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Chapter 2: Atoms, Ions and Compounds Problems: 2.1-2.80, 2.99-2.101, 2.104-2.105, 2.109-2.112.

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Presentation on theme: "Chapter 2: Atoms, Ions and Compounds Problems: 2.1-2.80, 2.99-2.101, 2.104-2.105, 2.109-2.112."— Presentation transcript:

1 Chapter 2: Atoms, Ions and Compounds Problems: , , ,

2 Chapter 2 – Atoms, Ions and Compounds 2.1 The Rutherford Model of Atomic Structure (or Nuclear Model) Joseph John (J. J.) Thomson (1897) carried out experiments with cathode rays. –His research group determined the charge-to- mass ratio of the particles in the rays. –They determined the particles were composed of tiny, negatively charged subatomic particles  electrons (e – )

3 Atomic Structure Eugen Goldstein (late 1880s) –Discovered canal (or anode) rays which were composed of positively charged subatomic particles  protons (p + ) –And decades later, James Chadwick won the Nobel Prize winner for his discovery (1935)  neutron (n) = neutral subatomic particle

4 Plum-Pudding Model of the Atom Thomson proposed that the atom was a uniform sphere of positively charged matter in which electrons were embedded – electrons are like raisins in a pudding of protons

5 The Nuclear Atom: Protons and the Nucleus Ernest Rutherford was a scientist who did many pioneering experiments in radioactivity He had members of his research group test Thomson’s Plum-Pudding Model using radioactive alpha (α) particles, basically helium atoms with a +2 charge and much bigger than an electron.

6 Rutherford's Alpha-Scattering Experiment Alpha (α) particles shot at a thin gold foil that’s only a few atoms thick –A circular detector is set up around the foil to determine what happens to the α particles. –If Plum-pudding Model was correct, the α particles should go through the foil like bullets through tissue paper.

7 Experimental results: Most of the a particles went straight through, but some were deflected or even bounced back! Rutherford's Alpha-Scattering Experiment

8 Consider what the scientists expected to observe given each model. Rutherford's Alpha-Scattering Experiment

9 Rutherford’s interpretation of the results Most alpha (α) particles pass through foil –Atom is mostly empty space with electrons moving around the space Some α are deflected or bounce back –Atom must also contain a dense region, and particles hitting this region are deflected or bounce back towards source. Dense region = atomic nucleus (contains atom’s protons and neutrons) That’s why this is called the NUCLEAR Model of the Atom Rutherford's Alpha-Scattering Experiment

10 Rutherford also estimated the size of the atom and its nucleus: An atom is 100,000 times (10 5 or 5 orders of magnitude) bigger than its nucleus. Rutherford's Alpha-Scattering Experiment

11 Example 1 An atom is 100,000 times (10 5 or 5 orders of magnitude) bigger than its nucleus.  If a nucleus = size of a small marble (~1 cm in diameter), indicate the length in meters then identify a common item that corresponds to that size for the following: a. 10 times bigger = _____________ m = ________________________ b. 100 times bigger = _____________ m = _______________________ c times bigger = _____________ m = ______________________ d. 10,000 times bigger = ____________ m = _____________________ e. 100,000 times bigger = ____________ m = ____________________

12 Properties of Protons, Neutrons, and Electrons Subatomic Particle ChargeLocationMass (amu) Proton+1inside nucleus Neutron0inside nucleus Electronoutside nucleus

13 2.2 Isotopes Each element always has the same number of protons, but the number of neutrons may vary. Atoms with different numbers of neutrons are called isotopes. –Carbon exists as carbon-12, carbon-13, and carbon-14 where each carbon atom has 6 protons but 6, 7, or 8 neutrons respectively Isotopes are identified with an element name followed by the mass number –Examples: uranium-235 (U-235), carbon-12 (C- 12), cobalt-60 (Co-60), etc.

14 Atomic Notation shorthand for keeping track of number of protons and neutrons in the nucleus –atomic number (Z): whole number of p + = number of e – in neutral atom –mass number (A): whole number sum of protons and neutrons in an atom Note: electrons contribute almost no mass to an atom E A Z element symbol mass number atomic number

15 Example Problem Isotope Mass number # of protons # of neutrons # of electrons strontium Mo zinc Ba P 3- Complete the following table:

16 2.4 The Periodic Table of the Elements Know which elements are metals, semimetals, nonmetals using the Periodic Table.

17 Properties of Metalloids (or Semimetals) Have properties of metals and nonmetals For example, silicon is shiny like a metal and acts as a semiconductor. Properties of Metals conduct heat & electricity malleable: can be flattened into thin sheets ductile: can be stretched into a wire Examples: aluminum, copper, gold Properties of Non-Metals dull appearance brittle non-conductor Examples: carbon (graphite in pencils), sulfur

18 SOLIDS, LIQUIDS, AND GASES: KNOW the physical state of each element at 25  C! At standard state conditions (25  C and 1 atm): –Only mercury (Hg) and bromine (Br) are liquid –H, N, O, F, Cl, and all Noble gases (group VIIIA) are gases –All other elements are solids

19 2.5 Trends in Compound Formation Chemical bond: holds atoms or ions together in a compound COVALENT BOND and MOLECULES Covalent bond: sharing of a pair of electrons by 2 nonmetal atoms Two or more covalently bonded atoms form a molecule Molecule:basic unit of a compound of covalently bonded atoms –Consider the HCl, H 2 O, NH 3, and CH 4 molecules below –Note how the chemical formula gives the actual number of each atom present in the following compounds:

20 Diatomic Molecules Recognize these elements that exist as diatomic molecules (X 2 ): H 2 N 2 O 2 F 2 Cl 2 Br 2 I 2 Consider these the “diatomic seven” since there are seven of them, and six of them form a 7 on the periodic table.

21 Ions When atoms lose or gain electrons, they form charged particles called ions. –Metals lose electrons  positively charged ions = cations –Nonmetals gain electrons  negatively charged ions = anions Main-group elements generally form ions—i.e. gain or lose electrons—to get the same number of electrons as a Noble gas. –Ions formed by main-group elements are usually isoelectronic with—i.e., have the same number of electrons as—one of the noble gases!

22 Ions Formed by Main-Group Elements Group IA elements  +1 charge: Li +, Na +, K +, etc. (“+” = “+1”) Group IIA elements  +2 charge: Mg 2+, Ca 2+, Sr 2+, Ba 2+,etc. Group IIIA elements  +3 charge: Al 3+ Group VA elements  -3 charge: N 3-, P 3- Group VIA elements  -2 charge: O 2-, S 2-, Se 2- Group VIIA elements  -1 charge: F –, Cl –, Br –, I –,etc.

23 Ionic Bonds and Ionic Compounds Another type of chemical bond: IONIC BOND Ions in an ionic compound are held together by ionic bonds. ionic bond: electrostatic attraction holding together positively charged metal cations and negatively charged nonmetal anions

24 Ionic Bonds and Ionic Compounds Formula unit: most basic entity of an ionic compound (eg. NaCl, Al 2 O 3, etc.) –The formula gives the ratio of ions (not actual #). –The 3D representation of NaCl at the right shows a network of Na+ (purple) and Cl– ions (green). –The formula, NaCl, indicates a 1-to-1 ratio of Na+ ions and Cl– ions present, not the presence of only one ion of each.

25 An ionic compound is a network of ions, with each cation surrounded by anions, and vice versa. To melt the solid, these bonds must be broken! Ionic compounds have very high melting points compared to molecules like water. Ionic Bonds and Ionic Compounds

26 2.6 Naming Compounds and Writing Their Formulas Every element can be identified using a name, symbol, or atomic number. Know the names and symbols for the first 18 elements on the Periodic Table as well as those elements included in Figure 2.17 on p. 57, and uranium (U). The Periodic Table that will be given on quizzes and exams will include only the element symbols, atomic number, and atomic mass. Spelling counts!

27 Ionic Compounds CATIONS: positively charged ions –Metal atoms lose valence electrons to form cations. I. Groups IA, IIA, IIIA elements, silver, zinc and cadmium form only one type of ion each. –Group IA elements  +1 charge always (e.g. Li + =lithium ion) –Group IIA elements  +2 charge always (e.g. Mg 2+ =magnesium ion) –Group IIIA elements  +3 charge always (e.g. Al 3+ =aluminum ion) –silver ion = Ag + ; zinc ion = Zn 2+ ; cadmium ion = Cd 2+

28 II. The Stock system is used to name transition metals, Pb, and Sn forming different charges. –iron (Fe), a transition metal, forms two different ions: Fe 2+ and Fe 3+ –lead (Pb), in Group IVA, forms two different ions: Pb 2+ and Pb 4+ Ionic Compounds

29 ANIONS: formed only by nonmetals When a nonmetal forms an ion, it is named: element stem name + -ide suffix + ion O = oxygen atom  O 2– = oxide ion N = nitrogen atom  N 3– = nitride ion Ionic Compounds

30 Polyatomic Ions Know the formulas and names of the following polyatomic ions: CrO 4 2– = chromate ion ClO – = hypochlorite ion C 2 H 3 O 2 – = acetate ion NO 3 – = nitrate ion ClO 2 – = chlorite ion PO 4 3– = phosphate ion NO 2 – = nitrite ion ClO 3 – = chlorate ion CN – = cyanide ion OH – = hydroxide ion NH 4 + = ammonium ion CO 3 2– = carbonate ion Hg 2 2+ = mercury (I) ion Cr 2 O 7 2– = dichromate ion HCO 3 – = hydrogen carbonate ion SO 4 2– = sulfate ion MnO 4 – = permanganate ion SO 3 2– = sulfite ion ClO 4 – = perchlorate ion SCN – = thiocyanate ion O 2 2– = peroxide ion

31 Naming Ionic Compounds 1. Get the individual ions for each compound 2. CATION NAME + ANION NAME, minus “ion”  Name of compound ZnS = __________________  ______________________________ individual ions name of compound Fe 2 (CrO 4 ) 3 =___________________________ BaSO 4 =___________________________ Ni(OH) 3 =___________________________ Hg(ClO 4 ) 2 =___________________________ Cu 3 (PO 4 ) 2 =___________________________ CdCO 3 =___________________________ PbO 2 =___________________________ K 2 O 2 =___________________________

32 Given the name of a compound, predict the formula: KNOW charges on ions formed by representative elements! KNOW how to use polyatomic ions and their charges when given to you! aluminum cyanide: __________________  ________________________ individual ions formula of compound copper(I) sulfide:________________  ____________________________ barium chlorate:________________  ____________________________ zinc phosphate:________________  ____________________________ cobalt(III) carbonate:________________  ____________________________ silver nitrate:________________  ____________________________ tin(II) permanganate:________________  ____________________________ calcium dichromate:________________  ____________________________ Naming Ionic Compounds

33 Binary Molecular Compounds # of atoms Greek prefix # of atoms Greek prefix 1mono6hexa 2di7hepta 3tri8octa 4tetra9nona 5penta10deca Binary molecular compounds are composed of 2 nonmetals NAMING: # of atoms of element indicated by Greek prefix before element name 1.For first element, Greek prefix + element name 2.For second element, Greek prefix + element name stem + "ide" –If only one atom present, “mono-” is usually omitted, except for CO=carbon monoxide.

34 Examples Example:CO 2 = carbon dioxide CCl 4 = ­_________________________ SF 6 =____________________________ P 4 O 10 =__________________________ Cl 2 O 5 = ____________________________ Some binary molecular compounds also have common names e.g. everyone knows (or should know) H 2 O is water Other common compounds and names you must know: NH 3 = ammonia CH 4 = methane H 2 O 2 = hydrogen peroxide

35 Acids ACIDS: Aqueous solutions of a compound that releases H + ions –usually have H in front, physical state indicated as aqueous (aq) –naming depends on the ion from which the acid forms F – = fluoride ion HF(aq) = hydrofluoric acid NO 2 – = nitrite ion HNO 2 (aq) = nitrous acid NO 3 – = nitrate ion HNO 3 (aq) = nitric acid For some acids, the stem name changes: PO 4 3– = phosphate ion H 3 PO 4 (aq) = phosphoric acid

36 Example Problems Cl –  CO 3 2–  SO 3 2–  C 2 H 3 O 2 –  Br –  SO 4 2–  ClO – 

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