Transcript Chapter Two Atoms & The Periodic Table

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Compound is composed of two or more
elements combined in a specific ratio and held
together by chemical bonds.
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Cannot be broken down by physical means.
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Want to achieve stable electron
configurations
 Noble Gas configurations
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Valence electrons (outer electrons) are
involved in bonding
 # of v.e. come from the group number
 Lewis Dot Symbols (or diagrams) show this
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Draw the element symbol
Represent v.e. with dots around the symbol
 4 sides of the element symbol
▪ Start on top of symbol
 No more than 2 dots per side (max 8 for element)
 Do not double up electrons until each side has 1
v.e.
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Can be drawn for atoms or ions
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Draw the Lewis Dot Diagrams for the
following:
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Silicon atom (Si)
Bromide ion (Br-)
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What group of the periodic table does
element X belong if its Lewis Dot Symbol was
X
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Contain cations (positive ions) and anions
(negative ions)
 Bonded together through ionic bonding:
electrostatic attraction between oppositely
charged particles
▪ Results from a transfer of electrons
 Ions cancel each other out producing a neutral
compound
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With ionic bonds, electrons are
TRANSFERRED
Na
+
Cl
-
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A chemical formula or just formula indicates
the ratio of elements in a compound
 NaCl
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Li2O
K3 N
Al2O3
Held together by a lattice work
Individual compounds in a
lattice work are called
formula units.

Lattice Energy is amount of energy required
to convert a mole of an ionic solid to its
constituent ions in the gas phase
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Ex: NaCl has a lattice energy of 788 kJ/mol
 It takes 788 kJ of energy to convert 1 mole of NaCl
to 1 mole of Na+(g) and Cl-(g)
 The higher the lattice energy, the more stable
the ionic compound is
1)
2)
Ionic Radius (distance between nuclei)
Magnitude of Charges
1) Ionic Radius
 When distance between nuclei increases,
attraction decreases, therefore weakening
lattice energy
1)
Ionic Radius (distance between nuclei)
2) Magnitude of Charges
 When distances are similar between nuclei,
we must look at charges
 The higher the charge, the stronger the
attraction, the stronger the lattice energy
2) Magnitude of Charges

Also known as nomenclature
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Naming monatomic ions
 Groups 1A, 2A, and 3A elements have charges
equal to their group # (recall e- configs)
 Most anions from elements in groups 4A-7A have
charges equal to their group # is 8
▪ Ex: Oxygen Group 6 – 8 = -2
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P. 123
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Cations:
 Named simply by putting the word ion to the name of
the element
▪ Ex: K+ is the potassium ion, Al3+ is the aluminum ion
 Certain metals form more than one charge
▪ Ex: Fe2+ and Fe3+
 Older system: Fe2+  Ferrous ion; Fe3+  Ferric ion
 Newer system: Fe2+  iron (II) ion; Fe3+  iron (III) ion
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Anions
 Named by changing the element’s ending to –ide
and adding the word ion
▪ Ex: Br- is the bromide ion; O2- is the oxide ion
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Cation and anion charges must equal ZERO
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Use the criss-cross method to write formulas
showing lowest whole number ratio

Ex: Potassium Bromide, Zinc Iodide,
Aluminum Oxide
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Most ionic compounds are binary (only 2
elements: metal and nonmetal—in other
words, cation & anion)
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Ionic compounds are named by using the
cation name followed by the anion name,
eliminating the words “ion”
Ex: NaBr  Sodium Bromide
CaF2 Calcium Fluoride
Al2N3  Aluminum Nitride
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What about ionic compounds with transition
metals???
 Use “reverse criss-cross” to determine original
charge of cation (be sure anion has correct charge)
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Ex: FeCl2
Cr2S3
MnO
If given the formula, write the name
If given the name, write the formula
Ag+
sodium ion
N3-
zinc bromide
K 2O
Strontium Carbide
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Bonding typically between two nonmetals
 Electrons are shared; both atoms “feel” like they
have a noble gas configuration
F
F
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A neutral combination of at least 2 atoms in
a specific arrangement held together by a
chemical bond/force
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Can be an element (H2) or a compound (H2O)
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Molecules with only 2 atoms
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Homonuclear: two of the same element
 H2, N2, O2, F2, Cl2, Br2, I2
 HOFBrINCl
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Heteronuclear: two different elements
 CO, HCl, etc.
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Allotrope: Elements with different distinct
forms (ex: Carbon graphite & diamonds)
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Name the first element as is
 Add the appropriate prefix to signify # of atoms
▪ Exclude mono for first element if there’s only one
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Name the second element with –ide ending
 Add the appropriate prefix to signify # of atoms
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Ex: N2P5
CS
SO3
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Sulfur Tetrafluoride
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Tetraphosphorus Decasulfide
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Tribromine Heptoxide
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Ex: SnCl4 and PbCl4
 Can be named using ionic or covalent naming
▪ Tin (IV) Chloride or Tin Tetrachloride; Lead (IV) Chloride
or Lead Tetrachloride
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B2H6
SiH4
NH3
PH3
H 2O
H 2S
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*Responsible for knowing the one’s in red
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Diborane
Silane
Ammonia
Phosphine
Water
Hydrogen Sulfide
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Acids typically have hydrogen at the
beginning
 When dissolved in water, they typically produce
hydrogen ions (H+)
 HCl in gas form is hydrogen chloride
 HCl dissolved in water is hydrochloric acid
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1)
2)
3)
How to name binary acids:
Change “hydrogen” into “hydro-”
Take the root of the second element and
add “-ic” to the ending.
End with acid
Ex: HBr(aq)
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Hydrofluoric acid
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H3N(aq)
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Hydrosulfuric acid
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H3P(aq)
Name/Write the following:
 SF6
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P2I4
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Diboron Heptoxide
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Hydroiodic acid
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H2Se(aq)
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Polyatomic ions: ions that contain multiple
atoms/elements (held together by covalent
bonding)
 Act as a group; can be treated as one
Name these compounds using
the same rules as we did for ionic
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You must know table 5.10 in your
text book!
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You are responsible for
knowing the name, the
formula, and the
charge for each
polyatomic ion in table
5.10 (this is a
SHORTENED list!)
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Ammonium Chloride
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Cobalt (II) Nitrate
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Try These
Calcium Phosphate
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Manganese (III) Carbonate
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Name the following compounds:
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Fe2(SO4)
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Al(OH)3
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NH4NO2
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Acids that contain oxoanions (polyatomic
ions with oxygen)
 Ex: NO3- makes the oxoacid HNO3(aq)
 Ex: SO42- makes the oxoacid H2SO4(aq)
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Continue adding hydrogen (+1) to the
beginning of the oxoanion until the net
charge is ZERO
 Just like ionic rules
Do NOT start with hydroSTART with the root of the polyatomic ion
Add the appropriate suffix
1)
2)
3)
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If the polyatomic ends in –ate, change to –ic
If the polyatomic ends in –ite, change to –ous
4) Add the word “acid
Example: H2CrO4
HNO2
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Monoprotic only has 1 ionizable hydrogen
 HNO3, HCl, HC2H3O2
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Polyprotic has more than 1 ionizable
hydrogen
 H2S, H3PO4
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Ionization will happen in steps
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Compounds with specific number of water
molecules within its solid structure
 Use greek prefixes to name water part
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CuSO4 * 5H2O
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Anhydrous salt is the compound w/o water
 Has distinct different properties than its hydrate
form
Name/Write the following
 Sulfurous Acid
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Iron (II) Acetate
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NaOH
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H2CO3(aq)
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Strontium Nitrate Tetrahydrate
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Determine the molecular mass of the
following compounds:
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Propane, C3H8
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Lithium Hydroxide, LiOH
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Percent by mass of every element in a
compound
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Divide the mass of element by the mass of
the entire compound
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Ex: What is the percent composition of H2O?
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Molecular Formulas: The TRUE formula for a
compound
 Shows exactly how many of each element is in a
compound via subscripts
 Ex: C6H12O6, H2O
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Empirical Formulas: The lowest WHOLE
number ratio of elements in a compound
 May or may not be the true formula
 Ex: CH2O, H2O
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Ex: An unknown compound was found to
have a percent composition as follows:
47.0 % potassium, 14.5 % carbon, and 38.5 %
oxygen. What is its empirical formula?
If the true molar mass of the compound is
166.22 g/mol, what is its molecular formula?
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Nitrogen and oxygen form an extensive
series of oxides with the general formula
NxOy. One of them is a blue solid that comes
apart, reversibly, in the gas phase. It contains
36.84% N. What is the empirical formula of
this oxide?
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Write the empirical formula for the following
compounds.
C8H18
C2H6O2
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A compound with an empirical formula of
C4H4O and a molar mass of 136 grams per
mole. What is the molecular formula of this
compound?