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Chapter 2

Chapter 2. Atoms, Molecules, and Ions. Atoms and Atomic Structure. Dalton’s Atomic Theory - 1808 - Elements are composed of small, nondivisible particles called atoms -Atoms of an element have identical properties and differ from those of other elements

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Chapter 2

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  1. Chapter 2 Atoms, Molecules, and Ions

  2. Atoms and Atomic Structure • Dalton’s Atomic Theory - 1808 • -Elements are composed of small, nondivisible particles called atoms • -Atoms of an element have identical properties and differ from those of other elements • -Atoms cannot be created, destroyed, or transformed into other atoms • -Compounds are formed when atoms of different elements combine in whole-number ratios • -Atom ratios are constant in a given compound • -Chemical reactions rearrange and recombine atoms but do not destroy them

  3. Structure of the Atom • -atom is mostly empty space • -consists of a very small, dense center called the nucleus • -nearly all of the atom’s mass is in the nucleus • -the nuclear diameter is 1/10,000 to 1/100,000 times less than the atom’s radius

  4. Structure of the Atom • -Sir John Joseph Thompson and Ernest Rutherford established a model of the atom still in use today • -Three fundamental particles make-up atoms:

  5. Elements • substances that cannot be decomposed into simpler substances via chemical reactions • Elemental symbols-abbreviation representing each element on periodic table • -First letter capital, second letter lower case • ie: C, Ca, Co • CO is not an element (it’s a compound) because there are two capital letters

  6. The Periodic Table • 1869 - Mendeleev & Meyer • Discovered the periodic law • -Organized based on related chemical reactivities, physical properties, other behaviors and trends • -The properties of the elements are periodic functions of their atomic numbers (not atomic masses) Law of Chemical Periodicity

  7. The Periodic Table • Groups or families • Vertical group of elements on periodic table • Similar chemical and physical properties • Period • Horizontal group of elements on periodic table • Transitions from metals to nonmetals • Three regions • Metals, nonmetals and metalloids

  8. Period • Horizontal group of elements on the periodic table • -Transition from metals to nonmetals

  9. Groups (families) • Vertical group of elements on the periodic table • -similar chemical and physical properties

  10. Metalloids separate metals from nonmetals Metals are to the left of stair step -Approximately 80% of the elements Nonmetals are to the right of stair step -Approximately 20% of the elements Elements box on the stair step have properties between metals and nonmetals

  11. The Periodic Table • Chemical properties of metals • -Outer (valence) shells contain few electrons • -Form cations by losing electrons • -Form ionic compounds with nonmetals • -Solid state characterized by metallic bonding -Conductors of electricity and heat -Malleable: can be hammered -Ductile: drawn into wire -Typically solids (except mercury)

  12. The Periodic Table • Chemical properties of nonmetals • -Outer shells contain four or more electrons • -Form anions by gaining electrons (- charge) • -Form ionic compounds with metals and covalent compounds with other nonmetals • -Form covalently bound molecules; noble gases are monatomic and have full electron shells • -Insulators meaning poor conductors • -Typically gasses or solids (Br2 liquid I2 solid)

  13. Periodic Table The Periodic Table: • Periodic trends Atomic number increase Size decreases Atomic number increase Size increases Form compounds with similar formulas

  14. Metals • Group IA metals: • Alkali metals- Li, Na, K, Rb, Cs, Fr • -solids, reactive, and never found in nature as free elements • -chemical formula from top-to-bottom are the same • Group IIA metals: • Alkaline earth metals- Be, Mg, Ca, Sr, Ba, Ra • -solids, reactive, and mostly found in nature as compounds • -chemical formula from top-to-bottom are the same

  15. Metals Group IIIA: B, Al, Ga, In, Tl • -Aluminum (Al) most abundant metal in the earth’s crust • -Boron (B) is the only nonmetal • -forms compounds of analogous chemical formulas Group IVA: C, Si, Ge, Sn, Pb -most variation in their properties -Carbon (C) is the “element of life” CO2 carbonate-limestone, coral, shells fossil fuels-coal, petroleum, natural gas -Silicon (Si) is found as gemstones, glass, and sand -Lead (Pb) was used as water pipes, paint, and in gasoline -form compounds of analogous chemical formulas

  16. Allotropes • elements (nonmetals) that exist in several different and distinct forms each having its own properties

  17. Group VA: N, P, As, Sb, Bi • Nitrogen (N) • -makes up ~75% of the earth’s atmosphere • -used to make fertilizer (NH3) • -found in biological systems in proteins and DNA • Phosphorous (P) • -has several allotropes most common are white and red phosphorus • -white ignites spontaneously in air and used to make phosphoric acid • -red phosphorus is used in striking strips of matchbooks • -form similar types of chemical compounds

  18. Group VIA: nonmetals • Chalcogens • O, S, Se, Te -Oxygen (O) is the power source of life on earth by combining with other substances, and has allotropes -Sulfur (S) (and even Selenium, Se) is fowl smelling, and S appears as allotropes -these elements are considered poisonous, but essential for human diets -Some variations is chemistries, but form analogous formulas

  19. Group VIIA nonmetals • halogens • F, Cl, Br, I -All exist in the form of diatomic molecules -At room temperature, Fluorine (F) and Chlorine (Cl) are gases where Bromine (Br) is a liquid and Iodine (I) is a solid -Some of the most reactive of all elements they react with metals and nonmetals to form compounds

  20. Group VIIIA nonmetals • noble, inert or rare gases • He, Ne, Ar, Kr, Xe, Rn -least reactive elements and for a long time considered unreactive -all are gases and none are abundant on earth -He is the second most abundant element in the universe

  21. Transition Metals • Most are found as compounds in nature • -Ag, Au, Pt are less reactive and can be found as pure substances • -These elements are commercially useful as building materials, in paints, catalytic converters, coins, batteries, and fireworks • -Play important roles in biological processes • -Bottom two rows of the periodic table are used in television picture tubes, atomic fuel, smoke detectors

  22. Atomic Number (Z) the number of protons in the nucleus • -numbered consecutively on the periodic table • atomic number determines the element: • -elements differ from one another by the number of protons in the nucleus • -the number of electrons in a neutral atom (no charge) is equal to the atomic number • -negatively charged species have more electrons • -positively charged species have less electrons

  23. Mass Number (A) • sum of the number of protons and neutrons • Z = atomic number (number of protons) • N = number of neutrons • A = Z + N • One common symbolism used to show mass and proton numbers is:

  24. Mass Number (A) • How many protons and neutrons are in the following?

  25. Mass Number (A) Give the number of protons, neutrons and electrons and the correct element symbol:

  26. Isotopes • Atoms of the same element but with different numbers of neutrons (or same atomic number, but different mass numbers) Isotopes have different masses and mass number(A) values but are the same element • Example: hydrogen isotopes or protium is the most common hydrogen isotope • -one proton and no neutrons or deuterium (D) is the second most abundant hydrogen isotope • -one proton and one neutron or tritium (T) is a radioactive hydrogen isotope • -one proton and two neutrons

  27. Atomic Weights • How do we know what the values of these atomic weights are?

  28. Atomic Weight • the weighted average of the masses of the elements stable isotopes • Example: Naturally occurring Cu consists of 2 isotopes. It is 69.1% 63Cu with a mass of 62.9 amu, and 30.9% 65Cu, which has a mass of 64.9 amu. Calculate the atomic weight of Cu to one decimal place.

  29. Atomic Mass • weighted average of the masses of an elements stable isotopes as listed on the periodic table • For example: hydrogen (H) = 1.008 amu • calcium (Ca) = 40.078 amu

  30. Molecules smallest unit of a pure substance that can be divided and still retain the composition and chemical properties of the substance Examples of molecules: • H2 • O2 • S8 • H2O • CH4 • C2H6O Molecular formulas:describe the composition of substances, but provide no structural information

  31. Classes of Substances: monatomic elements He, Au, Na diatomic elements (binary molecules) O2, H2, Cl2, F2, I2, N2,Br2 complex elements O3, S4, P8 Compounds (molecules) H2O, C12H22O11

  32. Molecular (Chemical) Formulas Compound Contains HCl H2O NH3 C3H8

  33. Chemical bonds Attractive forces that hold atoms together in compounds Chemical bond types: • Ionic bonding: resulting from electrostatic attractions between ions -formed by the transfer of one or more electrons from one atom to another -attraction of cations for anions typically form solids -most often formed by interactions between metals and nonmetals • Covalent bonding: results from sharing one or more electron pairs between two atoms - typically formed by interactions between nonmetals and nonmetals

  34. Ionic Compounds An ion is an atom or a group of atoms possessing a net electrical charge positive (+) ions or cations These atoms have lost 1 or more electrons Metals lose electrons to form cations (monoatomic cations) Consider Group IA and IIA metals negative (-) ions or anions These atoms have gained 1 or more electrons. Nonmetals gain electrons to form anions (monoatomic anions) Consider the nonmetals-groups IVA through VIIA

  35. Ionic Compounds • Coulomb’s Law: dictates the strength of ionic bonds – it is an inverse square law =

  36. Ionic Compounds Cations: Na+, Ca2+, Al3+ (monoatomic cations) NH4+ (polyatomic cation) Anions: F-, O2-, N3- (monoatomic anions) SO42-, PO43-, HCO3- (polyatomic anions) Transition metals forming cations No pattern exists for determining the charge Many metals form several different ions (charge states) H and noble gases are special cases

  37. Ionic Formulas Formulas of ionic compounds are determined by the charges of the ions -total charge of the cations must equal the total charge of the anions -The compound must be neutral NaCl sodium chloride (Na1+ & Cl1-) KOH potassium hydroxide (K1+ & OH1-) CaSO4 calcium sulfate (Ca2+ & SO42-) Al(OH)3 aluminum hydroxide (Al3+ & OH1-)

  38. Writing binary chemical formulas Charge becomes subscript. Then give the subscript as lowest common denominator

  39. Naming Ionic Compounds metal cation and a nonmetal anion • Name the cation • Name the anion, nonmetal stem with –ide ending LiBr lithium bromide magnesium chloride Li2S lithium sulfide Al2O3You do it!

  40. Naming Ionic Compounds LiBr lithium bromide MgCl2 magnesium chloride Li2S lithium sulfide Al2O3 aluminum oxide Na3P Mg3N2 Notice that binary ionic compounds with metals having one oxidation state They do not use prefixes or Roman numerals!!!

  41. Naming Cations Monoatomic cations use the name of the metal plus the word “cation” Examples: Al3+, Li+, Ca2+ Transition metals that have multiple charges: Two methods are used 1. Older method: -add suffix “ic” to element’s Latin name for higher oxidation state (higher positive charge) -add suffix “ous” to element’s Latin name for lower oxidation state (lower positive charge) 2. Modern method: use Roman numerals in parentheses to indicate metal’s oxidation state

  42. Naming Ionic Compounds ionic compounds containing metals with more than one oxidation state (charge) memorize them on your handout Metals exhibiting multiple oxidation states are: • most of the transition metals • metals in groups IIIA (except Al), IVA, & VA on the periodic table

  43. Naming Ionic Compounds CompoundOld SystemModern System FeBr2 ferrous bromide iron(II) bromide ferric bromide iron(III) bromide SnO tin(II) oxide SnO2 tin(IV) oxide CoCl2 cobaltous chloride CoCl3 cobaltic chloride plumbous sulfide lead(II) sulfide plumbic sulfide lead(IV) sulfide

  44. Naming Ionic Compounds There are polyatomic ions that form binary ionic compounds • OH- hydroxide • CN- cyanide • NH4+ ammonium KOH potassium hydroxide barium hydroxide Al(OH)3 aluminum hydroxide Fe(OH)2You do it!

  45. Naming Ionic Compounds KOH potassium hydroxide Ba(OH)2 barium hydroxide Al(OH)3 aluminum hydroxide Fe(OH)2 iron (II) hydroxide iron (III) hydroxide Ba(CN)2 ammonium sulfide NH4CN

  46. Naming Anions Monoatomic anions use the name of the nonmetal plus the word “-ide” at the end Examples: F-, Cl-, S2- Polyatomic anions Some guidelines for oxoanions (contain Oxygen) • If only 2 similar formula type anions exist, the one containing the greater number of oxygen atoms have an “–ate” ending, and the smaller number of oxygen atoms have an “-ite” ending NO3- NO2- SO42- SO32- • If more than two exist, the one with the largest number of oxygen atoms have a prefix “per-” and an “-ate” ending, and the smallest number of oxygen atoms have a prefix “hypo-” and an “-ite” ending ClO4- ClO3- ClO2- ClO- • Oxoanions containing H are named with the word hydrogen in front, if more than one H is contained in the oxoanion, then prefixes are used to indicate the number of hydrogen atoms HPO42- H2PO4-

  47. Naming Ionic CompoundsOn Your Own NaNO2 sodium nitrite sodium nitrate Na2SO3 sodium sulfate Na3PO4 MgSO4 Ca(NO3)2 BaCO3 potassium phosphate

  48. Ionic Formulas What is the name of K2SO3? Potassium sulfite What is charge on sulfite ion? -2 What is the formula of ammonium sulfide? (NH4)2S What is charge on ammonium ion? +1 What is the formula of aluminum sulfate? Al2(SO4)3 What is charge on both ions? +3 -2

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