Transcript Chapter 2

Chapter 5
Atoms, Molecules, and Ions
1
History
Greeks
 Democritus and Leucippus - atomos
 Aristotle- 4 elements.
 Alchemy-experimentation
 1660 - Robert Boyle- experimental
definition of element.
 Lavoisier- Father of modern chemistry.
 He wrote the book.

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Laws
Conservation of Mass
 In
any chemical reaction the total
mass of the reactants will equal the
total mass of the products
36.04 g Reactants

36.04 g Products
2H2 + O2
2H2O
4.04g + 32.00g = 36.04g
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Law of Definite Proportion The
law stating that a pure substance,
e.g. H2O, will always have the same
percent by weight, e.g. 11.2% H and
88.8% O.
 elements will combine in specific
ratios by mass.
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Law of Definite Proportions

H20 H 2 x 1.01g = 2.02g
O 1 x 16.00g = 16.00g
18.02g
2.02g H / 18.02 g H2O = 0.11 x 100 = 11%
16.00g H / 18.02 g H2O = 0.88 x 100 = 88 %
What are the proportions in a CO2 molecule?
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Law of Multiple Proportions
In chemistry, the law of multiple proportions is sometimes
called Dalton's Law after its discoverer, the English chemist
John Dalton.
The law is:
If two elements form more than one compound between
them, then the ratios of the masses of the second element
which combine with a fixed mass of the first element will be
ratios of small whole numbers.
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For example carbon oxide: CO and CO2
100 grams of carbon may react with 133 grams of
oxygen to produce carbon monoxide (CO)
100 grams of carbon may react with 266 grams of
oxygen to produce carbon dioxide (CO2)
The ratio of the masses of oxygen that can react
with 100 grams of carbon is 266:133 ≈ 2:1, a ratio
of small whole numbers.
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What?!
Compare Water H2O and Hydrogen Peroxide
H2O2
 H2O - H 2 x 1.01g = 2.02 g
O 1 x 16.00g = 16.00 g
16.00 g Oxygen / 2.02 g Hydrogen =
8gO/1gH


Water has 8 g of oxygen per 1 g of hydrogen.
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Law of Multiple Proportions
H2O2 – H 2 x 1.01g = 2.02g
O 2 x 16.00g = 32.00g
32.00 g Oxygen / 2.02g Hydrogen =
16gO/1gH
 Hydrogen peroxide has 16 g of oxygen per
g of hydrogen.
 Compare the 2 samples
 16g
/ 8g = 2 / 1 ratio
 Small whole number ratios.

H2O2
H2
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Dalton’s Atomic Theory
1) Elements are made up of atoms
2) Atoms of each element are identical.
Atoms of different elements are different.
3) Compounds are formed when atoms
combine. Each compound has a specific
number and kinds of atom.
4) Chemical reactions are rearrangement of
atoms. Atoms are not created or destroyed.
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Experiments to determine what
an atom was

J. J. Thomson- used Cathode ray tubes
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Thomson’s Experiment
Voltage source
-
+
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Thomson’s Experiment
Voltage source
-
+
13
Thomson’s Experiment
Voltage source

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end.
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Thomson’s Experiment
Voltage source

By adding an electric field
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Thomson’s Experiment
Voltage source
+
 By adding an electric field, he found that
the moving pieces were negative
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Thomsom’s Model
Found the electron.
 Couldn’t find positive
(for a while).
 Said the atom was like
plum pudding.
 A bunch of positive
stuff, with the
electrons able to be
removed.

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Rutherford’s Experiment
Used uranium to produce alpha particles.
 Aimed alpha particles at gold foil by
drilling hole in lead block.
 Since the mass is evenly distributed in
gold atoms alpha particles should go
straight through.
 Used gold foil because it could be made
atoms thin.

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Lead
block
Uranium
Florescent
Screen
Gold Foil
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What he expected
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Because
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Because, he thought the mass was
evenly distributed in the atom.
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What he got
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How he explained it
Atom is mostly empty
 Small dense,
positive piece
at center.
 Alpha particles
are deflected by
it if they get close
enough.

+
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+
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Modern View
The atom is mostly
empty space.
 Two regions
 Nucleus- protons
and neutrons.
 Electron cloudregion where you
might find an
electron.

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Sub-atomic Particles
Z - atomic number = number of protons
determines type of atom.
 A - mass number = number of protons +
neutrons.
 Number of protons = number of electrons if
neutral.

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Symbols
A
(p+ and n0) Mass #
(p+)
X
Z
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Na
11
Atomic#
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Periodic Table
Click Link
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Metals
Conductors
 Lose electrons
 Malleable and ductile
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Nonmetals
Brittle
 Gain electrons
 Covalent bonds
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Semi-metals or Metalloids
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Alkali Metals
33
Alkaline Earth Metals
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Halogens
35
Transition metals
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Noble Gases
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Inner Transition Metals
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+1+2
-3 -2 -1
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Ions
Atoms or groups of atoms with a charge.
 Cations- positive ions - metals losing
electrons(s).
 Anions- negative ions - nometals gaining
electron(s).
 Ionic bonding- held together by the opposite
charges.
 Ionic solids are called salts.

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Polyatomic Ions
Groups of atoms that have a charge.
 No, you don’t have to memorize them.
 List on page 97 or on the back of your
periodic table.
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Naming compounds



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Two types
Ionic - metal and non metal or polyatomics.
If the metal is in the 1st or 2nd column
(representative element). NO (roman numeral) in
the name
If the metal is a transition metal a (roman
numeral) may be needed (check your periodic
table)
Covalent- we will just learn the rules for 2 nonmetals.
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Ionic compounds




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If the cation is monoatomic- Name the metal
(cation) just write the name.
Mg 2+
Magnesium
If the cation is polyatomic- name it.
NH4+
Ammonium
If the anion is monoatomic- name it but change
the ending to –ide.
Cl1Chloride
If the anion is polyatomic- just name it
PO4 3Phosphate
Practice.
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Ionic Compounds

Have to know what ions they form
off table, polyatomic, or figure it out
Two types
– Binary ionic – 2 elements (cation + anion)
– Ternary ionic – 3 elements
– (cation + Polyatomic ion)
CaS
K2S

AlPO4
K2SO4

FeS
CoI3



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Ionic Compounds

Fe2(C2O4)

MgO

MnO

KMnO4

NH4NO3

Hg2Cl2

Cr2O3
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Ionic Compounds
KClO4
 NaClO3
 YBrO2
 Cr(ClO)6

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Writing Formulas
Two sets of rules, ionic and covalent
 To decide which to use, decide what the
first word or element is.
 If it is a metal or polyatomic use ionic.
 If it is a non-metal use covalent.

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Ionic Formulas
Charges must add up to zero.
 Get charges from table, name of metal ion,
or memorized from the list.
 Use parenthesis to indicate multiple
polyatomics.

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Ionic Formulas
Sodium nitride
 sodium- Na is always +1
 nitride - ide tells you it comes from the table
 nitride is N-3

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Ionic Formulas
Sodium nitride
 sodium- Na is always +1
 Nitride - ide tells you it comes from the
table
 nitride is N-3
 Doesn’t add up to zero.

+1
Na
-3
N
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Ionic Formulas
Sodium nitride
 sodium- Na is always +1
 nitride - ide tells you it comes from the table
 nitride is N-3
 Doesn’t add up to zero
 Need 3 Na

+1
Na
-3
N
Na3N
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Ionic Compounds
Sodium sulfite
 calcium iodide
 Lead (II) oxide
 Lead (IV) oxide
 Mercury (I) sulfide
 Barium chromate
 Aluminum hydrogen sulfate
 Cerium (IV) nitrite
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Chemical Bonds
The forces that hold atoms together.
 Covalent bonding - sharing electrons.
 Makes molecules.
 Chemical formula- the number and type of
atoms in a molecule.
 C2H6 - 2 carbon atoms, 6 hydrogen atoms,


Structural formula shows the connections,
but not necessarily the shape.
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H
H C
H
H
C H
H
There are also other models that attempt to
show three dimensional shape.
 Ball and stick.

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Binary Molecular Compounds
Two words, with prefixes.
 Prefixes tell you how many atoms.
 1-mono, 2-di,
3-tri, 4-tetr(a), 5-pent(a),
6-hex(a), 7-hept(a), 8 oct(a),9-non(a), 10-dec(a)

Name : P4Cl7
First
element whole name with the appropriate prefix, except
mono
P4
=
tetraphosphorus
Second
element, -ide ending with appropriate prefix
Cl7
Practice
=
heptachloride
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Naming Covalent Compounds
CO2
 CO
 CCl4

N2O4
 XeF6
 N4O4
 P2O10

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Binary Molecular (Covalent)
Compounds
The name tells you how to write the
formula
 Sulfur dioxide
 diflourine monoxide
 nitrogen trichloride
 diphosphorus pentoxide

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More Names and formulas
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Acids
Substances that produce hydrogen (H+) ions
when dissolved in water.
 All acids begin with H.
 Two types of acids:
 Oxyacids – oxygen in the formula
 Non-oxyacids – no oxygen in the formula

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Naming Acids
Hydrogen _______ide becomes hydro____ic acid
Hydrogen_______ate becomes _________ic acid
Hydrogen_______ite becomes _______ous acid
Hydrogen always #1 on periodic table
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Using naming acids table
Name : HCl from the periodic table ionic name is
hydrogen chloride so use:
Hydrogen _______ide
chlor
becomes hydro______ic
chlor acid
Transfer word root from one side to the other
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Using naming acids table
Write the formula for sulphuric acid: use
Hydrogen_______ate
sulph
becomes _________ic
sulphur acid
Ionic name is hydrogen sulphate
H+
SO4 2So acid formula is H2SO4
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Naming acids
If the formula has oxygen in it(oxyacid)
 write the name of the anion, but change
– ate to -ic acid
– ite to -ous acid
 Watch out for sulfuric and sulfurous

H2CrO4
 HMnO4
 HNO2

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Naming acids
If the acid doesn’t have
oxygen(nonoxyacid)
 add the prefix hydro change the suffix -ide to -ic acid
 HCl
 H2S
 HCN
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Formulas for acids
Backwards from names.
 If it has hydro- in the name it has no oxygen
 Anion ends in -ide
 No hydro, anion ends in -ate or -ite
 Write anion and add enough H to balance
the charges.
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Formulas for acids
hydrofluoric acid
 chromic acid
 carbonic acid
 hydrophosphoric acid
 fluorous acid
 perchloric acid
 phosphorous acid
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Hydrates
Some salts trap water crystals when they
form crystals.
 These are hydrates.
 Both the name and the formula needs to
indicate how many water molecules are
trapped.
 In the name we add the word hydrate with a
prefix that tells us how many water
molecules.

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Hydrates
In the formula you put a dot and then write
the number of molecules.
 Calcium chloride dihydrate = CaCl22O
 Chromium (III) nitrate hexahydrate =
Cr(NO3)3 6H2O

Simulation of a Hydrate
Hydrate Nomenclature Practice
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compounds
ternary
binary
molecular
hydrates
molecules
trapped water
All nonmetals
ionic
acids
formula unit
Metal or hydrogen and nonmetals
or Polyatomics
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