Download Chapter 2 (Hill/Petrucci/McCreary/Perry This chapter deals with

Chapter 2 (Hill/Petrucci/McCreary/Perry
This chapter deals with matter, in its various forms, on the nanoscale.
While we believe that all matter is composed of protons, electrons and, usually, neutrons, we
encounter matter as elements in the form of atoms (and, occasionally, molecules), as compounds
in the form of molecules and crystal lattices and as ions, charged particles.
Flow chart for classification of matter – this should be a review.
Alchemy (800-1200 A.D.) Study and classification of a large number of substances.
Separations: crystallization, distillation. Relationship between experiment and
theory. “Transmutation" of "base" metals into gold.
1600-1700
Galileo. Chemical changes involve rearrangements of parts too small to be seen Separations:
crystallization, distillation; Relationship between experiment and theory; “Transmutation" of
"base" metals into gold.
Francis Bacon
Heat originates from the motion of very small particles.
Robert Boyle (1627-1691)
Theory must be backed up by scientific "facts" (beginnings of modern chemistry). Boyle's
definition of an element: one must test an "element" to see if it can be chemically broken down
into simpler substances Use of atomic concepts to explain physical phenomena.
"Phlogiston" Theory (1650-1790) Georg Stahl
Phlogiston theory: an attempt to explain fire and burning process ("to set on fire")
Process of burning involved loss of "phlogiston" to air, a carrier for phlogiston
Combustion and rusting - the same
Rusting metals gained weight - "negative" phlogiston
Joseph Priestley (1733-1804)
Prepared/described a number of previously unknown gases - oxygen (1774)
mercury oxide + heat " mercury + O2
Discovery of oxygen critical to:
1.clear distinction between various pure gases and air
2.understanding most fundamental process: FIRE !!
Antoine Lavoisier (1743-1794)
Masses of chemicals which reacted in closed vessels were not changed
- Used Boyle's definition of elements to identify a number of new elements
- Elementary Treatise on Chemistry
- Combustion is a reaction between a substance and oxygen -- NO PHLOGISTON!!!
- Established chemistry as a quantitative science -- accurate weighings
Law of Conservation of Matter (Mass)
Matter neither gained nor lost in a chemical reaction
Substances can be “created” or “destroyed,” but matter cannot .
Joseph Proust (1799)
Chemical compounds have constant, unvarying compositions with respect to weights of elements
which make them up.
Law of Constant Composition (Definite Proportions)
"In a compound, elements are always present in a definite proportion by weight."
Thus, water, a compound is always
11 parts H and 89 parts O by weight (mass)
John Dalton’s Atomic Theory (1803-1807)
A comprehensive theory (a synthesis) to account for the observations of Lavoisier and Proust
Dalton's atomic theory had several postulates:
1. all matter is composed of small, invisible particles called atoms
2. in chemical reactions, atoms are neither created nor destroyed
3. atoms of each element have unique properties - all atoms of a given atom are identical and
have identical masses and other properties
4. chemical reactions involve the uniting or the separation of atoms of different elements
Dalton's experiments with various elements led him to propose relative atomic "weights" for
some of the more common elements: hydrogen, 1; carbon, 6; oxygen, 8.
Because Dalton did not and could not know that some elements (H2 , O2 , N2 , Cl2 , Br2 , I2 ) were
diatomic in nature, his atomic weights were wrong. Water, for example, which we now know is
H2 O, Dalton thought it was HO. Nevertheless, his “weights” did work for CO and for CO2
The other error in Dalton's thinking involved the idea that atoms of a given element had exactly
the same mass. They don’t ... necessarily ....
DIVISIBLE ATOMS
The atomistic view of matter was seen to be correct – but, if so, what were atoms? truly
indivisible?
As experiments toward the end of the 19 th Century were to demonstrate, atoms were divisible.
Protons
+1 charge, relative mass = 1.00
Neutrons
0 charge, relative mass = 1.00
Electrons
-1 charge, relative mass = 1/1836 (0.000545)
All neutral atoms of a given element have the same numbers of protons and electrons ... zero
charge!
ISOTOPES
We now know that isotopes exist for various elements. Isotopes are forms of an element whose
atoms differ in mass because their nuclei have different numbers of neutrons.
Using an instrument called a mass spectrometer, we can now determine atomic masses (in units
called atomic mass units or a.m.u. s) for various isotopes with extreme accuracy.
ISOTOPIC NOTATION
A way of describing atomic "makeup“ for nuclei of atoms
# PROTONS (P)
#NEUTRONS (N)
determine type of atom
determine isotope
EXAMPLE: 18 X8
Upper number is "mass number", A
A=#P + #N
Lower number is "atomic number," Z
Z = #P
Thus, A - Z = ???
and, A - Z = # N (number of neutrons)
Why is this important ?? because the number of neutrons determines what isotope we are
talking about.
EXAMPLE: 18 X8
A = 18
Z= 8
# P = 8 (atomic number)
# N = 18 - 8 = 10 a
The element is oxygen (O). Z = 8
See EXAMPLE and Exercises in Hill, p. 41
ATOMIC MASSES
If we measure the mass of a single hydrogen atom, it may vary, depending on the “kind” of
hydrogen atom we have.
1
2
3
Hydrogen has three isotopes: H1 , H1 , and H1 .
3
2
The isotope H1 , called tritium, has 1 proton and 2 neutrons; deuterium, H1 , has 1 proton and 1
1
neutrons, while protium, H1 , has only 1 proton.
The various isotopes of an element are present in fixed percentages in nature .. isotopic
abundances
The various isotopes of an element are present in fixed percentages, e.g. hydrogen (3 isotopes)
1H
1
2H
1
3H
1
0.99985
0.00015
Negligible
99.985%)
(0.015%)
1.007825 amu
2.0140 amu
3.01605
See and do “Average Atomic Weights” Exercise
Now look at Example 2.4 and Exercises 2.4A and 2.4B, p. 43
The Periodic Table - Observations
Mendeleev (1865)
Organization of known elements in order of increasing atomic mass and grouped by recurring
chemical and physical properties into groups or “families”
Prediction of new elements in “gaps”
“periodicity” concept
Mendeelev’s order failed (K, Ni, I)
The Periodic Table - Revisited
Moseley (1914)
Determined the number of protons in atoms using x-rays. Arranged elements by their numbers
of protons (order number) or atomic number
Elements now “fit” the table correctly
Modern Periodic Law: Chemical and physical properties of elements are periodic functions of
the number of protons in the nuclei of their atoms
Molecules And Compounds
Operational Definitions ...
molecule: a group of two or more atoms held together by chemical “bonds,” in which atoms
“share” electrons
molecular compound: a pure substance consisting of molecules, and in which different
elements are present in fixed proportions by mass and in fixed atomic ratios
molecular formula: a formula, consisting of elemental symbols and numerical subscripts, that
fully describes a unit of a molecular compound
empirical formula: the simplest whole-number formula for a molecular compound
Example.
C6 H6
C1 H1
molecular formula (benzene)
empirical formula for benzene
structural formula: a formula that shows how atoms are connected together in a molecule of a
molecular compound – this type of formula, however, does not necessarily tell us anything about
the 3-D shape of the molecule
See examples.
Chemical Nomenclature ...
Chemical Formulas and Chemical Names
Binary compounds: compounds that consist of only two elements, e.g H2 O (molecular) or NaCl
(ionic)
Naming Binary Molecular Compounds
Binary molecular compounds consist of a combination of two nonmetals, e.g. CO2 .
Numeric Prefixes – Indicate How Many Atoms
1
2
3
4
5
monoditritetrapenta-
6
7
8
9
10
hexaheptaoctanonadeca-
See Table 2.3, p. 49 ... be sure to learn these!
Recall: Binary molecular compounds consist of a combination of two nonmetals, e.g. CO2
Naming: prefix + nonmetal 1 + prefix + nonmetal 2 prefix + -ide
Examples.
P2 O5
C3 H8
diphosphorus pentoxide
tricarbon octahydride
See “Practice Problems for Ionic Formulas and Nomenclature,” #4 and #5
More Examples.
N2O
SiF 4
H2 O
NBr3
CO
CO2
Molecules and Compounds: Ions and Ionic Compounds
ion: a chemical species that has an excess or a deficit of electrons, and hence, a net charge
+
cation: an ion with a net positive charge, e.g. Na
A “monatomic” cation can be formed from a neutral atom when the atom loses one or mo re of its
electrons
0
Example. Na (11 protons (+), 11 electrons (-1) )
Net charge = +11 – 11 = 0
+
Na (11 protons (+), but only 10 electrons (-1) )
Net charge = +11 – 10 = +1
2-
anion: an ion with a net negative charge, e.g. S
A “monatomic” anion can be formed from a neutral atom when the atom gains one or more
additional electrons
0
Example. S (16 protons (+), 16 electrons (-1) )
Net charge = +16 – 16 = 0
2S (16 protons (+), but 18 electrons (-1) )
Net charge = +16 – 18 = - 2
See Figure 2.10 (p. 51) for monatomic ions
Naming Binary Ionic Compounds
Recall: Binary ionic compounds consist of a combination of a metal and a nonmetal, e.g. NaCl
Naming: metal name + nonmetal name + -ide
NO NUMERIC PREFIXES!!
Examples.
CaBr2 calcium bromide
Li2 S
lithium sulfide
FeCl2
CuI
iron(II) chloride
copper(I) iodide
Molecules and Compounds: Polyatomic Ions and Ionic Formulas
polyatomic ion: a cation or anion consisting of two or more atoms, bonded together, e.g. NO31See Table 2.4 (p. 54) for common polyatomic ions that you should know!
In ionic compounds, the total net charge is always 0.
To write the formula for an ionic compound, begin with the charge on the cation and the charge
on the anion – use the absolute value of the charge on the cation as subscript for the anion and
vice versa – by convention, write the cation first.
Molecules and Compounds: Writing Ionic Formulas
Example. Write the ionic formula for potassium sulfide.
1+
2K0 " K
and
S0 "S
Then, K2 S1
Try:
a
K2S
calcium oxide
iron(III) nitrate
Molecules and Compounds: Hydrates
hydrate: a type of compound whose formula includes a stated number of water- molecules
Example. CuSO4 • 5H2 O (blue crystals)
where the “dot” indicates “{waters of hydration.”
If we heat a compound like CuSO4 • 5H2 O, we can “drive off” the waters of hydration:
CuSO4 - 5H2 O + heat " CuSO4 (anhydrous)
blue crystals
white powder
a anhydrous ≡ without water
Molecules and Compounds: Acids
acids: compounds that ionize in water to form solutions dominated by H3 O+ ... acids generally
have H atoms that can come off as H+ ions
See Table 2.5 (p. 58) for common acids and their salts that you should know!
binary acids: acids that have H combined with one other nonmetallic element (see text
description, p. 58)
Example. HBr
hydrobromic acid
hydro - + element name – suffix + - ic + acid
Molecules and Compounds: Ternary Acids
ternary acids: acids that contain H, two nonmetal atoms, one of which is usually an O atom
(oxoacids)... the H atoms are attached to the O-atoms and not to the nonmetal atom. See Table
2.5 (p. 58) for common ternary acids and their salts that you should know! Note that HCl and
H2 S are the only binary acids in this table.
halogen-containing oxoacids: HXO n
Examples.
HClO
HClO 2
HClO 3
HClO 4
hypochlorous acid
chlorous acid
chloric acid
perchloric acid
Molecules and Compounds: Other Oxoacids
When there are only two oxoacids for a nonmetal, the one with more O atoms ends in –ic and the
one with fewer O atoms ends in -ous
Examples.
H2 SO3
H2 SO4
HNO2
HNO3
H3 PO3
H3PO4
sulfurous acid
sulfuric acid
nitrous acid
nitric acid
phosphorous acid
phosphoric acid
Molecules and Compounds: Salts
salt: a compound formed from the cation of a base (a metal ion) and the anion from an acid
To name salts of ternary acids and oxoacids, change –ous to –ite or –ic to-ate in the salt
Examples.
H2 SO3 (sulfurous acid)
Na2 SO3 (sodium sulfite)
H2 SO4 (sulfuric acid)
K2 SO4 (potassum sulfate)
HNO2 (nitrous acid)
HNO3 (nitric acid)
NaNO2 (sodium nitrite)
Ca(NO3 )2 (calcium nitrate)
How do we name these salts?
Examples. NaBrO 3 , ZnSO4 , K2CO3 , KI, AgNO3
Molecules and Compounds: Bases
bases: compounds that ionize in water to form solutions dominated by OH1- ... hydroxyl bases
generally have OH1- ions that can go into solution
See text (p. 57) for common bases from Group IA and IIA metals that you should know! Note
how we name these hydroxides ...
Example. NaOH
sodium hydroxide
metal name + hydroxide
2-
Other kinds of bases include ammonia (NH3 ) and carbonates (salts that contain the CO3 ion),
e.g. Na2 CO3
Molecules and Compounds: Organic Compounds
You should read this section (pp. 59-64), but only in a cursory fashion. We will not emphasize it
or discuss it in class.