8.2 The Chemical Earth
Focus 1:
The living and non-living components
of the Earth contain mixtures
Balancing Chemical Equations

Write the unbalanced equation.




Balance the equation.




Chemical formulas of reactants are listed on the left-hand side of the equation.
Products are listed on the right-hand side of the equation.
Reactants and products are separated by putting an arrow between them to show the direction of the
reaction. Reactions at equilibrium will have arrows facing both directions.
Apply the Law of Conservation of Mass to get the same number of atoms of every element on each side
of the equation. Tip: Start by balancing an element that appears in only one reactant and product.
Once one element is balanced, proceed to balance another, and another, until all elements are balanced.
Balance chemical formulas by placing coefficients in front of them. Do not add subscripts, because this
will change the formulas.
Indicate the states of matter of the reactants and products.





Use (g) for gaseous substances.
Use (s) for solids.
Use (l) for liquids.
Use (aq) for species in solution in water.
Write the state of matter immediately following the formula of the substance it describes.
Source: http://chemistry.about.com
Balancing Chemical Equations
Try these examples:
1)Mg + O2  MgO
2)Zn + HCl  ZnCl2 + H2
3)CaCO3  CaO + CO2
Balancing Chemical Equations
1)2Mg + O2  2MgO Balanced
2)Zn + 2HCl  ZnCl2 + H2 Balanced
3)CaCO3  CaO + CO2 Balanced
Elements, Compounds and Mixtures
-Elements are made of one type of atom and cannot be broken
down into simpler substances. Examples: Iron(Fe),
Oxygen(O2)
-Compounds are pure, homogeneous substances that can be broken
down into simpler substances, are made of two or more elements
and always have elements in the same ratio by mass. Examples:
table salt (NaCl), pure water (H2O)
-Mixtures contain two or more pure substances that are sometimes
heterogeneous and can be separated by physical means such as
filtering, boiling or the use of a magnet. Examples: iron filings
in sand, sugar dissolved in water
The Spheres of the Earth
The names of the four spheres are derived from the Greek words for stone (litho), air (atmo), water (hydro), and life
(bio).
Lithosphere
The lithosphere is the solid, rocky crust covering entire planet. This crust is inorganic and is composed of minerals.
It covers the entire surface of the earth from the top of Mount Everest to the bottom of the Mariana Trench.
Hydrosphere
The hydrosphere is composed of all of the water on or near the earth. This includes the oceans, rivers, lakes, and
even the moisture in the air. Ninety-seven percent of the earth's water is in the oceans. The remaining three
percent is fresh water; three-quarters of the fresh water is solid and exists in ice sheets
Biosphere
The biosphere is composed of all living organisms. Plants, animals, and one-celled organisms are all part of the
biosphere. Most of the planet's life is found from three meters below the ground to thirty meters above it and
in the top 200 meters of the oceans and seas.
Atmosphere
The atmosphere is the body of air which surrounds our planet. Most of our atmosphere is located close to the
earth's surface where it is most dense. The air of our planet is 79% nitrogen and just under 21% oxygen; the
small amount remaining is composed of carbon dioxide and other gasses.
Source: http://geography.about.com/od/physicalgeography/a/fourspheres.htm
The Spheres of the Earth
Mixtures in the Lithosphere:
-Rocks-mixtures of silicates, metals and other minerals
-Sand-mixture of silicon dioxide and shells
-Soils-mixture of clays, metals, sand, decomposing
matter
-Mineral ores-oxides, sulfides, carbonates, sulfates and
chlorides of metals
-Coal, oil and natural gas-mixtures of carbon
compounds
Mixtures in the Hydrosphere:
-Sea water- mixture of water and various salts such as
sodium, magnesium and calcium chlorides, and
other halides and sulfates
-Ground water- mixture of water and dissolved
chlorides and sulfates and suspended minerals
-Dissolved gases- nitrogen, oxygen and carbon dioxide
Mixtures in the Biosphere:
-Blood-mixture of plasma, red and white cells
-Animals, plants, bacteria-contain mixtures of carbon
compounds (carbohydrates, proteins, fats and
vitamins)
-Water with dissolved minerals
-Dissolved gases-oxygen, nitrogen, and carbon dioxide
Mixtures in the Atmosphere:
-Mixture of gases- elements of nitrogen, oxygen, argon
and a small amount of other gaseous compounds
such as water, carbon dioxide, carbon monoxide,
sulfur dioxide and nitrogen dioxide
Separation of Mixtures
Sieve
To separate solids of
different sizes
Filtration
To separate solids and
liquids/solutions
Separation of Mixtures
Evaporation (to dryness)
To separate dissolved solids
in liquids
Distillation
To separate liquids from
solutions (purification)
Separation of Mixtures
Separating Funnel
To separate two immiscible
liquids and for solvent
extraction. This technique
makes use of a difference in
densities
Separation by solubility
To separate mixtures of solids.
 One solid is soluble in a
solvent and the others are
not
 The insoluble components
are removed by filtration
 Evaporation is used to
recover the pure dissolved
substance (solute)
Separation of Mixtures
Liquification and fractional
distillation
To separate mixtures of gasesgases are cooled to liquefy
them, followed by fractional
distillation. Fractional
distillation allows for
separation of substances with
similar boiling points.
Other methods to separate gases
would make use of
differences in solubility in
liquids such as water.
Separation of Mixtures
Chromatography
is the separation of mixtures by selective adsorption (absorbing onto the surface)
onto a stationary phase. This technique is used to sort a mixture out into its
separate components.
There are several types for various mixtures and they include:
 Column chromatography
 Paper chromatography
 Thin layer chromatography
 Gas chromatography (GC)
All techniques make use of an inert substance such as alumina, silica or paper.
The components of a mixture adhere to the inert substance with different
strengths, which leads to separation.
Separation of Mixtures
Paper chromatography








This is the simplest form of chromatography.
The stationary phase is a special chromatography paper,
but often filter paper is used in schools.
The mobile phase is a solvent mixture, e.g. water and
ethanol.
The mixture under analysis is placed in a tiny,
concentrated dot near the bottom of the paper.
The paper is hung with the bottom dipped in solvent,
which rises up the paper to come in contact with the
mixture.
As the solvent rises further up the paper, the components
are separated as they are swept along.
The strip of paper is called a chromatogram.
Identification of the components is based on Rf values – a
ratio between the distance travelled by the component to
the distance travelled by the solvent front.
Solvent Front
Starting line
Separation of Mixtures
Gas chromatography (GC) uses a stationary phase and a mobile phase. The mobile
phase is a carrier gas and the stationary phase may be a liquid or a solid. GC is a
very rapid, highly sensitive and reliable form of analysis, but is limited to
compounds that can be vaporised without decomposing. Low-molecularweight organic compounds are ideal for this sort of analysis. The diagram on
the right shows a typical chromatogram.
Separation of Mixtures-summary of techniques
Separation Method
Property used to achieve
separation
Sieving
Particle size
Filtration
One substance is solid, the other is liquid
or solution
Evaporation
Liquid has a much lower boiling point
than the solid
Distillation
Large difference in boiling point
Fractional Distillation
Smaller difference in boiling point
Separating Funnel
Density (m/vol) of immiscible liquids
Adding a solvent then filtration
One substance is soluble in a solvent and
the others are not
Chromatography
Different adsorption to a stationary
phase
Separation of Mixtures-examples
Separation Method
Example of use
Sieving
To separate sand from gravel at a rock quarry
Filtration
Drinking water purification processes
Evaporation
Salt evaporation ponds for table salt
Distillation
Obtaining pure water from sea water
Fractional Distillation
Separation of crude oil components (petrol,
diesel, kerosene, waxes, etc.)
Separating Funnel
To remove oil from water, solvent extraction in
analytical testing (e.g. pesticides)
Adding a solvent then filtration
Removal of salt from sand with water
Chromatography
Analytical testing (e.g. water contaminants)
Chemical Analysis
Two general types:
1.
Qualitative Analysis

2.
to determine what substances are present in a sample
Quantitative Analysis

to determine how much of each substance there is in a
sample
Percentage composition
Quantitative Analysis of a substance involves the
determination of actual percentages present in a
sample.
This involves either:
1.
Volumetric analysis-involves measuring percentages
by volume.
2.
Gravimetric analysis-involves measuring percentages
by mass/weight.
In either case, the calculations will be similar
Gravimetric Analysis
There are a variety of reasons for determining the % composition
of a substance in a mixture including:




Determining the amount of pollutants present in drinking water.
Determining the amount of a metal present in an ore sample.
Quality control in the production of a variety of consumer
goods. (e.g. ensuring the correct quantities of N, P, and K in
fertilisers)
Soil testing to determine suitability for plant/crop growth.
Gravimetric Analysis
Gravimetric analysis involves the use of a variety of separation techniques,
followed by a simple calculation to determine the percentage composition of
a substance.
For example:
A sample of ore weighing 10.63g is found to contain 1.55g of nickel (Ni) and
0.76g of cobalt (Co). Calculate the % composition of Ni and Co.
%component = mass of component in sample x 100
total mass of sample
%Ni = 1.55g/10.63g x 100 = 14.58%
%Co = 0.76g/10.63g x 100 = 7.15%
Class Assignment
Choose a mixture from one of the 4 spheres of the Earth and gather information
about the following:




Industrial separation processes to separate the mixture
The properties of the mixture that are used in these separation processes.
The products of separation and their uses
The issues associated with wastes generated from these processes.
Present your information in Report Style with supporting diagrams, and a source
list.
8.2 The Chemical Earth
Focus 2:
Although most elements are found in
combinations on Earth, some elements are
found uncombined
Properties of the Elements
Elements are classified into three categories based on their physical properties.
The 3 categories are:



Metals
Non-metals
Semi-metals or metaloids
Some of the physical properties used in this classification:





Density (mass/volume)
Boiling point/melting point
Electrical and Thermal conductivity
State at room temperature (solid, liquid or gas)
Appearance
The Periodic Table
http://library.tedankara.k12.tr/chemistry/vol1/atomstr/trans50.jpg
http://www.dayah.com/periodic/Images/periodic%20table.png
Properties of the Elements
Metals: (e.g. Fe, Cu, Mg, Al, Au)
 solid at room temperature (except Hg) and usually dense/hard.
 usually high melting/boiling points.
 have a shiny (lustrous) appearance.
 are malleable (able to be hammered into sheets).
 are ductile (able to be drawn into wires).
 are good conductors of heat and electricity.
Uses: construction materials, utensils, electrical wiring, household
appliances, drink cans, etc.
Properties of the Elements
Non-metals: (e.g. C, S, He, Cl)
 can be solid liquid or gas at room temperature.
 usually have relatively low melting/boiling points.
 are usually not lustrous.
 are usually brittle, not malleable or ductile.
 Are poor conductors of heat and electricity (except for C in the
form of graphite).
Uses: carbon used as an electrode in dry cells and is the “lead” in
pencils, sulfur used in vulcanising rubber, neon is used in “neon”
signs and chlorine is used in bleach and swimming pools as well
as in the production of plastics such as PVC.
Properties of the Elements
Semi-metals: (B, Si, Ge, As, Sb)
 have properties that are a combination of metal and non-metal
properties.
 usually have high melting/boiling points.
 have variable conductivities depending upon temperature, but
are usually low.
 have variable appearance.
Uses: mixtures of silicon and germanium are used as semiconductors in transistors and computer chips. They can be
mixed with other elements (e.g. As and B) to increase their
conductivities.
Reactivity of the Elements
The elements vary greatly in their reactivity. How reactive an element is directly
related to how the electrons are arranged in the atom influencing what form it
will take in nature.



Some elements are not very reactive and are therefore found uncombined in
nature. These include: the noble gases (He, Ne, Ar, Kr, Xe, Rn), and the
metals Au, Ag, Pt and Cu (sometimes).
Some elements occur as molecules that contain only one type of atom. These
are referred to as molecular elements. These are also found combined with
other elements in compounds. These include: O2, N2, H2, Cl2, I2, P4
Most of the elements are reactive and therefore occur as compounds in
nature. These include: NaCl, H2SO4, SiO2.
General rule: The more reactive an element is, the less
of a chance it will be found uncombined in nature.
8.2 The Chemical Earth
Focus 3:
Elements in Earth materials are present
mostly as compounds because of
interactions at the atomic level
The particle nature of matter

Matter is often described as being made up of small particles that are
continuously moving and interacting. In each of the three states of matter
(solid, liquid, gas) the particles experience vibrational motion. Liquids and
gases experience translational (movement) motion as well. Gases experience
more translational motion than liquids as they have more energy.
solid
liquid
gas
The particle nature of matter
The primary "particle" in chemistry is the atom. Atoms are
defined as the smallest particle of an element. However, you
probably know that there is a substructure to an atom; that it is
made of protons, neutrons and electrons. You may also know
that protons and neutrons are each made of three quarks.
The particle nature of matter
Each element has a distinctive
atomic number and mass
number.
The atomic number (Z)
corresponds to the number
of protons in the nucleus.
The mass number (A)
corresponds to the total
number of neutrons and
protons in the nucleus.
Mathematically:
A = Z + number of neutrons
Structure of the Atom
The particles that make up the elements are called atoms.
All atoms of one element are the same, but they are
different from the atoms of all other elements. In other
words, each element has a distinct type of atom with a
specific number of protons, neutrons and electrons.



Protons have a +ve charge
Electrons have a –ve charge
Neutrons have no charge
Structure of the Atom
Protons (p) and neutrons
(n) are found in the
centre of the atom in the
nucleus
Electrons (e) are found in
the surrounding space
around the nucleus
moving randomly in
what is known as an
‘electron cloud’.
Relative Relative
mass
charge
electron
1/2000
(e)
-1
proton
(p)
1
+1
neutron
(n)
1
0
Structure of the Atom
Isotopes
All atoms of the same element
have the same number of
protons in the nucleus,
however they do not
necessarily have the same
mass. These atoms differ in
the number of neutrons and
therefore, the mass number
and are known as isotopes.
Some well-known isotopes
are in the table to the right.
Name
Hydrogen
Deuterium
Tritium
Carbon 12
Carbon 13
Carbon 14
Uranium 235
Uranium 238
#p
1
1
1
6
6
6
92
92
#n
0
1
2
6
7
8
143
146
#e
1
1
1
6
6
6
92
92
Structure of the Atom
The Bohr Model
Bohr’s model of the atom consists of
electrons in distinct energy levels or
‘shells’. The shells closest to the
nucleus are the lowest energy (n=1)
and ‘fill’ first.
The maximum number of electrons in each
shell can be calculated by 2n2.
Therefore,
n=1 maximum of 2 e
n=2 maximum of 8 e
n=3 maximum of 18 e
and so on…
The valence shell or outer shell can hold a
maximum of 8.
Structure of the Atom
Orbitals
Schrödinger used quantum mechanics to describe the shape of the ‘clouds’ within
each energy level. These are called orbitals and each energy level contains an
increasing number of orbitals to accommodate more electrons. All energy
levels contain ‘s’ orbitals, which are spherical (one lobe). All but the first
energy level contain 3 ‘p’ orbitals, which are dumbbell shaped (two lobes).
After the first two, each energy level contains 5 ‘d’ orbitals, most of which
have 4 lobes. Higher energy levels contain 7 ‘f’ orbitals. Each orbital can
accommodate 2 electrons. Therefore:




‘s’ orbitals hold 2 electrons
‘p’ orbitals hold 6 electrons
‘d’ orbitals hold 10 electrons
‘f’ orbitals hold 14 electrons
Note: For Interest Only! You are not required to
learn this information for the HSC
http://webfac1.enmu.edu/longro/www/orbitals/atorb.htm
Structure of the Atom
Below is a representation of the relative energy levels of electron
orbitals and how they appear around the nucleus.
Note: For Interest Only! You are not required to
learn this information for the HSC
Ions – loss or gain of eAn atom that loses or gains electrons is
called an ion.
There are two types:
1. Cations (+): have lost electrons, making them
positively charged (eg Mg2+ loss of 2e-)
2. Anions (-): have gained electrons, making
them negatively charged (eg O2- gain of 2e-)
Ions
The loss or gain of e- to form ions is directly related to the number of valence e- in an
atom. All atoms have a driving force towards a noble gas e- configuration as this is the
most stable configuration (i.e. 8 e- in the valence shell, unless we are talking about the
1st shell which only holds 2 e- as in He).
We can predict the ions that are formed by atoms by using the Periodic Table. The group
number (column number) indicates the number of e- in the valence shell.
Therefore:
Group I has one valence e- and will tend to lose 1e- forming a +1 ion and
Group VII has 7 valence e- and will tend to gain 1e- forming a -1 ion, etc.
The transition metals are more difficult to predict as many of these elements have a
variable e- configuration, however, these will all lose electrons to form positive ions.
In general:
Metals tend to form cations (+) and non-metals tend to form anions (-)
Ionic bonding
Ionic bonds are formed from the
transfer of electrons from one
atom to another. As previously
stated, this is to obtain an overall
noble gas configuration. The ratio
of atoms results in an electrically
neutral compound.
Because oppositely charged particles
attract to form these bonds, ionic
bonds tend to form between
metals and non-metals.
Note: ionic compounds do not form
discreet molecules, rather they tend
to form an array of anions and
cations in a fixed ratio which is
given in the empirical formula.
(See next slide)
Electrostatic attraction between
oppositely charged particles:
+
cation
Example:
Mg2+ + Cl-  MgCl2
anion
Ionic bonding
No discreet molecules are formed in ionic bonding due to
electrostatic forces holding the atoms together. More
information about these and their properties in 8.2.5.
-
+
-
+
-
+
-
+
-
+
-
+
-
+
-
+
-
+
-
+
-
Covalent Bonding
Covalent bonds are formed between two atoms sharing electrons.

In covalent bonding, there is no electrostatic attraction as in ionic bonding.
Atoms will ‘share’ a pair (single bond) or pairs (double or triple bonds) of eto gain a noble gas configuration. For example:

Cl with and electron configuration of (2,8,7) will covalently bond with
another Cl of (2,8,7) or with H of (1) to form Cl2 or HCl.

In the examples of Cl2 and HCl, all atoms have a full valence shell due to the
sharing of electrons. Cl has 8 e- and H has 2 e-. These compounds then exist
as individual particles or molecules and are known as covalent molecular
substances to distinguish them from covalent lattices such as in silicon
dioxide and diamond.

Other examples include water, ammonia and carbon dioxide.
Covalent Bonding
Covalent bonding leads to the formation of discreet
molecules (i.e. single units that are often weakly bonded
together by intermolecular forces). More about these
and their properties in 8.2.5.
Water
Chlorine
Hydrogen
Chloride
Lewis Dot Structures
Lewis dot structures are a way
of representing the valence econfiguration of an atom and
show how valence e- are
arranged in compounds.
Lewis dot structures can be used
to show the formation of
ions but are more commonly
used to show covalent
bonding.
The compounds formed to the
right are methane, ammonia,
water and hydrogen chloride
(hydrochloric acid).
8.2 The Chemical Earth
Focus 4:
Energy is required to extract
elements from their naturally
occurring sources
Physical vs. Chemical

Physical – changes that are associated with physical
properties which do not change the chemical
composition of a substance.


E.g. hardness, density, malleability, ductility, electrical and
thermal conductivities, melting point, boiling point, solubility
Chemical – changes that occur when a substance breaks
down or reacts with another substance in a chemical
reaction

A new substance is always formed and has different
properties than the original reactants.
Physical Changes - examples




Changing of state (melting iron, boiling water)
Changing the physical appearance (crushing ore
in a ball mill, drawing copper into wires)
Dissolving a solid in a liquid (sugar into water)
Separation of mixtures (filtering sand from
water, separating sea salt from water)
Physical changes – no new substances!
Chemical Changes - indications




A gas is evolved (iron and HCl generate H2 gas)
A solid (precipitate) is formed when two solutions are
added together (silver nitrate and sodium chloride
solutions produce a white solid of silver chloride).
A change in colour (purple potassium permanganate
(KMnO4) is added to hydrogen peroxide, the solution
turns colourless).
Change of temperature (magnesium is burned in air and
becomes very hot)
Chemical change – at least one new substance!
Physical vs. Chemical -Water
Water- a physical change
Water-a chemical change
Boiling water is an example of
relatively weak intermolecular
forces (Hydrogen bonds) breaking.
Electrolysis of water involves the
breaking of very strong covalent
bonds between H and O atoms.
Energy required = 44 kJ/mol
Energy required = 286 kJ/mol
Physical vs. Chemical -Water

Boiling – breaking H bonds

Electrolysis – breaking
covalent bonds
H H
H H
+
+
O
H
O
H
H bond
O
Physical vs. Chemical -summary
Chemical change
(reaction)
Physical change
At least one new substance
is formed
No new substances are
formed
Difficult to reverse (e.g.
unboiled egg?)
Easily reversed (e.g. melt ice
and freeze it again)
Generally requires a large
amount of energy
Generally, relatively small
quantities of energy
Decomposition Reactions
Energy absorbed
When a compound decomposes into two or more
other pure substances, energy is normally
absorbed in the form of heat, light or electricity.
Heat
Light
Electricity
+ AB  A + B
Decomposition Reactions-examples
Heat
Solid copper nitrate decomposes to solid copper oxide, nitrogen dioxide
and oxygen gases.
Electricity
Molten lead bromide (4000C) forms bromine gas at the +ve electrode
and liquid lead at the –ve electrode.
Light
Solid silver chloride decomposes to silver metal and chlorine gas.
(decomposition of silver compounds is the basis of photography
development).
Synthesis Reactions
Energy released
A synthesis or combination reaction involves the
combination of two or more pure substances.
When a compound is formed from its elements,
it is known as a direct combination reaction.
These reactions normally release energy.
A + B  AB + energy
Synthesis reactions - examples
Magnesium burns in air (oxygen) to produce
magnesium oxide (light and heat energy released)
Hydrogen and oxygen combine in an explosive
reaction to produce water (much energy released)
Copper metal combines with yellow sulphur when
heated to produce copper (I) sulphide. (much heat
energy is released)
Decomposition and Synthesis
Everyday Applications
Decomposition



Air bags – sodium azide (NaN3)
decomposes to sodium and
nitrogen gas by ignition with a
detonating cap.
Limestone (primarily CaCO3) –
decomposes to calcium oxide and
carbon dioxide by heating to make
lime (CaO), cement and glass.
Aluminium - the industrial process
of electrolysing aluminium oxide
produces aluminium metal.
Synthesis



Rust – iron and oxygen combine in
the presence of water to form iron
(III) oxide.
Burning coke (primarily carbon) is
used as a fuel in smelting iron ore
in a blast furnace during the steel
making process.
Pollutants - NO (nitric oxide or
nitrogen monoxide) and NO2
(nitrogen dioxide) are formed
inside the combustion chambers of
cars from nitrogen and oxygen
gases.
8.2 The Chemical Earth
Focus 5:
The properties of elements and
compounds are determined by
their bonding and structure
Properties of elements and their
compounds are very different
8Fe(s) + S8(s) + heat  8FeS(s)
substance
colour
Melting
point (0C)
Boiling
point (0C)
Density
(g/cm3)
magnetic
Iron
Grey
1535
2750
7.9
yes
Sulphur
Yellow
113
445
2.1
no
Iron (II)
sulphide
Yellow-gold 1194
-
4.84
no
Properties of elements, compounds
and mixtures are very different
Aluminium:
Physical properties
 M.P. = 6600C
 Density = 2.7 g/cm3
 Conductivity = 37.8 x106 S·m-1 (20 °C)
Bauxite ore (Al2O3·xH2O)
Physical properties
 M.P. = 20450C
 Density = 3.5-4 g/cm3
 Conductivity = non-conductor
Chemical properties
 4Al(s)+ 3O2(g)  2Al2O3(s)

2Al(s) + 6HCl(g)  2AlCl3(aq) +
3H2(g)
Chemical properties
 Al2O3(s) + 2NaOH + 3H2O(l) 
2Na[Al(OH)4](aq) (Bayer process)
Oxygen:
NB: Bauxite is essentially an impure
aluminium oxide. The major impurities
include iron oxides, silicon dioxide and
titanium dioxide. The impurities remain
as solids and do not react with NaOH.
This process removes these impurities.
Physical properties
 M.P. = -2190C
 Density = 0.0013 g/cm3
 Conductivity = non-conductor
Structure of metals
Metals can be described as three-dimensional lattices of
positive ions in a sea of delocalised electrons.
Covalent Bonding

Covalent Molecular-strong bonds,
weak intermolecular forces holding
molecules together

Covalent Network-covalent
bonding lattice that extends
indefinitely throughout the crystal
Ionic bonding
No discreet molecules are formed in ionic bonding
due to electrostatic forces holding the atoms
together. They form a continuing 3D lattice.
-
+
-
+
-
+
-
+
-
+
-
+
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Properties associated with bond types
Metallic bonding

High melting points-due to strong attraction between positively
charged metal ions and delocalized electrons. The higher the
valency, the stronger the bond e.g. Ca2+ is stronger than K+.

Good conductors of heat and electricity-due to the high mobility
of delocalized electrons. Electrons enter and leave a metal easily.

Malleable and ductile-due to delocalized electrons not belonging
to any particular metal atom. Therefore, one layer of ions can
slide over another without disrupting the bond between metal
atoms. The electrons and metal ions simply rearrange.

Hardness- tend to be hard due to tightly packed atoms.
Properties associated with bond types
Ionic bonding

High melting points-due to strong electrostatic attraction between anions and
cations.

Non-conductors of electricity in solid state-due to oppositely charged
particles, which are in fixed positions.

Conductors in the liquid (molten) state-due to the ions being able to move
freely through the liquid.

Hardness-due to strong electrostatic attraction between oppositely charged
particles.

Brittle-due to the fixed location of oppositely charged particles.
Displacement of ions moves them closer to ions of a similar charge, which
increases the repulsive forces along the fracture.
Properties associated with bond types
Covalent molecular bonding

Low melting points-due to generally weak attractive
forces between molecules. There are exceptions to this
rule (e.g. I2 melts at 1140C, but decomposes at 10000C).

Non-conductors-due to lack of mobile charged species
or delocalized electrons.

Soft-due to weak forces existing between molecules.
Properties associated with bond types
Covalent Network bonding

Very high melting points and boiling points-due to
strong covalent bonding, which form rigid 3-D
structures.

Non-conductors-due to lack of mobile charged species
or delocalized electrons.

Extremely hard- due to strong covalent bonding, which
form rigid 3-D structures.
The Chemical Earth
Compiled by: Robert Slider (2006)
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