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William L Masterton Cecile N. Hurley Edward J. Neth University of Connecticut Chapter 2 Atoms, Molecules,

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Presentation on theme: "William L Masterton Cecile N. Hurley Edward J. Neth University of Connecticut Chapter 2 Atoms, Molecules,"— Presentation transcript:

1 William L Masterton Cecile N. Hurley Edward J. Neth University of Connecticut Chapter 2 Atoms, Molecules, and Ions

2 Introduction Atoms Composed of electrons, protons and neutrons Molecules Combinations of atoms Ions Charged particles

3 Greeks: Empedocles and Democritus Suggested the concept of atoms but were not taken seriously or credited with an atomic theory

4 John Dalton: credited with the first atomic model

5 Figure John Dalton and Atomic Theory

6 Atomic Theory 1.An element is composed of tiny particles called atoms 2.All atoms of the same element have the same chemical properties 3.In an ordinary chemical reaction, atoms rearrange their bonds but atoms are not created or destroyed 4.Compounds are formed when two or more atoms of different element combine

7 Fundamental Laws of Matter

8 Law of Conservation of Mass Matter is conserved in chemical reactions This applies to all chemical reactions but DOES NOT include nuclear reactions

9 Law of Constant Composition Compound always contains the same elements in the same proportions by mass. Pure water has the same composition everywhere.

10 Law of Multiple Proportions The masses of one element that combine with a fixed mass of the second element are in a ratio of small whole numbers. Compare CO and CO 2

11 Figure A – The Law of Multiple Proportions Two different oxides of chromium

12 Components of the Atom Atomic theory raised more questions than it answered Could atoms be broken down into smaller particles 100 years after atomic theory was proposed, the answers were provided by experiment Finding the Electrons: Protons: Neutrons:

13 J.J. Thomson Discovered the electron

14 Figure 2.2 – J.J. Thomson and Ernest Rutherford

15 Figure 2.3 – Cathode Ray Apparatus

16 Electrons First evidence for subatomic particles J.J. Thomson in 1897 Rays emitted were called cathode rays Rays are composed of negatively charged particles called electrons Electrons carry unit negative charge (-1) and have a very small mass (1/2000 the lightest atomic mass)

17 J.J. Thomson’s Model Every atom has at least one electron Atoms are known that have one hundred or more electrons There is one electron for each positive charge in an atom Electrical neutrality is maintained

18 Ernest Rutherford: Discovered the nucleus of the atom

19 Gold Foil Experiment: Bombardment of gold foil with α particles (helium atoms minus their electrons) Expected to see the particles pass through the foil Found that some of the alpha particles were deflected by the foil Led to the discovery of a region of heavy mass at the center of the atom = nucleus

20 Figure 2.4 – Rutherford Backscattering

21 Nuclear Particles 1. Protons Mass nearly equal to the H atom Positive charge 2. Neutrons Mass slightly greater than that of the proton No charge

22 Atomic Mass The average mass of all of the isotopes of an element accounting for their relative abundances

23 Table 2.1 – Subatomic Particles

24 Terminology Atomic number, Z Number of protons in the atom Mass number, A Number of protons plus number of neutrons Mass # = p + + n 0

25 Nuclear symbolism A is the mass number Z is the atomic number X is the chemical symbol

26 Isotopes Isotopes are two atoms of the same element Same atomic number but differ in number of neutrons Different mass numbers Mass # = p + + n 0

27 Example 2.1

28 Radioactivity Radioactive isotopes are unstable (Radioactive decay is not a chemical process) 1. These isotopes decay over time 2. Emit other particles and are transformed into other elements Particles emitted 1. Beta (β) particles: High speed electrons 2. Alpha (α) particles: helium nuclei 3. Gamma (γ) rays: high energy light

29 Nuclear Stability depends on the neutron/proton ratio For light elements, n/p is approximately 1/1 For heavier elements, n/p is approximately 1.4/1

30 Figure 2.5 – The Nuclear Belt of Stability

31 2.3 Introduction to the Periodic Table Dmitri Mendeleev: Arranged elements by chemical properties Left space for elements unknown at the time Predicted detailed properties for several undiscovered elements: Sc, Ga, Ge By 1886, all these elements had been discovered, and with properties similar to those he predicted

32 Mendeleev’s P.T.

33 Introduction to the Periodic Table

34 Modern Periodic Table Period – a horizontal row on the periodic table Group – a vertical column on the periodic table Blocks – sections of elements with common properties Families – another name for group; emphasizes the similarity in properties within a group

35 Blocks in the Periodic Table Main group elements 1-2, OR roman numeral +A groups Transition Metals 3-12 OR non roman numeral groups Inner Transition/Rare Earth elements Bottom double rows

36 Families with Common Names (label on PT) Alkali Metals, Group 1(I) Alkaline Earth Metals, Group 2 (II) Halogens, Group 17 (VII) Noble Gases, Group 18 (VIII)

37 A Look at the Sulfur Group Sulfur (nonmetal), antimony (metalloid) and silver (metal)

38 Example 2.3

39 2.4 Molecules and Ions Molecule: Two or more atoms chemically combined 1. Atoms involved are often nonmetals 2. Covalent bonds are strong forces that hold the atoms together Molecular formulas: Number of each atom is indicated by a subscript Examples Water, H 2 O Ammonia, NH 3

40 Structural Formulas Structural formulas: a formulas that shows the bonding patterns within the molecule

41 Ions A charged particle that is the result of the loss or gain of electrons Cation – a positive ion (loss) Anion – a negative ion (gain) Examples: Na → Na + + e - O + 2e - → O 2-

42 Ionic Compounds Compounds formed from the electrostatic attraction of oppositely charged particles Sodium chloride (NaCl): Sodium cations and chloride anions associate into a continuous network

43 Forces: Ionic compounds are held together by strong forces Compounds are usually solids at room temperature High melting points often water-soluble

44 Solutions: When an ionic compound dissolves in water, the ions are released from each other conductivity – the ions in a solution support the transmission of an electric current Strong electrolytes – solutions that are very good conductors Weak electrolytes – solutions that are poor conductors Nonelectrolytes – solutions that do NOT conduct

45 Figure 2.12 – Electrical Conductivity

46 Formulas for Ionic Compounds Charge balance Each positive charge must have a negative charge to balance it Calcium chloride, CaCl 2 Ca 2+ Two Cl - ions are required for charge balance

47 Transition Metals Polyvalent – exhibit multiple positive charges depending on conditions Iron forms Fe 2+ and Fe 3+ Lead forms Pb 2+ and Pb 4+

48 Polyatomic Ions Groups of atoms may carry a charge; these are the polyatomic ions OH - NH 4 +

49 Noble Gas Connections Atoms that are close to a noble gas (group 18 or VIII) form ions that contain the same number of electrons as the neighboring noble gas atom +1, +2, +3 skip -3, -2, -1 Noble Gases

50 Example 2.5

51 2.6 Naming of Compounds Cations: element name Na +, sodium If polyvalent, a Roman numeral is used to denote the charge Fe 2+ iron(II)

52 Names of Compounds - Anions Monatomic anions are named by adding –ide to the element name Oxygen becomes oxide, O 2- Polyatomic ions keep their names

53 To name an ionic compound: name the cation first, then, name the anion (with the word 'ion' omitted). It is not necessary to indicate the number of cations and anions in the compound because it is understood that the total positive charges carried by the cations must equal the total negative charges carried by the anions.

54 KIpotassium ion + iodide ion = potassium iodide CoCl 2 cobalt(II) ion + two chloride ions = cobalt(II) chloride CoCl 3 cobalt(III) ion + three chloride ions = cobalt(III) chloride Hg 2 Cl 2 mercury(I) ion + two chloride ions = mercury(I) chloride AgNO 3 silver ion + nitrate ion = silver nitrate

55 Oxoanions Per _________ate ___________ate ___________ite Hypo_________ite

56 Table 2.3 – Oxoanions of Nitrogen, Sulfur and Chlorine

57 Binary Molecular Compounds Made of 2 nonmetal elements Never reduce subscripts Covalently bonded

58 Mono-1 di-2 tri-3 tetra-4 penta-5 hexa-6 hepta-7 octa-8 nona-9 Systematic naming 1. First name is the first element, with prefix to for number of atoms (EXCEPT NO MONO) 2. Second name is prefix with element name changed to –ide (INCLUDE MONO)

59 Some Examples Diphosphorus pentaoxide Sulfur dioxide Dinitrogen tetraoxide Hydrogen dioxide Carbon monoxide Phosphorus trichloride

60 Acids Ionic compounds with Hydrogen as the cation Naming: Common: (strong acids) HBrHIHCl H 2 SO 4 HNO 3 HClO 4 Br I Cl SO NO ClO 434

61 Oxyacids or Oxoacids: Acids with and oxoanion as the anion

62 Acids of Chlorine (example):

63 Examples: Hydrogen chloride (hydrochloric acid) Nitric acid Sulfuric acid Hypobromous acid Nitrous acid Phosphoric acid


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