and Empirical Formulas • The formulas for compounds that exist as molecules are called molecular formulas. Ex: Hydrogen peroxide, H2O2, consists of two H atoms and two O atoms. • The atomic ratio of hydrogen peroxide is one to one. • The simplest formula that would indicate this ratio would be HO. This simple formula is an empirical formula. ©Bires, 2002 Slide 22 Back Molecular and Empirical Formulas • Formulas for ionic compounds are almost always empirical. Ex: NaBr is empirical because it takes one sodium atom to combine with one bromine atom to obtain a neutral charge. ©Bires, 2002 Slide 23 Back Molecular and Empirical Formulas • C6H12O6 • CH2O is molecular formula for glucose is empirical because we could factor out a six from each subscript. • CH3COOH is molecular for acetic acid • CH2O is empirical (factor out a 2). * Note: glucose and acetic acid are very different molecules yet they have the same empirical formula. ©Bires, 2002 Slide 24 Back Analysis by Combustion • Formulas for organic substances can be determined by combustion analysis. Carbon is recovered as CO2, hydrogen as water. Oxygen can be determined by mass difference if only C, H and O are present. 2B discussed more later… ©Bires, 2002 Slide 25 Back Steps for Determining an Empirical Formula: 1. Start with the number of grams of each element, given in the problem. *If percentages are given, assume that the total mass is 100 grams so that the mass of each element = the percent given. 2. Convert the mass of each element to moles using the molar mass from the periodic table. 3. Divide each mole value by the smallest number of moles calculated. 4. Round to the nearest whole number. This is the mole ratio of the elements and is represented by subscripts in the empirical formula. ©Bires, 2002 Slide 26 What if you cannot round off to whole numbers? Back – If the number is too far to round (#.10 to #.90), then multiply each solution by the same factor to get the lowest whole number multiple. • e.g. If one solution is 1.5, then multiply each solution in the problem by 2 to get 3. • e.g. If one solution is 1.25, then multiply each solution in the problem by 4 to get 5. ©Bires, 2002 Slide 27 Back Molecular Formulas • Calculations that are based on % composition always lead to the empirical formula of a compound (because it is already in the most reduced ratio). • Once the empirical formula is found, the molecular formula for a compound can be determined if the molar mass of the compound is known. • The molecular formula differs from the empirical formula in that it is a whole number multiple of the empirical formula. The empirical formula is the most reduced ratio of the molecular formula. ©Bires, 2002 Slide 28 Back Calculating Molecular Formulas STEPS: 1. Calculate the mass of the empirical formula and divide the molar mass of the compound by the mass of the empirical formula to find the ratio between the molecular formula and the empirical formula. 2. Multiply all the atoms (subscripts) by this ratio to find the molecular formula. ©Bires, 2002 Slide 29 Back Formula for a hydrate • Finding the empirical formula for a hydrate involves finding the mass difference of the hydrate and anhydrous salt by driving off the excess water: • DEMO: http://www.chem.iastate.edu/group/Greenbowe/ sections/projectfolder/flashfiles/stoichiometry/e mpirical.html ©Bires, 2002 Slide 30 Back THE END ©Bires, 2002 Slide 31 Back The Periodic Table • The periodic table is a collection of all the known elements into a model that groups elements with similar properties. • Vertical columns represent Groups of elements with similar properties. • Horizontal Periods represent elements with similar electron configurations. ©Bires, 2002 Slide 32 Back The Periodic Table • The periodic table was created by Russian chemist Dmitri Mendeleev, who ordered the known elements according to atomic mass. • The modern periodic table is ordered according to atomic number. • The elements in the periodic table are arranged with Periodic Law which states that the chemical and physical properties of elements are periodic functions of their atomic numbers. • Periodic Law shows certain trends in the properties of elements (more in C6 and C7) ©Bires, 2002 Slide 33