Lecture notes Earth and Environmental Chemistry (CSM1031)
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Course
Environmental Geoscience (CSM1031)
Institution
University Of Exeter (UoEX)
This Earth and Environmental Chemistry course explores the chemical makeup of our planet and how it affects the environment. Students learn basics of topics such as organic chemistry, electrochemistry and radioactivity as well as being introduced to new concepts like phase diagrams. This is all lin...
Introduction
1. To introduce the concept of atoms and atomic structure
2. To review bonding of atoms in materials
3. To understand different types of chemical reactions
4. To discuss techniques for chemical analysis
5. To review fundamental principles of thermodynamics
6. To understand radioactive decay and isotopes• Introduction to acids, bases & pH
7. Introduction to organic chemistry
8. To review fundamental principles in electrochemistry and its applications
9. Introduction to aqueous geochemistry
10. Introduction to the chemistry of the Earth, rocks and oceans
11. To gain experience with chemical laboratory practice
Lecture Notes
Uses of geochemistry
Everything is made from atoms
Volcanic eruptions yield lava and gases
Ore deposits form from geochemical interactions between fluids and rock
Elements are used in energy generation
Environmental pollution
Understanding how microbes interact with rocks
Understanding the relative impacts of human activity and natural cycles
Understanding climate change
UN sustainable development goals
Atomic Theory
Features of atoms, moles and isotopes Electron filling order and the periodic table Defining
the position of electrons in atoms
Lecture Notes
,Earth science involves the study of rocks, water, air, soil, sediment and dust. All of these
materials are chemicals compounds [or aggregates of] which are made up of atoms that are
joined together to make elements.
-John Dalton proposed that all matter consists of minute particles called atoms, which cannot
be created or destroyed. Chemical reactions occur when atoms are rearranged and compounds
are formed from different definite numerical proportions of atoms/ elements.
-Atoms are made up of protons [positively charge], neutrons [uncharged] and electrons
[negatively charged]. Electrons orbit the nucleus in shells whilst the protons and neutrons are
contained within the nucleus.
-Atomic number defines the number of protons [which is also the number of electrons]. -
Mass number is the number of protons + the number of neutrons in the nucleus.
-A mole is the amount of a substance that contains as many elementary entities as there are
atoms in 12g of carbon-12. -In one mole there are 6.02 x10^23 elementary entities- this
number is known as Avogadro's constant [measured as atoms mol^-1] -The molar mass of an
element is the same as its atomic mass;
1Da=1gmol^-1
1Da = 1 proton = 1 neutron\
1 electron = 1/1850 Da
The mass of an atom is very small, roughly equal to the sum of the protons and neutrons in
the nucleus.
-Isotopes are elements with varied numbers of neutrons in the nucleus. E.g. uranium can have
143 or 146 neutrons.
—JJ Thomson proposed the 'plum pudding' model of the atom
-The Rutherford-Bohr model of the atom describes how electrons orbit the nucleus in
increasing energy levels as they get further from the nucleus.
-Electrons are envisaged s occupying orbitals, in which the statistical probability of finding
the electron is highest.
-Rutherford's gold foil experiment led to the 'solar system' model, ie 99.99% of everything is
nothing.
-Electrons can move into higher energy levels by excitation- through any form of energy
[ light, heat etc]. Energy is released on return to ground state.
-Orbitals are filled starting with the lowest energy level. if there happen to be two orbitals of
the same energy level, both will receive one electron before either receives a second.
-Periods= number of electron shells [levels]
-Increasing atomic number corresponds to progressive filling of electron shells.
, The position of an electron can be defined by 4 quantum numbers;
1. Principal quantum number; any positive integer corresponding to the gross energy
levels.
2. Angular momentum quantum number; describes the shape of an orbital and that there
is a maximum number of electrons that can occupy a quantum group.
3. Magnetic quantum number; describes orbital shape relative to the nucleus.
4. Spin projection quantum number; describes the direction of spin of the electron ie
clockwise or anticlockwise.
The Periodic Table
The Periodic Table The Periodic Table and electronic configuration Properties Atomic radius
Ionisation energy Electron affinity Electronegativity and electropositivity Element Groups
Lecture Notes
Chemistry Skills
1. SI units
2. Scientific notation
3. Significant figures
4. Molar masses
5. Atomic masses and isotopes
6. Electron configurations
1. SI Units
all measurements must have units to describe what and how much is measured
The International System of Units is metric based and in decimal form
attaching prefixes to a base unit alters the magnitude [by a power of 10]
SI Units
Prefixes
1L = 1000mL = 1dm^3
1mL = 1cm^3
1000cm^3 = 1L
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