Download Chapter 2 Atoms, Molecules, and Ions

CHAPTER 2
ATOMS, MOLECULES AND IONS
Over the centuries people have often wondered what matter is. Now, most agree that it is
composed of:
Atoms
The smallest particle that can be obtained chemically and is recognizable as an element (and is in
turn composed of neutrons, protons and electrons).
Molecules
Many substances are composed of identical aggregates of atoms we call molecules, eg. H2O, CO2.
Ions
Some substances contain charged species called ions. The ions may be charged atoms or
molecules but overall the charges cancel to give a neutral substance. NaCl, Na+Cl–, NaNO3 ,
Na+NO3–
Some “oft
met” ions:
Cations (+ve)
Anions(-ve)
Na+
sodium ion
Cl–
chloride ion
Ca2+
calcium ion
NO3–
nitrate ion
Al3+
aluminium ion
SO42–
sulfate ion
CH3COO–
ethanoate (acetate) ion
NH4+ ammonium ion
2:1
Daltons Atomic Theory
John Dalton (in 1808)
An element is composed of tiny particles called atoms. Atoms of the same element have identical
properties that are different to those of the atoms of other elements.
In a chemical reaction atoms are neither created, destroyed or charged into other atoms.
Compounds are formed when two or more elements combine.
When these postulates are combined they explain the earlier postulates that had been developed:
The Law of Constant Composition.
The Law of Conservation of Mass
Mass is neither created nor destroyed in chemical reactions.
Proust’s Law(1799) Law of Definite Proportions.
Different samples of the same chemical substance always contain the same proportion of elements
by mass.
The Law of Multiple Proportions
If two elements can combine in different ways to form different substances, the mass ratio are
small ,whole number multiples of each other.
The Law of Multiple Proportions
Where elements combine to form more than one compound they combine so that a fixed mass of
one combines with masses of the other in simple whole number proportions.
CO
C
12 g
CO2
O
16 g
C
12 g
1:1
O
32 g
1:2
12 g C
16 g O
Mass ratio in CO
=
12 g C = 2
Mass ratio in CO 2
32 g O
2:2
Structure of the Atoms
The Electron
When electricity is passed through a gas at low pressure, cathode rays are produced. Thomson
(1856 - 1940) in 1900 showed that the cathode rays consisted of particles that were deflected by
both magnetic and electric fields and of much lower mass than that of least heavy atom (hydrogen,
H) - so they are subatomic particles. He called them electrons. Later the mass was calculated (R.
1
A. Millikan) and shown to be about 1840
of the mass of the H atom or 9.109390 x 10-28 g.
charge on electron e–
mass of electron
-1
about 1840 of H atom
1
“Prodding atoms with sticks or firing cannon balls at paper?
2:3
Rutherford (1910) bombarded a gold foil with α-particles (He2+ nuclei from the radioactive element
radium, Ra). The α-particles were known to be about 7000 times as heavy as the electron and
their path was not expected to be affected much by a thin gold foil. The surprise was that while
most of the rays were hardly affected by the gold foil (the expected result) a few were strongly
deflected and sometimes even reflected. The only way that the observations could be explained
was that α-particles had come near some small, but massive, positively charged particle that he
called the nucleus.
Calculations showed that the nucleus must contain most of the mass of the atom and must be
very small compared to the volume occupied by the atom. The positively charged particle present
in the nucleus was called a proton. The nucleus of the hydrogen atom carries one positive charge
and is a proton. The nuclei of other atoms contain more than one proton.
Unfortunately it was nor possible to account for all of the mass of the nuclei of other atoms
there appearing to be some additional particle present. Rutherford called these neutrons. Neutrons
were observed a few years later.
protons
mass/u
1.00728
charge
+1
neutrons
1.00867
0
nucleus
Atomic Number
The atomic number is the number of protons in the nucleus and defines the element. All atoms
with a particular atomic number are atoms of the same element.
Z
atomic number
23
Na
11
mass number
A
For a neutral atom the number of protons must equal number of electrons. The nucleus of a
hydrogen atom is a proton.
Using sodium as an example:
charge
Na
12 neutrons
0
11 protons
+11
11 electrons
-11
0
2:4
total
All atoms of an element have the same atomic number, Z. However, it soon became clear that not
all have the same mass number, A. Atoms with the same atomic number, Z, but different mass
number, A, are called ISOTOPES.
The two natural isotopes of chlorine.
chlorine -35
35
17 Cl
17 p + 18 n
chlorine -37
37
17 Cl
17 p + 20 n
The Periodic Table
In the early nineteenth century chemists became interested in the physical and chemical
similarities that exist between elements. Dobereiner (1817 and 1829) published his ideas about the
existence of triads of similar elements.
Newlands (1863 - 1866) developed his “law of octaves” . Newlands arranged the element in
order of increasing atomic weight and recognized that the fist element was similar to the eighth, the
second to the ninth and so on. Mendeleev
periodic table. Todays periodic table lists the element in order of increasing atomic number (not
atomic weight/mass)
GROUPS
or
FAMILIES
PERIODS
H
Li
Be
B
C
N
O
F
The periodic table in the front of your text book has additional columns inserted.
The GROUPS have been numbered in various ways but we will use the modern system of
numbering 1 - 18 or old 1 - 8. (Make sure you understand the system)
We sub-divide the periodic table into:
Main Groups - Groups 1, 2 and 13 to 18
Transition Metals 4 to 11 or d-block (3-12) (explained later)
2:5
He
Ne
Metals on the left and non-metals on the right
Metals
Metalloids
Nonmetals
Some of the groups have names in common usage
Group 1
Group 2
Group 16
Group 17
Group 18
alkali
metals
alkaline
earths
Oxygen
family
halogens
noble gases
Molecules and Ions
Molecules
Molecular formula
Smallest complete unit of molecular substance
carbon dioxide is written as CO2. The number after the symbol tells how many of that kind of atom
is present in the molecule..
sulfur trioxide SO3;
methanol CH4O;
nitric acid HNO3 (in aqueous solution).
These are the formulas and tell little about the real structure of the molecule.
H
The structural formula for methanol is
H
C
O
H
and this tells how the atoms are connected.
H
Nearly the same information can be conveyed with the condensed molecular formula CH3OH - 3 H
connected to the C. O connected to the C and an H connected to the O. A ball and stick model
gives a much better representation.
2:6
Structural formula
H
H C O H
O
O C O
S
O
O
propanone "acetone"
H
H
C
H
C O
H C
H
H
Sometimes the atoms can be connected together in several different ways. The compound
represented are different and called ISOMERS.
ethanol C2H6O
H
H
H H
H C O C H
H
H
H C C O
H
H H
Structural isomers
CH3CH2OH
CH3OCH3
condensed formulas
Monatomic ions
23
+
11 Na
23
11 Na
+ e−
11p+ + 10e–
11p+ + 11e–
cation
16
8O + 2
16 2−
8O
e–
8p+ + 8e–
8p+ + 10e–
anion
Monatomic Ions
13
14
15
Group I
Group II
Group III
Group IV
Group V
+1
+2
Al3+
-
-3
2:7
16
17
Group VI Group VII
-2
-1
18
Group 0
0
If you count the electrons you will find that when the ions form, they usually have the same number
of electron as the previous or next noble gas. This is not always the case and some ions
(monatomic) may exhibit more than one charge. When this is the case it is necessary to distinguish
between the ions using the roman numeral after the name to indicate the particular ion intended.
chromium(III), Cr3+
iron(III), Fe3+
copper(II), Cu2+
chromic
ferric
cupric
chromium(II), Cr2+
iron(II), Fe2+
copper(I), Cu+
chromous
ferrous
cuprous
chromium
iron
copper
Anions (See page 45) chlorine gives chloride ion, Cl– etc.
Polyatomic ions
For example
NO −3
nitrate ion
NO −2
nitrous ion
OH–
hydroxide ion
NH +4
ammonium ion
Learn these - See Table 2.2 and 2.3
Th
d oxoanions.
2:8
Molecular Elements
Some elements in the normal state at room temperature are molecular. Examples are nitrogen, N2;
oxygen, O2; the halogens, eg. Cl2.
Nomenclature
Binary Compounds
Binary compounds are easy to name. The name consists of two parts:
NaCl
name of element
part of name of element + IDE
sodium chloride
HF
hydrogen fluoride
CaO calcium oxide
Al2O3 aluminum oxide
All have the name ending IDE. There is a problem where several different binary compounds are
formed by the pair of elements. In this case the element names are prefixed by a Greek multiplier
prefix.
mono
tetra
hepta
deca
1
4
7
10
di
penta
octa
undeca
2
5
8
11
tri
hexa
nona
dodeca
3
6
9
12
So, SCl2 is sulfur dichloride, SCl4 is sulfur tetrachloride and S2Cl2 is disulfur dichloride. Some binary
compounds have common names that replace the systematic names.
water, H2O; ammonia, NH3; methane, CH4. There are more examples in the textbook.
Acids
You will meet a group of compounds called acid as you do your labs. They are usually only
correctly described as acids in aqueous solutions in which they are more or less ionized.
2:9
Pure Substance
HCl(g)
hydrogen chloride
HBr(g)
hydrogen bromide
HI(g)
hydrogen iodide
HNO3
HNO2
Aqueous solution
H+(aq), Cl–(aq)
hydrochloric acid
H+(aq), Br–(aq)
hydrobromic acid
H+(aq), I–(aq)
hydriodic acid
Oxoanions
hydrogen nitrate
H+(aq), NO3–(aq)
nitric acid
H+(aq), NO2–(aq)
nitrous acid
hydrogen nitrite
etc. see page 56 of your text
Naming Ionic Compounds
Ionic compounds are named cation name first then anion name.
Thus
KNO3(s)
Ca(NO3)2(s)
FeSO4(s)
Fe2(SO4)3(s)
potassium nitrate
calcium nitrate
iron(II) sulfate
iron(III) sulfate
Practice naming compounds.
2:10
Anion name
chloride
bromide
iodide
nitrate
nitrite