Chemie 1
Inhoudsopgave
Week 1 De bouw van een atoom en Atoomtheorieën .................................................................................... 2
2.1 composition of the atom .............................................................................................................................. 2
2.2 development of atomic theory ..................................................................................................................... 3
2.3 Light, atomic structure and the Bohr Atom .................................................................................................. 4
Week 2 Periodiek system, elektronenconfiguratie, Mol en molaire massa...................................................... 6
2.4 The periodic law and the periodic table ....................................................................................................... 6
2.5 electron arrangement and the periodic table .............................................................................................. 6
4.1 the mole concept and atoms ........................................................................................................................ 8
Week 3 Octetregel, stoomgrootte, bindingen ............................................................................................... 10
2.6 Valence electrons and the octet rule .......................................................................................................... 10
2.7 trends in the periodic table ........................................................................................................................ 12
3.1 Chemical bonding ....................................................................................................................................... 13
Week 4 nomenclatuur, isomeren, resonatie en stabilitiet van moleculen. .................................................... 16
3.2 naming compounds and writing formulas of compounds .......................................................................... 16
3.4 Drawing Lewis structures of molecules and polyatomic ions ..................................................................... 18
Week 5 eigenschappen van molecule, reactievergelijkingen......................................................................... 24
3.3 properties of ionic and covalent compounds ............................................................................................. 24
3.5 properties based om molecular geometry and intermolecular forces ....................................................... 24
4.3 the chemical equation ad the information it conveys ................................................................................ 25
4.4 balancing chemical equations .................................................................................................................... 26
Week 6 Redox reacties en molariteit ............................................................................................................ 26
4.8 oxidation-reduction reaction ...................................................................................................................... 26
4.9 Calculations using the chemical equation .................................................................................................. 26
8.5 oxidation-reduction processes ................................................................................................................... 27
,Week 1 De bouw van een atoom en Atoomtheorieën
2.1 composition of the atom
Electrons, protons and neutrons
The basic structure of an element is the atom. The atom is composed of three main
particles; the proton, the neutron and the electron. The atom is composed of two distinct
regions:
1. the nucleus is a small, dense, positively charged region in the center. It’s composed
of positively charged protons and uncharged neutrons.
2. Surrounding the nucleus is a diffuse region of negative charge populated by
electrons.
We present an element symbolically as follows:
The atomic number is equal to the number of protons in the atom, the mass number is
equal to the sum of the number of protons and neutrons. Electrons are so small the it is
insignificant in comparison to the mass of the nucleus.
Mas number= protons + neutrons
Number of neutrons= mass number – protons
Number of protons= number of electrons.
Isotopes
Isotopes are atoms of the same element having different masses, because they have
different amounts of neutrons. Isoptopes are often written with the name of the element
followed by the mass number. For example C-12 and C-14.
As stated earlier, the atomic number is the whole number associated with every element on
the periodic table. The other number associated with each element on the periodic table is
its atomic mass, the weighted average of the masses of each isotope that makes up the
element. Atomic mass is measured in atomic mass units (amu)
1 amu = 1.66054 x 10^-24 gram
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, 2.2 development of atomic theory
Dalton’s theory
The first experimentally based theory of atomic structure was proposed in the early 1800s by
john Dalton, an English schoolteacher, Dalton proposed the following:
1. all matter consist of tiny particles called atoms
2. an atom cannot be created, divided, destroyed, or converted to any other type of
atom.
3. Atoms of a particular element have identical properties.
4. Atoms of different elements have different properties.
5. Atoms of different element combine In simple whole-number ratios to produce
compound.
6. Chemical change involves joining, separating or rearranging atoms.
Although Dalton’s theory was founded on meager and primitive experimental information,
we regard much of it as correct today.
Evidence for subatomic particles: electrons, protons and neutrons
The next discoveries occurred almost a century later. William Crookes and Eugen Goldstein,
indicated that the atom is composed of charged particals. Crookes connected two metal
electrons at opposite ends of a sealed glass vacuum tube. When the electricity was turned
on, rays of light were observed to travel between two electrodes. They were called cathode
rays because they traveled from the cathode (the negative electrode) to the anode (the
positive electrode).
Later experiments by J.J> Thomson demonstrated the electrical and magnetic properties of
cathode rays. The rays were deflected toward the positive pole of an external electic field.
Because opposite charges attract, this indicates the negative character of the rays. In 1897,
Thomson announced that cathode rays are streams of negative particles of energy. These
particles are electrons.
The existence of the neutron was first proposed in the early 1920s, but is was not until 1932
that James Chadwick experimentally demonstrated its existence.
Evidence for the nucleas
In the early 1900s, it was believed that protons and electrons were uniformly distributed
throughout the atom. However an experiment by Hans Geiger led ernest Rutherford to
propose that the majority of the mass and positive charge of the atom was actually located
in a small, dense region, the nucleus, with small negatively charged electrons occupying a
much larger volume outside of the nucleus.
Ruthford and others had demonstrated that some atoms spontaneously ‘decay’ to produce
three types of radiation: alpha, beta and gamma. This process is known as natural
radioactivity. Geiger used materials that were naturally radio active as projectile sources,
‘firing’ the alpha particals produced at a thin metal foil target. He then observed the
interaction of the metal and alpha particles with detection screen and found that, most
alpha particals passed through the foil without being deflected. A small franction of the
maprticals were deflected.
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