Chapter 3: Elements, Compounds and the Periodic Table

Report
Chapter 3:
Elements, Compounds,
and the Periodic Table
Chemistry: The Molecular Nature
of Matter, 6E
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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
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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
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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
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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
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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
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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
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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
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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
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Properties of Subatomic Particles
Nucleus (protons
+ neutrons)
 3 Kinds of subatomic
particles of principal
interest to Chemists
Electrons
Particle
Mass (g)
Electrical
Charge
Electron
9.109391028
1
Proton
1.672641024
+1
0
1e
1
1
1 H, 1 p
Neutron
1.674951024
0
1
0n
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Chemistry: The Molecular Nature of Matter, 6E
Symbol
10
Atomic Notation
Atomic number (Z)
 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 (11 p )
 Different number of neutrons (10n )
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Atomic Notation
Isotope Mass number (A)
 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
A
 Symbolized by X
Z
Ex. What is the atomic symbol for helium?
He has 2
e –,
2n&2
Jespersen/Brady/Hyslop
p+
Z = 2, A = 4
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2 He
12
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)
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Example:
What is the isotopic symbol for Uranium235?
 Number of protons (p+) = 92
= number of electrons in neutral atom
 Number of neutrons (1n) = 143
 Atomic number (Z) = 92
 Mass number (A) = 92 + 143 = 235
 Chemical symbol = U
 Summary for uranium-235: 235
92
U
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Learning Check:
 Fill in the blanks:
symbol
neutrons
60Co
33
81Br
46
65
29 Cu
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36
protons
27
electrons
27
35
35
29
29
Chemistry: The Molecular Nature of Matter, 6E
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Your Turn!
206
An atom of 82 Pb has ___ protons, ___
neutrons, and ___ electrons.
A. 82, 206, 124
B. 124, 206, 124
C. 124, 124, 124
D. 82, 124, 82
E. 82, 124, 124
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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
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Calculating Atomic Mass
 Generally, elements are mixtures of isotopes
Ex. Hydrogen
Isotope
1H
2H
Mass
1.007825 u
2.0140 u
%Abundance
99.985
0.015
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
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Learning Check
Naturally occurring magnesium is a mixture of 3
isotopes; 78.99% of the atoms are 24Mg (atomic mass,
23.9850 u), 10.00% of 25Mg (atomic mass, 24.9858 u),
and 11.01% of 26Mg (atomic mass, 25.9826 u). From
these data calculate the average atomic mass of
magnesium.
0.7899 * 23.9850 u = 18.946 u
0.1000 * 24.9858 u = 2.4986 u
0.1101 * 25.9826 u = 2.8607 u
Total mass of average atom =
24.3053 u rounds up to 24.31 u
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Chemistry: The Molecular Nature of Matter, 6E
24Mg
25Mg
26Mg
19
Your Turn!
A naturally occurring element consists of two
isotopes. The data on the isotopes:
isotope #1
68.5257 u
60.226%
isotope #2
70.9429 u
39.774%
Calculate the average atomic mass of this element.
A. 70.943 u
0.60226 * 68.5257 u = 41.270 u
B. 69.487 u
0.39774 * 70.9429 u = 28.217 u
C. 69.526 u
69.487 u
D. 69.981 u
E. 69.734 u
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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
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Modern Periodic Table
with group labels and chemical families identified
Actinides
Note: Placement of elements 58 – 71 and 90 – 103 saves space
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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)
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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
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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
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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
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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
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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
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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
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Metals, Nonmetals, or Metalloids
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Metals
 Most elements in periodic table
Properties
 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
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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
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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
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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
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 8 Elements
Metalloids
 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)
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Your Turn!
Which of the following statements is correct?
A. Cu is a representative transition element
B. Na is an alkaline earth metal
C. Al is a semimetal in group IIIA
D. F is a representative halogen
E. None of these are correct
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Your Turn!
All of the following are characteristics of metals
except:
A. Malleable
B. Ductile
C. Lustrous
D. Good conductors of heat
E. Tend to gain electrons in chemical reactions
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Ions
Ions & Ionic Compounds
 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
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Formation of Ionic Compounds
Metal + Non-metal  ionic compound
2Na(s) + Cl2(g)
 2NaCl(s)
Na + Cl
+
Na

+ Cl
NaCl(s)
e
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Cations
Ionic Compounds
 Positively charged ions
 Formed from metals
 Atoms lose electrons
Ex. Na has 11 e– & 11 p+
Na+ has 10 e– & 11 p+
Anions
 Negatively charged ions
 Formed from non-metals
 Atoms gain electrons
Ex. Cl has 17 e– & 17 p+
Jespersen/Brady/Hyslop
Cl– has 16 e– & 17 p+
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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
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Ions of Representative Elements
 Can use periodic table to predict ion charges
 When we use North American numbering of
groups: Cation positive charge = group #
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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
N
O
F
Ne
 O2–
 What is expected charge on N?
 Move 3 spaces to right
 N3 –
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Rules For Writing Ionic Formulas
1. Cation given first in formula
2. Subscripts in formula must produce
electrically neutral formula unit
3. Subscripts must be smallest whole numbers
possible

Divide by 2 if all subscripts are even

May have to repeat several times
4. Charges on ions not included in finished
formula unit of substance

If no subscript, then 1 implied
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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
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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
46
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Chemistry: The Molecular Nature of Matter, 6E
Your Turn!
Which of the following is the correct formula for
the formula unit composed of potassium and
oxygen ions?
A. KO
B. KO2
C. K2O
D. P2O3
E. K2O2
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Chemistry: The Molecular Nature of Matter, 6E
Your Turn!
Which of the following is the correct formula for
the formula unit composed of Fe3+ and sulfide
ions?
A. FeS
B. Fe3S2
C. FeS3
D. Fe2S3
E. Fe4S6
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Chemistry: The Molecular Nature of Matter, 6E
Cations of Transition Metals
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
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Cations of Post-transition Metals
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+
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Ions of Some Transition Metals &
Post-transition Metals
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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
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Table 3.4 Polyatomic Ions
(Alternate Name in
parentheses)
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Learning Check
Ex. What is the formula of the ionic compound
formed between ammonium and phosphate
ions?
 Ammonium = NH4+
 Phosphate = PO43–
(NH4)+ (PO4)3–
(NH4)3PO4
Ex. Between strontium ion and nitrate ion?
 Strontium = Sr2+
 Nitrate = NO32–
Sr2+ (NO3)–
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Sr(NO3)2
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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
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Naming Ionic Compounds
Cations:
 Metal that forms only 1 positive ion
 Cation name = English name for metal
 Na+
sodium
 Ca2+
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
 Cr2+
chromium(II)
Jespersen/Brady/Hyslop
Cr3+
chromium(III)
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Naming Ionic Compounds
Anions:
 Monatomic anions named by adding
“–ide” suffix to stem name for element
 Polyatomic ions use names in Table 3.5
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Learning Check: Name The
Following
 K2O
potassium oxide
 NH4ClO3
ammonium chlorate
 Mg(C2H3O2)2
magnesium acetate
 Cr2O3
chromium(III) oxide
 ZnBr2
zinc bromide
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Learning Check: Determine The
Formula
 Calcium hydroxide
 Ca(OH)2
 Manganese(II) bromide
 MnBr2
 Ammonium phosphate
 (NH4)3PO4
 Mercury(I) nitride
 (Hg2)3N2
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Your Turn!
Which is the correct name for Cu2S?
A.
B.
C.
D.
E.
copper sulfide
copper(II) sulfide
copper(II) sulfate
copper(I) sulfide
copper(I) sulfite
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Chemistry: The Molecular Nature of Matter, 6E
Your Turn!
Which is the correct formula for ammonium sulfite?
a) NH4SO3
b) (NH4)2SO3
c) (NH4)2SO4
d) NH4S
e) (NH4)2S
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Naming Hydrates
 Ionic compounds
 Crystals contain water molecules
 Fixed proportions relative to ionic substance
 Naming
 Name ionic compound
 Give number of water molecules in formula using
Greek prefixes
monoditritetrapenta-
=
=
=
=
=
1
2
3
4
5
Jespersen/Brady/Hyslop
hexaheptaoctanonadeca-
=
=
=
=
=
6
7
8
9
10
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Learning Check: Naming Hydrates
 CaSO4 · 2H2O
 calcium sulfate dihydrate
 CoCl2 · 6H2O
 cobalt(II) chloride hexahydrate
 FeI3 · 3H2O
 iron(III) iodide trihydrate
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Your Turn!
What is the correct formula for copper(II) sulfate
pentahydrate?
A. CuSO4 · 6H2O
B. CuSO3 · 5H2O
C. CoSO4 · 4H2O
D. CoSO3 · 5H2O
E. CuSO4 · 5H2O
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Molecular Compounds
Molecules
 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
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Molecules vs. Ionic Compounds
Molecules
 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
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Molecular Compounds
 Formed when nonmetals combine
 C + O2  CO2
2H2 + O2  2H2O
 Millions of compounds can form from a few nonmetals
 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)
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urea (CON2H4)
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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
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O
F Ne
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3-D Shapes of Molecules
 Space filling models
 Used to give shapes of simple nonmetal hydrides
 Blue = nitrogen
 Red = water
 Yellow = fluorine
 White = hydrogen
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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
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Table
3.8
Hydrocarbons Belonging to the
Alkane Series
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Alkanes
 Boiling point increases as number of carbon
atoms increases
 Space filling models of alkanes
 Black = carbon
 White = hydrogen
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Your Turn!
Which is the correct name for C4H10?
A. methane
B. ethane
C. propane
D. pentane
E. butane
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Chemistry: The Molecular Nature of Matter, 6E
Other Hydrocarbons
Alkenes
 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
 Name = number prefix + ene
Ex. C2H2 = ethyne (acetylene)
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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)
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Your Turn!
What is the name of C4H9OH?
A. hexanol
B. propanol
C. pentanol
D. tetranol
E. butanol
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Chemistry: The Molecular Nature of Matter, 6E
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
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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?
A. C8H14
B. C8H16
C. C8H18
D. C8H17OH
E. C8H15OH
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Your Turn!
What is the correct structural formula for octane?
a) CH3CH2CH2CH2CH2CH2CH2CH3
b) CH3CH2CH2CH2CH2CH2CH3
c) C8H18
d) CH3CH2CH2CH2CH2CH2CH2CH2CH3
e) CH3CH2CH2CH2CH2CH2CH2CH2OH
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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
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Naming Binary Molecular Compounds
1. hydrogen chloride
1 H 1 Cl
2. phosphorous pentachloride
1 P 5Cl
3. triselenium dinitride
3 Se 2N
HCl
PCl5
Se3N2
 Mono always omitted on 1st element
 Often omitted on 2nd element unless more than one
combination of same 2 elements
Ex. Carbon monoxide
Carbon dioxide
CO
CO2
 When prefix ends in vowel similar to start of
element name, drop prefix vowel
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Learning Check: Name Each
Format:
 Number prefix + 1st element name
 Number prefix + stem + –ide for 2nd element
 AsF3
 HBr
=
=
arsenic trifluoride
 N2O4
 N2O5
 CO
=
=
=
dinitrogen tetroxide
 CO2
=
carbon dioxide
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hydrogen bromide
dinitrogen pentoxide
carbon monoxide
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Your Turn!
Which is the correct formula for nitrogen
triiodide?
A. N3I
B. NI3
C. NIO3
D. N(IO3)3
E. none of the above
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Chemistry: The Molecular Nature of Matter, 6E
Your Turn!
Which is the correct name for P4O10?
A. phosphorus oxide
B. phosphorous decoxide
C. tetraphosphorus decoxide
D. tetraphosphorus oxide
E. decoxygen tetraphosphide
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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
H 2O
 Ammonia
NH3
Jespersen/Brady/Hyslop
▪ Sucrose
C12H22O11
▪ Phosphine PH3
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Summary of Naming
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