Transcript The Chemical Earth

The Chemical Earth
Why study chemistry?
The Earth includes the biosphere, lithosphere,
hydrosphere and atmosphere. Each of these is a
mixture of thousands of different substances, many
of which are useful to us if we:
 have an understanding of the properties of the
elements and compounds that make up the Earth’s
materials
 develop efficient processes for separating useful
materials
Classification of matter

Pure substances have a fixed composition and
fixed properties. They cannot be decomposed by
simple physical separation techniques.

Mixtures have variable composition and variable
properties. They can be separated into their
components by various physical separation
techniques.
Elements
Pure substances can be further classified into
elements and compounds.
 Elements are the simplest pure substances
consisting of only one type of atom. They
cannot be broken down (or decomposed).
Compounds

Compounds are also pure substances. They are
composed of two or more elements that are
chemically bonded together. They are composed
of a fixed number of atoms of each component
element. They can be decomposed into their
component elements or into simpler compounds.
Mixtures
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The particles of each component in
homogeneous mixtures are distributed
uniformly (eg: sugar dissolved in water).
The particles in heterogeneous mixtures are
not distributed uniformly (eg: concrete)
Chemical analysis
Chemical analysis is the process of finding out what
is present in a particular chemical sample.
Chemical analysis can be:
 Qualitative – determining what substances are
present
 Quantitative – determining how much of each
substance is present in a sample.
Gravimetric analysis
Gravimetric analysis involves separating the
components of a material and accurately
determining their mass. The percentage composition
of the material can then be calculated. Gravimetric
analysis can be used to determine the:
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composition of a mixture using physical separation
techniques
% composition of a compound using chemical and
physical separation techniques.
Elements
Elements consist of atoms of the same type. Many
elements exist in nature as molecules.
 Monatomic molecules: consist of only one atom
(eg: noble gases)
 Diatomic molecules: a molecule in which two
atoms are bonded together.
 Polyatomic molecules: a molecule of more than
two atoms bonded together
Reactivity of elements
Elements vary in their tendency to react. Elements
that react readily (eg: calcium, sodium) are usually
found combined with other elements as
compounds. On the other hand elements such as
gold which have low reactivity are often found in
their pure form.
Metals
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Relatively high densities
Good conductors of heat and electricity
Malleable and ductile
Shiny surface when freshly cut of cleaned
Relatively high melting points
Non-metals
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State and form is variable (oxygen is a gas,
bromine a liquid, sulfur a solid)
Usually not lustrous
Poor conductors heat and electricity
Not malleable or ductile
Variable melting points
Atomic theory
An atom is made up of three fundamental particles:
 Protons: positively charged particle
 Neutrons: neutral
 Electrons: negatively charged particle
In an uncharged atom:
number of protons = number of electrons
The nucleus
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Contains protons and neutrons.
Has a positive charge equal to the number of
protons
Contains about 99.9% of the mass of the atom
Extremely dense
Electrons move in space outside the nucleus
Atomic number and atomic mass
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Atomic number (Z): number of protons in
the nucleus and is fixed for any one element.
(NB: In an uncharged atom this will also be
the number of electrons)
Atomic mass (A): sum of the number of
protons and neutrons in the nucleus
Atomic number and atomic mass
Atomic mass = atomic number + number of neutrons
OR
A
=
Z
+
number of neutrons
-
Z
=
number of neutrons
Therefore:
A
Atomic number and atomic mass
A
Z
X
X is the element symbol
A is the atomic mass
Z is the atomic number
Electrons
Electrons are believed to exist in energy levels or
shells. The maximum number of electrons in each
shell is determined by the formula 2n2 (where n =
shell number 1, 2, 3, 4, etc). So:
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1st shell: can hold a maximum of 2 electrons
2nd shell: can hold a maximum of 8 electrons
3rd shell: can hold a maximum of 18 electrons
4th shell: can hold a maximum of 32 electrons
Electron configuration
The pattern of electrons in each shell is called the
electron configuration. When determining the
electron configuration of an atom the general rule is:
Starting from the innermost shell, each electron
shell or energy level must be filled before moving
to the next energy level or shell.
NB: potassium and calcium are exceptions
Sodium
23
Na
11
This means an atom of Sodium has 11 protons,
11electrons and and 12 neutrons.
Therefore electrons are arranged as:
 First shell = 2 electrons
 Second Shell = 8 electrons
 Third shell = 1 electron
Valence energy level
The outermost shell of an atom is referred to as
the valence energy level. Similarly, the electrons
that occupy the outermost shell are called
valence electrons.
In the periodic table elements with the same
number of valence electrons occur in the same
column or group.
Octet rule
In general, atoms are
most stable when they
have 8 electrons in their
outer-most shell. This
accounts for the lack of
reactivity of the noble
gases.
Ions
Elements can achieve stable electron configurations
by losing or gaining electrons. In doing so they form
ions.
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Positively charged ions (cations) are formed when
one or more electrons are removed from an atom.
Negatively charged ions (anions) are formed when
an atom gains one or more electrons
Electron dot diagrams
Electron dot diagrams are a way of showing the
arrangement of valence electrons in atoms. For
example:
1 valence
electron
2 valence
electrons
7 valence
electrons
6 valence
electrons
Electron dot diagrams and ionic
bonds
Electron dot diagrams can also be used to show
the formation of ions.
Question: why will these two ions be attracted to
each other and form an ionic bond?
Polyatomic ions
Polyatomic ions are groups of atoms bonded to
one another that have a net positive or negative
charge. The carbonate ion is an example of a
The superscript 2- indicates
polyatomic ion.
CO23
Polyatomic ions often
have the suffix “ate” or
“ite”.
that there are two more
electrons than the total
number of protons
possessed by the four atoms
The subscript indicates that
there are 3 oxygen atoms
Covalent bonds
Covalent bonds are formed when adjacent atoms
share electrons. For example, a chlorine atom has
the electron configuration (2, 8, 7). Two chlorine
atoms can combine to form a chlorine molecule Cl2
by sharing a pair of electrons (each atom
contributes one electron).
http://www.gcsescience.com/Chlorine-Formation.gif
Electron dot diagrams and
covalent bonds
Molecules are a group of two or more atoms held
together by covalent bonds.
Bonds in which two electrons are shared are called
single covalent bonds and can be represented by a
line drawn between the atoms H–H.
Electron dot diagrams and
covalent bonds
The number of covalent bonds formed by an atom
depends on the number of valence electrons.
In forming the molecule O2, each oxygen accepts a
share of two electrons from the other atom. Hence
four electrons are shared by the two oxygen atoms.
This is called a double covalent bond: O=O
Ionic or covalent?
If one member of a pair of atoms wants to gain
electrons while the other wants to lose electrons
then the pair will form an ionic bond. If both
members want to gain electrons then they will form
covalent bonds.
Question:
 Will Group I elements tend to form ionic bonds or
covalent bonds?
 What about Group 6 elements?
Valency
The valency of an element is a measure of its
“combining power” (the number of bonds it can form).
 When an element forms ionic compounds the
valency is the charge the atom carries.
Ex: Na+ = +1 valency
 When an element forms covalent compounds
valency is the number of covalent bonds the atom
forms.
Ex: water is H-O-H, the valency of O = 2 and H = 1
Formula ionic compounds
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Ionic compounds are electrically neutral –
therefore the number of negative charges must
equal the number of positive charges.
Ex: NaCl – the numbers of sodium and chloride
ions is equal.
If the charges on the ions are not equal then
there will be more ions with the smaller charge.
Ex: the compound formed between Ca2+ and Clis CaCl2 (there are 2 Cl ions for each Ca ion)
Naming ionic compounds
The following naming rules apply to ionic
compounds:
1.
The cation (positive ion) is named first
2.
The anion (negative ion) is named second
3.
The suffix ‘ide’ is added to the non-metal in
simple binary compounds (compounds made up
of only two elements)
Ex: NaCl = sodium chloride
Formula covalent molecular
compounds
In covalent compounds the formula represents the
number of atoms of each element in one molecule of
the compound. This is also called the molecular
formula.
Ex:
H2O – two atoms of hydrogen and one atom of
oxygen
Naming molecular compounds
1.
2.
3.
The name of the element closer to the bottom or
left-hand side of the periodic table is written first.
The the suffix ‘-ide’ is added to the end of the name
of the second element.
The number of atoms of each element is indicated
by the prefixes ‘mono-’, ‘di-’, ‘tri-’, ‘tetra-’, ‘penta-’ or
hexa-’, which stand for 1, 2, 3, 4, 5 and 6
respectively.
NB: prefix ‘mono-’ is not used for the first-named element.
Chemical equations
In a chemical reaction the arrangement of atoms
is changed to produce new substances but atoms
are neither destroyed or created eg: mass is
conserved.
Chemical reactions are represented by chemical
equations.
Writing chemical equations
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Reactants are on the left and products are on the
right eg:
magnesium + oxygen
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magnesium oxide
Coefficients written IN FRONT of formulas show
the number of particles of that substance eg:
2Mg = 2 atoms of magnesium
Physical state of reactants and products is shown
by (g), (l), (s) – gas, liquid, solid
Writing chemical equations
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The number of atoms of each element must be
the same on the left and right hand sides of the
equation eg:
2Mg + O2
2MgO
2 x Mg
2xO
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2 x Mg
2xO
The sum of the electrical charges on the left must
equal that on the right
Balancing equations
Balancing a chemical equation is done by
changing the coefficients in front of the formulas.
1.
Write the word equation
2.
Write the formula for all elements and
compounds present
3.
Alter the coefficients to balance the number
of each type of atom on both sides of the
equation
4.
Write in the physical states
Decomposition reactions
Decomposition is a chemical reaction in which a
compound is broken down into their constituent
elements or simpler compounds. This is achieved
by adding energy as:
 Heat (thermal decomposition)
 Light
 Electricity (electrolysis)
Synthesis reactions
Synthesis is the process of forming a compound
from its component elements or other compounds in
a laboratory. It leads to the formation of a more
complex substance. For example, ammonia can be
synthesised directly by combining nitrogen and
hydrogen gases at high temperatures and
pressures.
Bond energy
A chemical change generally involves the
absorption or release of greater quantities of
energy than a physical change.
Reason:
A chemical change involves the
breaking of chemical bonds
Covalent & ionic bonding
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In covalent bonds electrons are shared
between atoms. The energy required to
separate atoms joined by a covalent bond is
referred to as the bond energy.
Ionic bonds are formed by the electrostatic
attraction between oppositely charge ions. The
energy required to break ionic bonds is
referred to as the lattice energy.
Bonding and Physical Properties
All substances are made up of atoms, molecules or
ions. It is the organisation of these particles that
determines the physical properties of a
substance. Solids can be classified into four
groups on the basis of their physical properties:
 Ionic compounds
 Covalent molecular compounds
 Covalent network compounds
 Metallic substances