Transcript Chapter 2 Atoms, Molecules, and Ions

Chapter 2
Atoms, Molecules, and Ions
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2.
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6.
Atoms and the atomic theory
Components of the atom
The periodic table
Molecules and Ions
Ionic Formulas
Names of Compounds
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Atoms and the Atomic Theory
Elements consist of tiny particles called atoms.
Atoms retain their identity in reactions.
In a compound, atoms combine in fixed ratios of
small whole numbers. ( Water = 2 H, 1 O )
Key Figures:
 Rutherford Discovered nucleus [Gold foil
experiment]
 JJ Thompson
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Electrons
 Thomson’s discovery of the electron
 Negatively charged
 Smaller than proton (1/2000) 0.005
AMU
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Components of Atoms
Relative mass
Relative charge
Location
Proton
1
+1
Nucleus
Neutron
1
0
Nucleus
Electron
0.0005
-1
Electron cloud
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Atomic Number, Mass Number
Atomic Number
= # of protons in nucleus
= # of electrons in a neutral atom. (not an ion)
Atomic Number
is characteristic of a particular element. (all
Hydrogen atoms have 1, Helium 2)
Mass Number
discovery of the nucleus
= # of protons + # of neutrons
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Isotopes:
Atoms of the same element with a
different mass number.
Protons
Neutrons
Atomic
Number
Carbon-12
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6
6
Carbon-14
6
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Nuclear
Symbol
Mass
Number
12
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Isotopes of Hydrogen
No neutrons
protium
1 neutron
2 neutrons
deuterium
tritium
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Nuclear stability
(stable isotopes)
 Small elements (Up to
atomic # 20) the stable
proton:neutron ratio = 1:1
 Carbon 12, C-12 is a
stable carbon isotope.
 C-16 is unstable.
After element 20…
Then more neutrons are
needed to mute the
repulsive force of the
protons in the nucleus.
 For heavy elements
Atomic Number =80+ the
stable P:N ratio is 1:1.5
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Groups
Groups
Groups
periods
Non
Metals
Metals
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Molecules
Usually made up of non-metal atoms
Held together by covalent bonds
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Types of formulas
Using ethyl (drinking) alcohol as an example:
Molecular Formula: C2H6O
Gives # and type of each element
Structural Formula:
Shows how atoms are bonded
Condensed Structural: CH3CH2OH
Gives structural hints
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Ions
Formation of monatomic ions
Gain or lose in order to obtain a noble gas
electron configuration.
Lose electrons: (metals)
Na  Na+ + eGain electrons: (non-metals)
F + e-  F –
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Monatomic ion charges
Cations (+)
 Group 1
 Group 2
 Aluminum
+1
+2
+3
Anions (-)
 Group 17
 Group 16
 Nitride
-1
-2
-3
Find their locations
on the periodic table
and label them with
their familiar names.
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Monatomic ions of Transition Metals
and POST-transition metals
 Many are polyvalent





i.e. multiple possible
charges.
Fe2+, Fe3+
Cu+, Cu2+
Pb2+, Pb4+
Sn2+, Sn4+
 Key monovalent ions:


Silver Ag+
Zinc Zn2+
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Polyatomic Ions
 Group of several atoms acting as an ionic
unit.

Ex. NO3 - Nitrate
 NH4+ and Hg22+ are the only common
polyatomic ions with a positive charge.
(Cations)
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Determine if the following are ionic or molecular
 KCl
 Ionic
Ionic: metal and non metal
(or polyatomic ion)
 NaNO3  Ionic
 CO2
 Molecular
 PBr3
 Molecular
 CoO
 Ionic
 CCl4
 Molecular
Molecular: 2 or more non
metals
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Formulas of Compounds
 Ionic compounds:
 Apply principle of electroneutrality.
Cation
Ca 2+
Anion
F-
Al3+
Na+
Zn2+
NO3H2PO4C2H3O2-
Formula
CaF2
Net charge
Al(NO3)3
NaH2PO4
Zn(C2H3O2)2
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Ionic Compounds dissolved in water:
“electrolytes”
 Electrolytes can carry a current to complete a
circuit.
 Ionic compounds are electrolytes
They may be strong or weak.
 Molecular compounds are non-electrolytes.
They will NOT carry a current to complete a
circuit.
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Names of compounds
Ionic:
 Join together the names of the 2 ions

Na+ (sodium) Br- (bromide)= NaBr =
sodium bromide
 Polyvalent transition metals include
charge in the name

Fe3+ (Iron III) O2- (Oxide) = Fe2O3= Iron
(III) Oxide
 Oxo-anions
 -ate, -ite, per-, hypo-
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Systematic naming of oxo-anions
Hypo-
Nitrogen
Sulfur
Chlorine
Bromine
Iodine
x
x
Hypochlorite
Hypobromite
Hypoiodide
Nitrite
Sulfite
Chlorite
Bromite
Iodide
Nitrate
sulfate
Chlorate
Bromate
Iodate
x
x
perchlorate
perbromate
periodate
(least)
-ite
(less)
-ate
(more)
Per(most)
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Practice Naming
 Na3N
 Sodium Nitride
 Cu(NO3)2  Copper (II) Nitrate
 LiBrO2
 Lithium Bromite
 LiF
 Lithium Fluoride
 BeIO4
 Beryllium Periodate
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Names of compounds
Molecular:

Use greek prefixes to indicate # of atoms
 1st element: Name + greek # if more than
one.
 2nd element: Greek # prefix + “ide” version of
element name.

CO2 = Carbon Dioxide
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Greek Number Prefixes 1-10
1. Mono
2. Di
3. Tri
4. Tetra
SF6
N2O5
CO
5. Penta
Sulfur Dioxide
6. Hexa
7. Hepta
PCl5
8. Octa
Phosphorus
Trichloride
9. Nona
10. deca
H2O
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Molecular compounds with common names
Formula
Common
Name
Formula
Common
Name
H2O
Water
PH3
Phosphine
H2O2
Hydrogen
Peroxide
AsH3
Arsine
NH3
Ammonia
NO
Nitric Oxide
N2H4
Hydrazine
N2O
Nitrous Oxide
CH4
Methane
C2H2
(a rocket fuel)
Acetylene
(welding gas)
(laughing gas)
(natural gas)
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Acids
 Compounds with an “H” that ionizes in water.
 HCl, in water is an electrolyte of H+ and Cl- ions.
Acid naming
Binary (2 types of atoms)
use the “–ic” suffix
HCl
Hydrochloric acid
HI
Hyrodiodic acid
HBr
Hyrdrobromic acid
HF
Hyrdrofluoric acid
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Acids
 Oxo acids:
-ate salt = -ic acid
NO3- is nitrate so HNO3 is

Nitric Acid
-ite salt = -ous acid
NO2- is nitrite so HNO2 is

Nitrous Acid
HClO
HClO2
HClO3
HClO4
Hypoclorous Acid
Chlorous Acid
Chloric Acid
Perchloric Acid
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