Download Elements and Compounds

Elements and Compounds
Composition of Matter
Elements
1
Element Identification
2
Source of Light Elements
Solar Spectra
3
Detecting Elements
Molecular Spectrum of CO
4
Chemical and Physical
Properties
• Elements
• An Element is the basic form of matter – it
cannot be broken down into smaller parts
by chemical methods
Each element has a unique identity
Compounds
• A Compound is a unique substance that
is formed from two or more elements and
has different properties from all other
substances
5
Pure Substances
• Pure substances have unchanging,
definite compositions and properties.
• Table Salt: crystalline white substance that
is commonly used to modify food taste
NaCl : 1 sodium atom + 1 chlorine atom.
Two different elements, unique properties.
Properties of Substances
• A physical property is one that can be
observed without performing a change to
the substance:
• Color
• Density
• Hardness
• etc.
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Names of the Elements
• Latin names from historical context
produce unusual symbols for some
elements:
first letter is capitalized, second is lower
case
– Ferrum: Fe (Iron)
– Aurum: Au (Gold)
– Cobalt: Co
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Atomic Theory
• John Dalton used Greek ideas with
experimentation to reveal the atomic
nature of the elements
• The atom is defined as the smallest
particle of an element that retains the
properties of the bulk element
• Daltons atomic theory allows the
prediction of composition
Daltons Atomic Theory
• Matter is composed of atoms
• All atoms of each element are identical
and possess the same properties
• Chemical compounds are composed of
atoms of different elements in whole
number ratio
• Reactions produce re-arranged atoms
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Law of Constant Composition
• Any compound is made up of elements in
the same proportion by mass.
• A compound is identified by a specific
formula
Compounds and Molecules
• Two or more elements bonded to each
other are molecules or molecular
compounds
• Elements bonded together are polyatomic
i.e. Cl2
• Molecules may be formed by both types:
compounds or polyatomic elements
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Chemical Properties
• Chemical properties are related to the
identity of the element or the compound
• The ability of a substance to undergo a
chemical change: “reaction”
• Chemical change is a transformation of a
substance into a new substance
Mixtures of elements and
compounds
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Mixtures
• Heterogeneous: a non-uniform mixture
containing two or more phases
• Sample of mixture produces different
amounts of each substance in container
Homogeneous Mixtures
• One physical state with uniform properties
• May be a Solution or an Alloy
– Solution: homogeneous mixture with one
liquid phase
– Alloy: homogeneous mixture of solids with
one phase (solid solution)
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Electrons
• 1897 J.J. Thompson proves existence of
the electron
• Is common to all atoms
• Has a charge of -1
• Used “Plum Pudding” analogy : electrons
and a positive medium mixed as a solid
particle
Rutherford Experiment
• Used positively charged helium ions to
bombard thin gold foil
• Showed a dense positively charged
center- not a conglomeration of positive
and negative charges
• Calculated charge to mass ratio =-1.76 x
108 coulombs/gram From an electron
beam
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Electron Properties
• R.A. Millikan (Univ. Of Chicago)
determined charge magnitude of electrons
which led to the determination of the mass
• q= -1.6 x 10-19 Coulombs
• Mass = 9.11 x 10-28 grams (from charge to
mass ratio)
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The Nuclear Atom
Counting The Particles
• All atoms of each element are not quite
the same
• Mass of atoms of an element is variable, a
result of a new particle, the neutron
• The nuclear mass is the sum of protons
and neutrons: isotopes
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Calculating Isotopic Values
• Atomic Number = element identity
•
= Protons in nucleus
• neutrons=mass number – protons
• 90Sr has 38 protons:
• 90 – 38 = 52 neutrons
• Still has 38 electrons!
atomic
number
Every atom with an
atomic number of 1
is a hydrogen atom.
1 proton in the
nucleus
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atomic
number
Every atom with an
atomic number of 6
is a carbon atom.
6 protons in the
nucleus
atomic
number
Every atom with an
atomic number of
92 is a uranium
atom.
92 protons
in the
nucleus
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Isotopes of the
Elements
• Atoms of the same element can have
different masses.
• They always have the same number of
protons, but they can have different
numbers of neutrons in their nuclei.
• The difference in the number of
neutrons accounts for the difference in
mass.
• These are isotopes of the same
element
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Isotopes of the Same
Element Have
Equal numbers of protons
Different numbers of
neutrons
Isotopic Notation
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Isotopic Notation
6 protons + 6 neutrons
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6C
6 protons
Isotopic Notation
6 protons + 8 neutrons
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6C
6 protons
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Isotopic Notation
8 protons + 8 neutrons
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8O
8 protons
Isotopic Notation
8 protons + 9 neutrons
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8O
8 protons
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Isotopic Notation
8 protons + 10 neutrons
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8O
8 protons
Hydrogen has three isotopes
1 proton
1 proton
1 proton
0 neutrons
1 neutron
2 neutrons
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Examples of Isotopes
Element Protons Electrons
Symbol
Neutrons
Hydrogen
Hydrogen
Hydrogen
1
1
1
0
1
2
Uranium
235U
92
Uranium
238U
92
92
92
143
92
92
146
1
1
1
1
1H
1
2H
1
3H
Atomic Mass
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• The mass of a single atom is too small to
measure on a balance.
• Using a mass spectrometer, the mass
of
the
hydrogen
atom
was
determined.
A Modern Mass Spectrometer
Positive ions
formed from
sample.
Electrical field
From the intensity and positions
at slits
A mass
of the lines Deflection
on the mass
of
accelerates
spectrogram
spectrogram,
the different
ions
positive
ions. positive
is recorded.
isotopes and
their at
relative
occurs
amounts can
be determined.
magnetic
field.
5.8
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A typical reading from a mass spectrometer. The two
principal isotopes of copper are shown with the
abundance (%) given.
5.9
Using a mass spectrometer, the mass
of one hydrogen atom was determined
to be 1.673 x 10-24 g
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This number is very small.
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
The mass of a hydrogen atom is
very small.
Numbers
of thisthis
sizeproblem
are too small
for of
To overcome
a system
practical
use.
relative atomic
masses using “atomic
mass units” was devised to express the
masses of elements using simple
numbers.
1.673 x 10-24 g
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The standard to which the masses of all
other atoms are compared to was chosen
to be the most abundant isotope of
carbon.
12
6
A mass of exactly 12 atomic mass
units (amu) was assigned to
12
6
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1
1 amu is defined as exactly equal to
12
the mass of a carbon-12 atom
1 amu = 1.6606 x 10-24 g
12
6
Average atomic mass 1.00797 amu.
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Average atomic mass 39.098 amu.
Average atomic mass 248.029 amu.
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Average Relative
Atomic Mass
• Most elements occur as mixtures of
isotopes.
• Isotopes of the same element have
different masses.
• The listed atomic mass of an element
is the average relative mass of the
isotopes of that element compared to
the mass of carbon-12 (exactly
12.0000…amu
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To calculate the atomic mass multiply
the atomic mass of each isotope by its
percent abundance and add the
Average atomic
results.
mass (amu)
Isotope
Isotopic mass
(amu)
Abundance (%)
63
29
Cu
62.9298
69.09
65
29
Cu
64.9278
30.91
63.55
(62.9998 amu) 0.6909 = 43.48 amu
(64.9278 amu) 0.3091 = 20.07 amu
63.55 amu
Relationship Between Mass
Number and Atomic Number
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The mass number minus the atomic
number equals the number of neutrons in
the nucleus.
mass
number
atomic
number
109
47
atomic
number
47
mass number 109
Ag
=
=
number of
neutrons
62
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Atomic Mass Units
• Isotopic mass is determined by
comparison with a standard:
• AMU = 1/12 the mass of a 12C atom
• Boron isotopes:
– 10B is 0.83442 times the mass of 12C:
– Therefore has 10.013 AMU
– 11B has 11.009 AMU
Table Values for Atomic Mass
• Since elements exist in nature with
isotopic masses, the reported value is the
weighted average of all of these isotopes
based on their percent abundance
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Calculating atomic masses
• Nature contains 19.9% 10B and 80.1%
11B. What is the weighted average?
• 10B: 0.199 x 10.013 AMU = 1.99 AMU
• 11B: 0.801 x 11.009 AMU = 8.82 AMU
•
10.81 AMU
Binary Compounds
• Metal-Nonmetal
• Single Oxidation
• Multiple
oxidation
• Polyatomic ions
Al2O3
CaCl2
Fe2O3 ; FeCl2
Mg3(PO4)2
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Practice: Nomenclature
•
•
•
•
•
•
•
•
•
•
•
•
NaCl
Sodium Chloride
HCl (blue Stuff)
Hydrogen Chloride
H2O (blue Stuff)
Dihydrogen oxide
FeS
Iron(II) oxide
Al2O3
Aluminum oxide
SiO2 (blue stuff)
Silicon dioxide
Binary Compounds
•
•
•
•
•
•
•
•
NaCl
Sodium chloride
CaBr2
Calcium bromide
NH4Cl
Ammonium chloride
N2O
Dinitrogen oxide
•
•
•
•
•
•
•
•
FeS
Iron(II) sulfide
Mn2O3
Manganese(III) oxide
Sn(NO2)4
Tin(IV) nitrite
Fe(OH)3
Iron(III) hydroxide
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Non-metal Binary Compounds
• Use Greek Prefixes (table 4-3)
• Write first element which is closest to
metal
• Hydrogen is written second
• First element written as english name
• Second element(s) end with –ide
• Never use mono- for first element; only
remaining element
Oxides of Nitrogen
•
•
•
•
•
•
•
•
N2O
Dinitrogen monoxide
NO
Nitrogen monoxide
N2O3
Dinitrogen trioxide
N2O5
Dinitrogen pentoxide
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Binary Acids
• Hydrogen plus halogen (group VII) or
other anions
• Name of compound ends with –ic acid
• HCL – hydrogen choride
• Hydrochloric acid
• HBr- hydrogen bromide
• Hydrobromic acid
Oxy Acids
• Formed by the combination of hydrogen
and most polyatomic ions
• Name after the root of the anion plus the
following endings:
• -ic acid if the polyatomic ion ended with –
ate
• -ous if the anion ended with –ite
• Mineral acids use common names
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Oxyacids
•
•
•
•
•
•
•
•
C2H3O2- acetate ion
CO32- carbonate ion
NO3- nitrate ion
PO43- phosphate ion
ClO- hypochlorite ion
ClO2- chlorite ion
ClO3- chlorate ion
ClO4- perchlorate ion
•
•
•
•
•
•
•
•
HC2H3O2 acetic acid
H2CO3 carbonic acid
HNO3 nitric acid
H3PO4 phosphoric a.
HClO hypochlorous a.
HClO2 chlorous acid
HClO3 chloric acid
HClO4 perchloric acid
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