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Chapter 3: Elements, Compounds, and the Periodic Table

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1 Chapter 3: Elements, Compounds, and the Periodic Table
Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop

2 Discovery of Subatomic Particles
Late 1800s & early 1900s Cathode ray tube experiments showed that atoms are made up of subatomic particles Discovered negatively charged particles moving from Cathode – negative electrode to Anode – positive electrode 3.1 | Internal Structure of the Atom

3 Discovery of Electron JJ Thomson (1897) Modified cathode ray tube
Made quantitative measurements on cathode rays Discovered negatively charged particles Electrons (e) Determined charge to mass ratio (e/m) of these particles e/m = 1.76 x 108 coulombs/gram

4 Millikan Oil Drop Experiment
Determining charge on Electron Calculated charge on electron e = 1.60 x 1019 C Combined with Thomson’s experiment to get mass of electron m = 9.09 x 1028 g

5 Discovery of Atomic Nucleus
Rutherford’s Alpha Scattering Experiment Most alpha () rays passed right through gold A few were deflected off at an angle 1 in 8000 bounced back towards alpha ray source Gave us current model of nuclear atom Alpha particles were known to be positively charged and about 4000 times as massive as an electron. (Actually this experiment was performed by 2 post-doctoral students in Rutherford’s Lab, Geiger and Marsden) According to Rutherford, “it is almost as if you had fired a 15-inch shell at a piece of tissue paper and had it come back at you.”

6 Discovery Of Proton Discovered in 1918 in Ernest Rutherford’s lab
Detected using Mass Spectrometer Hydrogen had mass 1800x mass of electron Masses of other gases whole number multiples of mass of hydrogen Proton Smallest positively charged particle

7 Rutherford’s Nuclear Atom
Demonstrated that nucleus: has almost all of mass in atom has all of positive charge is located in very small volume at center of atom Very tiny, extremely dense core of atom Where protons (p+) & neutrons (1n) are located

8 Atomic Structure Electrons (e)
Very low mass Occupy most of atom’s space Balance of attractive & repulsive forces controls atom size Attraction between protons (p+) & electrons (e) holds electrons around nucleus Repulsion between electrons helps them spread out over volume of atom In neutral atom Number of es must equal number of p+s Diameter of atom ~10,000 × diameter of nucleus Nucleus to atom diameter like an english pea in a football stadium.

9 Discovery of Neutron (1n)
First postulated by Rutherford & coworkers Estimated number of positive charges on nucleus based on experimental data Nuclear mass based on this number of protons always far short of actual mass About ½ actual mass Therefore, must be another type of particle Has mass about same as proton Electrically neutral Discovered in 1932 by Chadwick Caused free neutron to be created

10 Properties of Subatomic Particles
Nucleus (protons + neutrons) Electrons 3 Kinds of subatomic particles of principal interest to Chemists Particle Mass (g) Electrical Charge Symbol Electron 1028 1 Proton 1024 +1 Neutron 1024 Table 3.1 /Figure 3.5

11 Atomic Notation Atomic number (Z) Isotopes
Number of protons that atom has in nucleus Unique to each type of element Element is substance whose atoms all contain identical number of protons Z = # protons Isotopes Atoms of same element with different masses Same number of protons ( ) Different number of neutrons ( )

12 Atomic Notation Isotope Mass number (A) Atomic Symbols
A = (# protons) + (# neutrons) A = Z + N For charge neutrality, number of electrons & protons must be equal Atomic Symbols Summarize information about subatomic particles Every isotope defined by 2 numbers Z & A Symbolized by Ex. What is the atomic symbol for helium? He has 2 e–, 2 n & 2 p+ Z = 2, A = 4

13 Isotopes Most elements are mixtures of 2 or more stable isotopes
Each isotope has slightly different mass Chemically, isotopes have virtually identical chemical properties Relative proportions of different isotopes are essentially constant Isotopes distinguished by mass number (A): Ex. 3 isotopes of hydrogen (H) 4 isotopes of iron (Fe)

14 Example: What is the isotopic symbol for Uranium- 235?
Number of protons (p+) = = number of electrons in neutral atom Number of neutrons (1n) = 143 Atomic number (Z) = 92 Mass number (A) = = 235 Chemical symbol = U Summary for uranium-235:

15 Learning Check: 33 27 27 46 35 35 Fill in the blanks: 60Co 81Br
symbol neutrons protons electrons 60Co 81Br 33 27 27 46 35 35

16 Your Turn! An atom of has ___ protons, ___ neutrons, and ___ electrons. 82, 206, 124 124, 206, 124 124, 124, 124 82, 124, 82 82, 124, 124

17 Carbon-12 Atomic Mass Scale
Need uniform mass scale for atoms Atomic mass units (symbol u) Based on carbon: 1 atom of carbon-12 = 12 u (exactly) 1 u = 1/12 mass 1 atom of carbon-12 (exactly) Why was 12C selected? Common Most abundant isotope of carbon All atomic masses of all other elements ~ whole numbers Lightest element, H, has mass ~1 u

18 Calculating Atomic Mass
Generally, elements are mixtures of isotopes Ex. Hydrogen Isotope Mass %Abundance 1H u 2H u How do we define Atomic Mass? Average of masses of all stable isotopes of given element How do we calculate Average Atomic Mass? Weighted average. Use Isotopic Abundances & isotopic masses

19 Learning Check Naturally occurring magnesium is a mixture of 3 isotopes; 78.99% of the atoms are 24Mg (atomic mass, u), 10.00% of 25Mg (atomic mass, u), and 11.01% of 26Mg (atomic mass, u). From these data calculate the average atomic mass of magnesium. * u = u Mg * u = u 25Mg * u = u 26Mg Total mass of average atom = u rounds up to u

20 Your Turn! A naturally occurring element consists of two isotopes. The data on the isotopes: isotope # u % isotope # u % Calculate the average atomic mass of this element. u u u u u * u = u * u = u u

21 Periodic Table Early Versions of Periodic Tables Modern Periodic Table
Summarizes periodic properties of elements Early Versions of Periodic Tables Arranged by increasing atomic mass Mendeleev (Russian) & Meyer (German) in 1869 Noted repeating (periodic) properties Modern Periodic Table Arranged by increasing atomic number (Z): Rows called periods Columns called groups or families Identified by numbers 1 – 18 standard international 1A – 8A longer columns & 1B – 8B shorter columns 3.2 | The Periodic Table

22 Modern Periodic Table with group labels and chemical families identified Actinides Note: Placement of elements 58 – 71 and 90 – 103 saves space

23 Representative/Main Group Elements
A groups—Longer columns Alkali Metals 1A = first group Very reactive All Metals except for H Tend to form +1 ions React with oxygen Form compounds that dissolve in water Yield strongly caustic or alkaline solution (M2O)

24 Representative/Main Group Elements
A groups—Longer columns Alkaline Earth Metals 2A = second group Reactive Tend to form +2 ions Oxygen compounds are strongly alkaline (MO) Many are not water soluble Accumulate in ground

25 Representative/Main Group Elements
A groups—Longer columns Halogens 7A = next to last group on right Reactive Form diatomic molecules in elemental state 2 gases 1 liquid 2 solids Form –1 ions with alkali metals—salts

26 Representative/Main Group Elements
A groups—Longer columns Noble Gases 8A = last group on right Inert—very unreactive Only heavier elements of group react & then very limited Don’t form charged ions Monatomic gases

27 Transition Elements B groups—shorter columns All are metals
In center of table Begin in fourth row Tend to form ions with several different charges Ex. Fe2+ and Fe3+ Cu+ and Cu2+ Mn2+, Mn3+, Mn4+, Mn5+, Mn6+, Mn7+ Note: Last 3 columns all have 8B designation

28 Inner Transition Elements
Lanthanide elements Elements 58 – 71 Actinide elements Elements 90 – 103 At bottom of periodic table Tend to form +2 and +3 ions. All Actinides are radioactive

29 Metals, Nonmetals, or Metalloids
Elements break down into 3 broad categories Organized by regions of periodic table Metals Left-hand side Sodium, lead, iron, gold Nonmetals Upper right hand corner Oxygen, nitrogen, chlorine Metalloids Diagonal line between metals & nonmetals Boron to astatine 3.3 | Metals, Nonmetals, or Metalloids

30 Metals, Nonmetals, or Metalloids

31 Metals Properties Most elements in periodic table Metallic luster
Shine or reflect light Malleable Can be hammered or rolled into thin sheets Ductile Can be drawn into wire Hardness Some hard – iron & chromium Some soft – sodium, lead, copper

32 Properties of Metals Conduct heat & electricity
Solids at Room Temperature Melting points (mp) > 25 °C Hg only liquid metal (mp = –39 °C) Tungsten (W) (mp = 3400 °C) Highest known for metal Chemical reactivity Varies greatly Au, Pt very unreactive Na, K very reactive

33 Nonmetals 17 elements Upper right hand corner of periodic table Exist mostly as compounds rather than as pure elements Many are Gases Monatomic (Noble) He, Ne, Ar, Kr, Xe, Rn Diatomic H2, O2, N2, F2, Cl2 Some are Solids: I2, Se8, S8, P4, C 3 forms of Carbon (graphite, coal, diamond) One is liquid: Br2

34 Properties of Nonmetals
Brittle Pulverize when struck Insulators Non-conductors of electricity and heat Chemical reactivity Some inert Noble gases Some reactive F2, O2, H2 React with metals to form ionic compounds

35 Metalloids 8 Elements Properties Semiconductors
Located on diagonal line between metals & nonmetals B, Si, Ge, As, Sb, Te, Po, At Properties Between metals & nonmetals Metallic shine Brittle like nonmetal Semiconductors Conduct electricity But not as well as metals Silicon (Si) & germanium (Ge)

36 Your Turn! Which of the following statements is correct?
Cu is a representative transition element Na is an alkaline earth metal Al is a semimetal in group IIIA F is a representative halogen None of these are correct

37 Your Turn! All of the following are characteristics of metals except:
Malleable Ductile Lustrous Good conductors of heat Tend to gain electrons in chemical reactions

38 Ions & Ionic Compounds Ions Ionic compound Formula unit
Transfer of 1 or more electrons from 1 atom to another Form electrically charged particles Ionic compound Compound composed of ions Formed from metal & nonmetal Infinite array of alternating Na+ & Cl ions Formula unit Smallest neutral unit of ionic compound Smallest whole-number ratio of ions

39 Formation of Ionic Compounds
Metal Non-metal  ionic compound 2Na(s) Cl2(g)  2NaCl(s)

40 Ionic Compounds Cations Anions Na+ has 10 e– & 11 p+
Positively charged ions Formed from metals Atoms lose electrons Ex. Na has 11 e– & 11 p+ Anions Negatively charged ions Formed from non-metals Atoms gain electrons Ex. Cl has 17 e– & 17 p+ Na+ has 10 e– & 11 p+ Cl– has 16 e– & 17 p+

41 Experimental Evidence for Ions
Electrical conductivity requires charge movement Ionic compounds: Do not conduct electricity in solid state Do conduct electricity in liquid & aqueous states where ions are free to move Molecular compounds: Do not conduct electricity in any state Molecules are comprised of uncharged particles

42 Ions of Representative Elements
Can use periodic table to predict ion charges When we use North American numbering of groups: Cation positive charge = group #

43 Ions of Representative Elements
Noble gases are especially stable Nonmetals Negative () charge on anion = # spaces you have to move to right to get to noble gas Expected charge on O is Move 2 spaces to right O2– What is expected charge on N? Move 3 spaces to right N3 – N O F Ne

44 Rules For Writing Ionic Formulas
Cation given first in formula Subscripts in formula must produce electrically neutral formula unit Subscripts must be smallest whole numbers possible Divide by 2 if all subscripts are even May have to repeat several times Charges on ions not included in finished formula unit of substance If no subscript, then 1 implied

45 Determining Ionic Formulas
Ex. Formula of ionic compound formed when magnesium reacts with oxygen Mg is group 2A Forms +2 ion or Mg2+ O is group 6A Forms –2 ion or O2– To get electrically neutral particle need 1:1 ratio of Mg2+ & O2– Formula: MgO

46 Determining Ionic Formulas
“Criss-cross” rule Make magnitude of charge on one ion into subscript for other When doing this, make sure that subscripts are reduced to lowest whole number. Ex. What is the formula of ionic compound formed between aluminum & oxygen ions? Al3+ O2– Al2O3

47 Your Turn! Which of the following is the correct formula for the formula unit composed of potassium and oxygen ions? KO KO2 K2O P2O3 K2O2

48 Your Turn! Which of the following is the correct formula for the formula unit composed of Fe3+ and sulfide ions? FeS Fe3S2 FeS3 Fe2S3 Fe4S6

49 Cations of Transition Metals
Center (shorter) region of periodic table Much less reactive than group 1A & 2A Still transfer electrons to nonmetals to form ionic compounds # of electrons transferred less clear Form more than 1 positive ion Can form more than 1 compound with same non-metal Ex. Fe + Cl FeCl2 & FeCl3

50 Cations of Post-transition Metals
9 metals Ga, In, Sn, Tl, Pb, Bi, Uut, Uuq, Uub After transition metals & before metalloids 2 very important ones – tin (Sn) & lead (Pb) Both have 2 possible oxidation states Both form 2 compounds with same nonmetal Ex. Ionic compounds of tin & oxygen are SnO & SnO2 Bismuth Only has +3 charge Bi3+

51 Ions of Some Transition Metals & Post-transition Metals

52 Compounds with Polyatomic Ions
Binary compounds Compounds formed from 2 different elements Polyatomic ions Ions composed of 2 or more atoms linked by molecular bonds If ions are negative, they have too many electrons If ions are positive, they have too few electrons Formulas for ionic compounds containing polyatomic ions Follow same rules as ionic compounds Polyatomic ions are expressed in parentheses

53 Table 3.4 Polyatomic Ions (Alternate Name in parentheses)

54 Learning Check (NH4)+ (PO4)3– (NH4)3PO4 Sr2+ (NO3)– Sr(NO3)2
Ex. What is the formula of the ionic compound formed between ammonium and phosphate ions? Ammonium = NH4+ Phosphate = PO43– Ex. Between strontium ion and nitrate ion? Strontium = Sr2+ Nitrate = NO32– (NH4)+ (PO4)3– (NH4)3PO4 Sr2+ (NO3)– Sr(NO3)2

55 Nomenclature (Naming)
IUPAC system to standardize name of chemical compounds One system so that anyone can reconstruct formula from name We will look at naming Ionic Compounds of Representative metals Transition metals Monatomic ions Polyatomic ions Hydrates 3.5 | Nomenclature of Ionic Compounds

56 Naming Ionic Compounds
Cations: Metal that forms only 1 positive ion Cation name = English name for metal Na+ sodium Ca calcium Metal that forms more than 1 positive ion Use Stock System Cation name = English name followed by numerical value of charge written as Roman numeral in parentheses (no spaces) Transition metal Cr chromium(II) Cr chromium(III) 3.5 | Nomenclature of Ionic Compounds

57 Naming Ionic Compounds
Anions: Monatomic anions named by adding “–ide” suffix to stem name for element Polyatomic ions use names in Table 3.5

58 Learning Check: Name The Following
K2O NH4ClO3 Mg(C2H3O2)2 Cr2O3 ZnBr2 potassium oxide ammonium chlorate magnesium acetate chromium(III) oxide zinc bromide Note: chromium is a transition metal that has a variable charge, hence we must specify. Since Zn has only 1 common charge, this is unnecessary. Other transition metals with only one common charge are Cd2+, and Ag+

59 Learning Check: Determine The Formula
Calcium hydroxide Ca(OH)2 Manganese(II) bromide MnBr2 Ammonium phosphate (NH4)3PO4 Mercury(I) nitride (Hg2)3N2

60 Your Turn! Which is the correct name for Cu2S? copper(I) sulfide
copper sulfide copper(II) sulfide copper(II) sulfate copper(I) sulfide copper(I) sulfite

61 Your Turn! Which is the correct formula for ammonium sulfite? NH4S
NH4SO3 (NH4)2SO3 (NH4)2SO4 NH4S (NH4)2S

62 Naming Hydrates Ionic compounds Naming mono- = 1 hexa- 6 di- 2 hepta-
Crystals contain water molecules Fixed proportions relative to ionic substance Naming Name ionic compound Give number of water molecules in formula using Greek prefixes mono- = 1 hexa- 6 di- 2 hepta- 7 tri- 3 octa- 8 tetra- 4 nona- 9 penta- 5 deca- 10

63 Learning Check: Naming Hydrates
CaSO4 · 2H2O calcium sulfate dihydrate CoCl2 · 6H2O cobalt(II) chloride hexahydrate FeI3 · 3H2O iron(III) iodide trihydrate

64 Your Turn! What is the correct formula for copper(II) sulfate pentahydrate? CuSO4 · 6H2O CuSO3 · 5H2O CoSO4 · 4H2O CoSO3 · 5H2O CuSO4 · 5H2O

65 Molecular Compounds Molecules Chemical bonds Molecular formulas
Electrically neutral particle Consists of two or more atoms Chemical bonds Attractions that hold atoms together in molecules Arise from sharing electrons between 2 atoms Group of atoms that make up molecule behave as single particle Molecular formulas Describe composition of molecule Specify # of each type of atom present 3.6 | Molecular Compounds

66 Molecules vs. Ionic Compounds
Discrete unit Water = 2 hydrogen atoms bonded to 1 oxygen atom Ionic Compounds Ions packed as close as possible to each other Sodium chloride = Each cation has 6 anions; each anion has 6 cations No one ion “belongs” to another

67 Molecular Compounds Formed when nonmetals combine
C + O2  CO H2 + O2  2H2O Millions of compounds can form from a few non- metals Organic chemistry & Biochemistry Deal with chemistry of carbon + H, N & O A few compounds have only 2 atoms Diatomics: H2, O2, Cl2, HF, NO Most molecules are far more complex Sucrose (C12H22O11) urea (CON2H4)

68 Hydrogen-containing Compounds
Nonmetal hydrides Molecule containing nonmetal + hydrogen Number of hydrogens that combine with nonmetal = number of spaces from nonmetal to noble gas in periodic table N O F Ne

69 3-D Shapes of Molecules Space filling models
Used to give shapes of simple nonmetal hydrides Blue = nitrogen Red = water Yellow = fluorine White = hydrogen

70 Organic Compounds Carbon compounds
Carbon + hydrogen, oxygen, & nitrogen Originally thought these compounds only came from living organisms Now more general Hydrocarbons Simplest organic compounds Contain only C & H Always have ratio of atoms CnH2n+2 Named using prefix designating number of C atoms All have –ane suffix

71 Table Hydrocarbons Belonging to the Alkane Series

72 Alkanes Boiling point increases as number of carbon atoms increases
Space filling models of alkanes Black = carbon White = hydrogen

73 Your Turn! Which is the correct name for C4H10? methane ethane propane
pentane butane

74 Other Hydrocarbons Alkenes Alkynes
Hydrocarbons with two less H’s than alkanes CnH2n Name = number prefix + ene Ex. C2H4 = ethene (ethylene) Alkynes Hydrocarbons with four fewer H’s than alkanes CnH2n – 2 Ex. C2H2 = ethyne (acetylene)

75 Other Organic Compounds
Hydrocarbons are basic building blocks of organic chemistry Many other classes of compounds derived from them Alcohols Replace H in alkane with -OH group Name = number prefix + anol Ex. CH3OH = methanol (methyl alcohol) C2H5OH = ethanol (ethyl alcohol) Note that ethene is commonly referred to ethylene.

76 Your Turn! What is the name of C4H9OH? hexanol propanol pentanol
tetranol butanol

77 Writing Formulas for Organic Compounds
Molecular formula Indicates # of each type of atom in molecule Ex. C2H6 for ethane or C3H8 for propane Order of atoms Carbon | Hydrogen | Other atoms alphabetically Ex. sucrose is C12H22O11 Emphasize alcohol – write OH group last C2H5OH Structural formula Indicate how carbon atoms are connected Ethane = CH3CH3 Propane = CH3CH2CH3

78 Your Turn! Octane is a hydrocarbon with 8 C atoms that is the major component of gasoline. What is the correct molecular formula for octane? C8H14 C8H16 C8H18 C8H17OH C8H15OH

79 Your Turn! What is the correct structural formula for octane?
CH3CH2CH2CH2CH2CH2CH2CH3 CH3CH2CH2CH2CH2CH2CH3 C8H18 CH3CH2CH2CH2CH2CH2CH2CH2CH3 CH3CH2CH2CH2CH2CH2CH2CH2OH

80 Nomenclature of Molecular Compounds
Goal is a name that translates clearly into molecular formula Naming Binary Molecular Compounds Which 2 elements present? How many of each? Format: First element in formula Use English name Second element Use stem & append suffix –ide Use Greek number prefixes to specify how many atoms of each element 3.7 | Nomenclature of Molecular Compounds

81 Naming Binary Molecular Compounds
hydrogen chloride phosphorous pentachloride triselenium dinitride Mono always omitted on 1st element Often omitted on 2nd element unless more than one combination of same 2 elements Ex. Carbon monoxide CO Carbon dioxide CO2 When prefix ends in vowel similar to start of element name, drop prefix vowel 1 H 1 Cl HCl 1 P 5Cl PCl5 3 Se 2N Se3N2

82 Learning Check: Name Each
Format: Number prefix + 1st element name Number prefix + stem + –ide for 2nd element AsF3 = HBr = N2O4 = N2O5 = CO = CO2 = arsenic trifluoride hydrogen bromide dinitrogen tetroxide dinitrogen pentoxide carbon monoxide carbon dioxide

83 Your Turn! Which is the correct formula for nitrogen triiodide? N3I
NIO3 N(IO3)3 none of the above

84 Your Turn! Which is the correct name for P4O10? phosphorus oxide
phosphorous decoxide tetraphosphorus decoxide tetraphosphorus oxide decoxygen tetraphosphide

85 Exceptions to Naming Binary Molecules
Binary compounds of nonmetals + hydrogen No prefixes to be used Get number of hydrogens for each nonmetal from periodic table Hydrogen sulfide = H2S Hydrogen telluride = H2Te Molecules with Common Names Some molecules have names that predate IUPAC systematic names Water H2O ▪ Sucrose C12H22O11 Ammonia NH ▪ Phosphine PH3

86 Summary of Naming


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