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  1. 1. ATOMS, MOLECULES AND IONSQBA Miguel A. Castro Ramírez
  2. 2. THE ATOMIC THEORY OF MATTERThe theory that atoms are the fundamental building blocksof matter reemerged in the early 19th century,championed by John Dalton.
  3. 3. THE ATOMIC THEORY OF MATTERDalton atomic theory:• Each element is composed of extremely small particles calledatoms• All atoms of a given element are identical to one another in massand other properties, but the atoms of one element are differentfrom the atoms of all other elements• Atoms of an element are not changed into atoms of a differentelement by chemical reactions; atoms are neither created nordestroyed in chemical reactions. E.g : electrolysis• Compounds are formed when atoms of more than one elementcombine; a given compound always has the same relative numberand kind of atoms.
  4. 4. THE ATOMIC THEORY OF MATTER Law of Chemical Combination:• Law of Constant Composition: In a given compound, the relativenumbers and kinds of atom are constant - The basis of Dalton’s postulates 4.• Law of the Conservation of Mass (Law of Conservation of Matter): - The total mass of materials present after a chemicalreaction is the same as the total mass present before reaction. - Atoms can neither be created nor destroyed in a chemicalrxn; thus the total mass remains unchanged Therefore, Total mass of reactants = Total mass of the products E.g. the total mass of element A and B is equivalent to thetotal mass of C and D in the reaction: A+B C+D - Basis for Dalton’s postulates 3.
  5. 5. THE ATOMIC THEORY OF MATTER• Law of Multiple proportions : If two elements A & Bcombine to form more than one compound, the masses of Bthat can combine with a given mass A are in the ratio of smallwhole numbers. Example:1) CO and CO2 For this example, the number of carbon atom isconstant (fixed mass of the both element). Therefore, theratio of oxygen atoms in CO and CO2 compound is always 1:22) H2O and H2O2 For this example, the number of hydrogen atom isconstant (fixed mass of both element). Therefore, the ratio ofoxygen atoms in H2O and H2O2 compound is always 1:2
  6. 6. THE DISCOVERY OF ATOMIC STRUCTUREBehavior of charged particles: Particles with the same chargerepel one another, whereas particles with unlike chargesattract one another. The electron • Streams of negatively charged particles were found to emanate from cathode tubes. • J. J. Thompson is credited with their discovery (1897). • Thompson measured the charge/mass ratio of the electron to be 1.76 x 108 coulombs/g.
  7. 7. THE DISCOVERY OF ATOMIC STRUCTUREMillikan Oil Drop Experiment • Once the charge/mass ratio of the electron was known (previous finding), determination of either the charge or the mass of an electron would yield the other. • Robert Millikan (University of Chicago) determined the charge on the electron (1.602 x 109 C) in 1909. • So, he calculated the mass of the electron: 1.602 x 109 C Electron mass = = 9.10 x 10- 28 g 1.76 x 108 C/g
  8. 8. THE DISCOVERY OF ATOMIC STRUCTURE Radioactivity• Radioactivity: Spontaneous emission of radiation by an atom• It was first observed by Henri Becquerel.• Marie and Pierre Curie also studied it• Three types of radiation were discovered by Ernest Rutherford:  α , β and γ • Each types differ in its response to an electric field. • Rutherford showed, both α and β rays consist of fast moving particles, which were called α and β particles. • α- positive charge, β – negative charge and γ - no charge
  9. 9. THE DISCOVERY OF ATOMIC STRUCTURE The Atom Circa 1900 • Thompson: Atom consisted of a uniform positive sphere of matter in which the electrons were embedded. • The prevailing theory was that of the “plum pudding” model. • It featured a positive sphere of matter with negative electrons embedded in it. • Ernest Rutherford proved this model is wrong
  10. 10. THE MODERN VIEW OF ATOMIC STRUCTURE• An atom comprises of protons, neutrons and electrons. (also called subatomic particles)• Charge of electron: -1.602 x 10-19 C• Charge of proton: + 1.602 x 10-19 C• 1.602 x 10-19 = Electric charge• For convenience, the charges of atomic and subatomic particles are expressed as multiple of this charge rather than in coulomb.• Thus, electron = 1-, proton= 1+• Neutron has no charge• Atoms: Equal number of electrons n protons, so atoms have no electrical charge• Protons and neutrons are packed into a tiny positively charged core of the atom known as the nucleus.
  11. 11. THE MODERN VIEW OF ATOMIC STRUCTURE• Atomic mass unit (AMU) : 1 amu = 1.66054 x 10 -24 g• Atoms are also small• Convenient way to express atomic dimensions is agstrom (Å) - 1 angstrom = 10-10 m
  12. 12. THE MODERN VIEW OF ATOMIC STRUCTURE Atomic Numbers, Mass Numbers and Isotopes• What makes an atom of one element different from another element? • Bcoz of the subatomic compositions • Atomic Number: Number of protons in the nucleus of an atom of any particular element • No of protons = no. of electrons • Mass number: Total number of protons + electrons in the atom X = Atomic symbol of the element A = mass number; Proton + Neutron Z = atomic number (the number of protons or electrons)THE SYMBOL OF THE ATOM OR ISOTOPE
  13. 13. THE MODERN VIEW OF ATOMIC STRUCTURE Isotopes • Atoms with identical atomic numbers but different mass numbers • Isotopes have different numbers of neutrons. 11 12 13 14 6 C 6 C 6 C 6 C
  14. 14. ATOMIC WEIGHTS The Atomic Mass Scale• 1 amu = 1.66054 x 10-24 and 1g= 6.02214 x 1023• The atomic mass unit is presently defined by assigninga mass of exactly 12 amu to an atom of the 12C isotopeof the carbon.• In these units, an 1H atom has a mass of 1.0078 amuand an 16O atom has a mass of 15.9949 amu.
  15. 15. ATOMIC WEIGHTS Average Atomic Masses • We can determine the average atomic mass of an element by using the masses of its various isotopes and their relatives abundance • Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations How to Calculate: The 3 naturally occurring isotopes of neon are; neon-20, 90.51%, 19.99244 amu; neon-21, 0.27%, 20.99395 amu; neon-22, 9.22%, 21.99138 amu. Calculate the weighted average atomic mass of neon. Solution: (0.9051 x 19.99244 amu) + (0.0027 x 20.99395 amu) + (0.00922 x 21.99138 amu) = 20.1794 amu• Average atomic mass (expressed in atomic mass units) is also known as itsatomic weight.
  16. 16. ATOMIC WEIGHTS Average Atomic MassesEXERCISE The 2 naturally occurring isotopes of copper are copper-63, mass 62.9298u, and copper-65, mass 64.9278u. What must be the percent natural abundances of the 2 isotopes if the atomic mass of copper listed in a table of atomic masses is 63.546u?
  17. 17. THE PERIODIC TABLE1. The modern P.T. is based on Mendeleev’s earlier work.2. Elements are organized and categorized according to their physical and chemical properties and in the order of increasing atomic number.3. Each vertical column is called a group. The main group is Group 1 to 8. Elements from the same group have similar physical and chemical properties.4. Periods are the horizontal rows of elements, labeled 1 to 7.
  19. 19. THE PERIODIC TABLE• For each element in the table, the atomic number and atomicsymbol are given.• May notice slightly difference and variations in PTs from 1book to another. These are simply matters of style, or theymight concern the particular information included. NOFUNDAMENTAL DIFFERENCES.• Elements that belong to the same group often exhibitsimilarities in physical and chemical properties. Example: ‘Coinage metal’- Cu, Ag and Au- belong to group 1B. Used throughout the world to make coins.
  20. 20. THE PERIODIC TABLE• Elements on the left and the middle side of the PT are metallicelements or metal• Share characteristic properties such as luster and highelectrical and heat conductivity.• All metals, with the exception of mercury (Hg), are solids atroom T.• The metals are separated from nonmetallic elements ornonmetals by a diagonal step-like line.• The elements that lie along the line that separates metalsfrom non-metals, such as (Sb), have properties that fallbetween those of metals and those of non-metals.• Referred as metalloids. - Tend to be semiconductors – Conduct electricity, but not nearly so well as metals.
  21. 21. THE PERIODIC TABLENames of Some Groups in the Periodic Table
  22. 22. MOLECULES AND MOLECULAR COMPOUNDS Molecules and Chemical Formulas• Many elements are found in nature in molecular form Examples: Oxygen – O2 Ozone- O3 • Compounds that are composed of molecules contain more than one type of atom and are called molecular compounds. Example: Water –H2O Hydrogen peroxide- H2O2 • Even though composed of same elements, both water and hydrogen peroxide have different properties. • Most molecular substances that we will encounter contains only non-metals.
  23. 23. MOLECULES AND MOLECULAR COMPOUNDS Molecular and Empirical Formula• Chemical formulas that indicate the actualnumbers and types of atoms in a moleculeare called molecular formula.• Empirical formulas give the lowest whole-number ratio of atoms of each element in acompound Example: Molecular Formula: H2O2 Empirical Formula: HO• Molecular formula provide moreinformation about molecules than doempirical formula
  24. 24. MOLECULES AND MOLECULAR COMPOUNDS Picturing Molecules• The structural formula of a substanceshows which atoms are attached to whichwithin the molecule.• Perspective drawings also show the three-dimensional array of atoms in a compound.• Examples: Water and hydrogen peroxide: O HO OH H H• Atoms are represented by their chemicalsymbols and lines are used to represent thebonds.
  25. 25. MOLECULES AND MOLECULAR COMPOUNDS• Structure formula does not depict the actual geometry of themolecule.• Perspective drawings can be written to show the three-dimensional array of atoms in a compound• Ball-and-stick model show atoms as spheres and bonds assticks. - The advantage of accurately representing the angles atwhich the atoms are attached to one another within themolecule.• Space-filling model: Molecule would look like if the atoms werescaled up in size. - These models show the relative sizes of the atoms, but the angles between atoms, which help define their molecular geometry, are often more difficult to see than in ball-and-stick models.
  26. 26. IONS AND IONIC COMPOUNDS Predicting Ionic Charges• The nucleus of an atom is unchanged by chemical processes, but someatoms can readily gain or lose electrons.• When atoms lose or gain electrons, they become ions. – Cations are (+) and are formed by elements on the left side of the periodic chart. e.g: Na+, Ca2+ – Anions are (-) and are formed by elements on the right side of the periodic chart. e.g: Cl- , Br-
  27. 27. IONS AND IONIC COMPOUNDS Ionic Compounds• Sodium Chloride (NaCl): Example of an ionic compound. - Contains both + and – charges • Ionic Compounds: Generally combinations of metals and nonmetals (NaCl). •In contrast, molecular compounds are generally composed of non-metals only (H2O)
  28. 28. IONS AND IONIC COMPOUNDS The Formation of Ionic Compounds• The atoms transfer electrons between them when they react e-• Na + Cl  Na+ + Cl-
  29. 29. IONS AND IONIC COMPOUNDS The Formation of Molecular Compound• Molecular compounds: Formed when non-metallic elementscombine• The attraction that hold atoms to each other in molecules,called chemical bonds, are electrical in nature.•They arise from sharing of electron between one atom toanother.
  30. 30. NAMING INORGANIC COMPOUNDS• System used in naming the substances: Chemical Nomenclature• Organic Compounds: Contain C, usually in combination withH,N,S or O. (CH4)• Inorganic Compounds: All others (Al2O3) Names and Formulas of Ionic Compounds1) Positive Ions (cation) • Cation formed from metal atoms have the same name as the metal. • Na+ sodium ion •If a metal can formed different cations, the + charge is indicated by a Roman numeral in parentheses following the name of the metal • Fe2+ iron(II) ion Fe3+ iron(III) ion
  32. 32. NAMING INORGANIC COMPOUNDS- Polyatomic or cations with variable - Monoatomic ions charges - Do not have variable charges Formula of the most common ions
  34. 34. NAMING INORGANIC COMPOUNDS• Anions derived by adding H+ to an oxyanion are namedby adding as a prefix the word hydrogen or dihydrogen,as appropriate: •CO32- carbonate ion • HCO3- hydrogen carbonate ion (bicarbonate) • H+ reduced the – charge of parent anion
  35. 35. NAMING INORGANIC COMPOUNDS3) Ionic Compound • Cation followed by anion • CaCl calcium chloride • Al(NO3)3 aluminium nitrate Exercisea. Give the names of the following ions: (i) O22- and (ii) SO42-b. Write down the formulas of the following ions: (i)ammonium and (ii) nitrate
  36. 36. NAMING INORGANIC COMPOUNDSPROBLEM: Give the systematic names or the formula or the formulas for the names of the following compounds: (a) Fe(ClO4)2 (b) sodium (c) Ba(OH)2 8H2O sulfite PLAN: Note that polyatomic ions have an overall charge so when writing a formula with more than one polyatomic unit, place the ion in a set of parentheses.SOLUTION: (a) ClO4- is perchlorate; iron must have a 2+ charge. This is iron(II) perchlorate. (b) The anion sulfite is SO32-; therefore you need 2 sodiums per sulfite. The formula is Na2SO3. (c) Hydroxide is OH- and barium is a 2+ ion. When water is included in the formula, we use the term “hydrate” and a prefix which indicates the number of waters. So it is barium hydroxide octahydrate.
  37. 37. NAMING INORGANIC COMPOUNDSNames and Formulas of Acids • If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- . – HCl: hydrochloric acid – HBr: hydrobromic acid – HI: hydroiodic acid
  38. 38. NAMING INORGANIC COMPOUNDS • If the anion in the acid ends in -ite, change the ending to -ous acid. – HClO: hypochlorous acid – HClO2: chlorous acid
  39. 39. NAMING INORGANIC COMPOUNDS • If the anion in the acid ends in -ate, change the ending to -ic acid. – HClO3: chloric acid – HClO4: perchloric acid
  40. 40. NAMING INORGANIC COMPOUNDSNames and Formulas of Binary Molecular Compounds1. The name of the less electro(-) - written 1st - exception for compound contain oxygen. - O always written last except when combined with fluorine 2. If both elements are in the group in PT, higher atomic no comes 1st. 3. Name of 2nd element (more electro(-))is given –ide ending 4. Greek prefixes are used to indicate the no of each element CO2: carbon dioxide CCl4: carbon tetrachloride5) If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one. N2O5: dinitrogen pentoxide
  41. 41. SOME SIMPLE ORGANIC COMPOUNDS• Organic chemistry is the study of carbon.• Organic chemistry has its own system of nomenclature.• The simplest hydrocarbons (compounds containing only carbon and hydrogen) are alkanes.• The first part of the names above correspond to the number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).
  42. 42. SOME SIMPLE ORGANIC COMPOUNDS• When a hydrogen in an alkane is replaced with something else (a functional group, like -OH in the compounds above), the name is derived from the name of the alkane.• The ending denotes the type of compound. – An alcohol ends in -ol.
  43. 43. •Alkene•alkuna