AP TOPIC 2: Atoms, Ions & Nomenclature
The History of Atomic Theory
Circa. 400-5 BC. Greek philosopher Democritus proposes the idea of matter being made up of
small, indivisible particles (atomos).
Late 18th Century. Lavoisier proposes the Law of conservation of mass and Proust proposes the
Law of constant composition.
Early 19th Century. Using the previously unconnected ideas above, John Dalton formulates his
Atomic Theory.
Dalton’s Atomic Theory
(i) Elements are made from tiny particles called atoms.
(ii) All atoms of a given element are identical. (see isotopes below).
(iii) The atoms of a given element are different to those of any other element.
(iv) Atoms of different elements combine to form compounds. A given compound always has
the same relative numbers and types of atoms. (Law of constant composition).
(v) Atoms cannot be created or destroyed in a chemical reaction they are simply rearranged
to form new compounds. (Law of conservation of mass).
Structure of the Atom and The Periodic Table
Several experiments were being carried out in the 19th and 20th centuries that began to identify
the sub-atomic particles that make up the atom. A summary of those experiments is given below.
In the first part of the 20th Century, following Chadwick’s discovery of the neutron, Bohr proposed
the idea that the atom was made up of the nucleus containing protons and neutrons that was being
orbited by electrons in specific, allowed orbits. This particle model of the electron and atom
was expanded a few years after Bohr’s original ideas to incorporate the wave nature of the
electrons.
The atomic numbers (in the table below printed above the symbol and sometimes referred to as
Z) and mass numbers (in the table below printed below the symbol and sometimes referred to as
A) that appear on the periodic table have specific meanings.
Since all atoms are neutral it also tells us the number of electrons surrounding the nucleus. (note:
When atoms lose or gain electrons they become charged and form ions).
In this example Al is a metal, Si is a semi-metal (metalloid) and P is a non-metal.
Isotopes
Atoms with the same number of protons and electrons, but different numbers of neutrons are
called isotopes. This leads to the modification of Dalton’s Atomic Theory point (ii) above, to read.
All Atoms of the same element contain the same number of protons and electrons but may have
different numbers of neutrons.
Since it is the electrons in atoms that affect the chemical properties of a substance, isotopes of
the same element have the same chemical properties.
On some periodic tables you will find atomic mass numbers that are not integers. What does this
mean? A good starting point is to analyze what it does not mean.
For example, the atomic mass of Cl is often quoted on a periodic table as 35.5 and can be
represented by the following symbol;
35.5Cl
17
This does not mean that there are 17 protons, 17 electrons and 18.5 neutrons in an atom of
chlorine. It is not possible to have a fraction of a neutron, there can only be a whole number of
neutrons in an atom.
So what does it mean, and where does the 0.5 come from? Here is the explanation.
The non-integer values mean that there is more than one isotope of chlorine that exists in nature,
in this case 35Cl and 37Cl. A quick calculation will tell you that these two species have the same
number of protons and electrons, but different (whole) numbers of neutrons (18 and 20
respectively). That is, they are isotopes of one another. These isotopes happen to exist naturally
in the following abundance; 35Cl 75% and 37Cl 25%.
A simple calculation can be applied to work out the average atomic mass when considering all
the isotopes present in a natural sample.
Radioactivity
Radioactivity is the spontaneous decay of certain atoms with the evolution of alpha, beta and
gamma particles. The radiation comes from the nucleus of the atom i.e. it is a nuclear reaction.
Radioactive decay reactions
Alpha decay (the loss of a Helium nucleus)
Beta decay (a neutron splits to give a proton and an electron)
Gamma decay (a rearrangement of the nuclear particles only)
Other nuclear reactions
Positron emission (a positron has the same mass an electron, but has a positive charge)
Electron capture (a captured electron combines with a proton in the nucleus to form a neutron)
Half-life
The half-life of a radioactive nucleus is the time taken for half of the atoms to decay. It is
independent of the initial quantity of atoms. There are three methods of determining half-life.
(i) Graphically
(ii) Use of the expression
No = count rate initially
K = radioactivity constant
(iii) Use of the expression
n = number of half-lives
N = count rate at time t
t = time
Transmutation of elements
It is possible, via nuclear reactions, to artificially produce elements.
1919 Rutherford. Alpha particle bombardment.
1932 Cockcroft and Walton. H+ used as the bombarding particle, it has less charge than the alpha
particles and so feels less repulsion from the target nucleus.
Neutron induced transmutation. Neutrons being neutral feel no repulsion from the positive target
nucleus.
1970 Use of accelerated heavier nuclei. A machine will accelerate the positive nucleus toward the
target.
Mass Deficit
When atoms are formed by the combination of protons, neutrons and electrons the mass of the
atom is found to be less than that of the sum of the individual particles. This appears to contradict
the law of conservation of mass. The explanation for the mass deficit is that when the particles
combine a small amount of the mass is converted to energy (binding energy) and released to the
surroundings.
Predicting stability
As a general guideline only, stable (non-radioactive) nuclei tend to have neutron: proton ratios
closer to 1:1. Nuclei that have higher neutron: proton ratios tend to want to lower the ratio by
converting a neutron to a proton and an electron. The electrons are released as β particles.
Nuclear and electron rearrangement
Radioactivity is a nuclear process that involves rearrangement of the nuclei of atoms, whereas
chemical reactions involve the rearrangement of the electrons.
Nuclear Fission
The process of heavy nuclei capturing neutrons, splitting to form other, smaller nuclei and
releasing more neutrons is known as nuclear fission. In the process large amounts of energy can
be released, exacerbated by the production of more neutrons each time leading to a potential
chain reaction.
Nuclear Fusion
The combination of smaller nuclei into larger ones with the release of energy is known as nuclear
fusion. These reactions are less easy to perform than nuclear fission since they involve the
combination of two nuclei that are both positively charged and therefore repel one another.
Uses of Radioactivity
Some examples of the uses of radiation are listed below
1. Medicine. For example, 131I for thyroid and brain imaging, 58Co for diagnosis of pernicious
anaemia and 67Ga for lung function.
2. Isotopic dating. See questions above.
3. Thickness control in Engineering. A radioactive source is placed on one side of a sheet of
paper or metal as it is made and a detector is placed on the other side. Any change in the
thickness of the material will be reflected in a change in the observed radioactivity count.
4. Leak detection. Radioactive sources can be injected into pipe-work and will be detected
emerging where the leak is.
5. Nuclear Fission (power and atomic bomb). Uranium nuclei can be bombarded with
neutrons and converted to other nuclei. The process occurs with the loss of massive
amounts of energy in a chain reaction. These huge amounts of energy can be used in a
constructive or destructive way.
Molecules and Ions
Molecules
Molecules are formed when a definite number of atoms are joined together by chemical bonds. A
molecule can consist of the atoms of only one element, or the atoms of many different elements
but always in a fixed proportion. This means that molecules can be elements or compounds.
Molecules are usually formed between non-metal elements. Formulae show the number of each
type of atom present written as subscripts. The lack of a subscript means only one of that type of
atom is present.
Examples of molecules
Ions
Atoms have equal numbers of protons and electrons and consequently have no overall charge.
When atoms lose or gain electrons, the proton: electron numbers are unbalanced causing the
particles to become charged. These charged particles are called ions. Positive ions (where the
number of protons is greater than the number of electrons) are called cations, and negative ions
(where the number of electrons is greater than the number of protons) are called anions. Metals
tend to form cations and non-metals tend to form anions. These oppositely charged ions form
ionic compounds by attracting one another. An ion made up of only one type of atom is called a
monatomic ion; one made up from more than one type of atom is called a polyatomic ion.
Examples of ions
Nomenclature of inorganic compounds
Binary compounds of metals and non-metals (Ionic compounds)
Binary compounds are those formed between two elements. In compounds where one is a metal
and one a non-metal an ionic compound is formed. Ionic formulae and names can be determined
by considering the charge on the ions. To find the formula of an ionic compound the positive and
negative charges must be balanced, i.e. there must be no net charge. To name a binary
compound of a metal and a non-metal, the unmodified name of the positive ion is written
first followed by the root of the negative ion with the ending modified to -ide. For example,
sodium chloride.
A few common ions, their charges and formula are listed below.
Note: Most transition metal ions and a few other metal ions, include a number written as a Roman
numeral after the name. These metals form ions with varying charges and this numeral identifies
the charge in each case.
Binary acids
For the purposes of nomenclature an acid can be defined as a compound that produces
hydrogen ions when it is dissolved in water. Binary acids are formed when hydrogen ions
combine with monatomic anions. To name a binary acid use the prefix hydro followed by the
other non-metal name modified to an –ic ending. For example, hydrochloric acid.
Polyatomic ions and oxoanions
Polyatomic ions are those where more than one element are combined together to create a
species with a charge. Some of these ions can be named systematically, others names must be
learned.
Some common polyatomic ions, their charges and formula are listed below.
Polyatomic anions where oxygen is combined with another non-metal are called oxoanions and
can be named systematically. In these oxoanions certain non-metals (Cl, N, P and S) form a
series of oxoanions containing different numbers of oxygen atoms. Their names are related to the
number of oxygen atoms present, and are based upon the system below.
Where there are only two members in such a series the endings are –ite and –ate. For example,
sulfite (SO32-) and sulfate (SO42-). When there are four members in the series the hypo- and per- prefixes
are used additionally.
Some oxoanions contain hydrogen and are named accordingly, for example, HPO 42-, hydrogen
phosphate. The prefix thio- means that a sulfur atom has replaced an atom of oxygen in an anion.
To name an ionic compound that contains a polyatomic ion, the unmodified name of the
positive ion is written first followed by unmodified name of the negative ion. For example,
potassium carbonate.
Oxoacids
Oxoacids are formed when hydrogen ions combine with polyatomic oxoanions. This gives a
combination of hydrogen, oxygen and another non-metal. Such acids are called oxoacids. To
name an oxoacid use the name of the oxoanion and replace the -ite ending with –ous or
the -ate ending with -ic. Then add the word acid. For example, sulfuric acid.
To illustrate the names of these oxoanions and oxoacids consider the following example using
chlorine as the non-metal.
Binary compounds of two non-metals (Molecules)
If the two elements in a binary compound are non-metals, then the compound is molecular, i.e.
they form a molecule. To name a molecular compound of two non-metals, the unmodified
name of the first element is followed by the root of the second element with ending
modified to -ide. In order to distinguish between several different compounds with the
same elements present use the prefixes mono, di, tri, tetra, penta and hexa to represent
one, two, three, four, five and six atoms of the element respectively. For example, sulfur
dioxide.
Some other examples are given below;
Note that the prefix mono is only applied to the second element present in such compounds, if the
prefix ends with a or o and the element name begins with a or o then the final vowel of the prefix
is omitted, and that some compounds have trivial names that have come to supercede their
systematic names, for example, ‘water’, not ‘dihydrogen monoxide’.
Hydrates
Hydrates are ionic formula units with water associated with them. The water molecules are
incorporated into the solid structure of the ions. To name a hydrate use the normal name of the
ionic compound followed by the term hydrate with an appropriate prefix to show the
number of water molecules per ionic formula unit. For example, copper(II) sulfate
pentahydrate.
Strong heating can generally drive off the water in these salts. Once the water has been removed
the salts are said to be anhydrous (without water).