8.2 The Chemical Earth

1
8.2 The Chemical Earth

• Focus 1
• The living and non-living components of the Earth
  contain mixtures

2
Balancing Chemical Equations

• Write the unbalanced 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.
• Balance the equation.
• 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

3
Balancing Chemical Equations

• Try these examples
• 1)Mg O2 ? MgO
• 2)Zn HCl ? ZnCl2 H2
• 3)CaCO3 ? CaO CO2

4
Balancing Chemical Equations

• 1)2Mg O2 ? 2MgO Balanced
• 2)Zn 2HCl ? ZnCl2 H2 Balanced
• 3)CaCO3 ? CaO CO2 Balanced

5
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

6
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/physicalgeog
  raphy/a/fourspheres.htm

7
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
8
Separation of Mixtures

• Sieve
• To separate solids of different sizes

• Filtration
• To separate solids and liquids/solutions

9
Separation of Mixtures

• Evaporation (to dryness)
• To separate dissolved solids in liquids

• Distillation
• To separate liquids from solutions (purification)

10
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)

11
Separation of Mixtures

• Liquification and fractional distillation
• To separate mixtures of gases-gases 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.

12
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.

13
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
14
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-molecular-weight organic
  compounds are ideal for this sort of analysis.
  The diagram on the right shows a typical
  chromatogram.

15
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
16
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)
17
Chemical Analysis

• Two general types
• Qualitative Analysis
• to determine what substances are present in a
  sample
• Quantitative Analysis
• to determine how much of each substance there is
  in a sample

18
Percentage composition

• Quantitative Analysis of a substance involves the
  determination of actual percentages present in a
  sample.
• This involves either
• Volumetric analysis-involves measuring
  percentages by volume.
• Gravimetric analysis-involves measuring
  percentages by mass/weight.
• In either case, the calculations will be similar

19
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.

20
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

21
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.

22
8.2 The Chemical Earth

• Focus 2
• Although most elements are found in combinations
  on Earth, some elements are found uncombined

23
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

24
The Periodic Table
http//library.tedankara.k12.tr/chemistry/vol1/ato
mstr/trans50.jpg
25
http//www.dayah.com/periodic/Images/periodic20ta
ble.png
26
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.

27
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.

28
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 semi-conductors in transistors and computer
  chips. They can be mixed with other elements
  (e.g. As and B) to increase their conductivities.

29
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.
30
8.2 The Chemical Earth

• Focus 3
• Elements in Earth materials are present mostly as
  compounds because of interactions at the atomic
  level

31
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.

32
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.

33
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
34
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

35
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 mass Relative charge
electron (e) 1/2000 -1
proton (p) 1 1
neutron (n) 1 0
36
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 p n e
Hydrogen 1 0 1
Deuterium 1 1 1
Tritium 1 2 1
Carbon 12 6 6 6
Carbon 13 6 7 6
Carbon 14 6 8 6
Uranium 235 92 143 92
Uranium 238 92 146 92
37
Structure of the Atom

• The Bohr Model
• Bohrs model of the atom consists of electrons in
  distinct energy levels or shells. The shells
  closest to the nucleus are the lowest energy
  (n1) and fill first.
• The maximum number of electrons in each shell can
  be calculated by 2n2.
• Therefore,
• n1 maximum of 2 e
• n2 maximum of 8 e
• n3 maximum of 18 e
• and so on
• The valence shell or outer shell can hold a
  maximum of 8.

38
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
39
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
40
Ions loss or gain of e-

• An atom that loses or gains electrons is called
  an ion.
• There are two types
• Cations () have lost electrons, making them
  positively charged (eg Mg2 loss of 2e-)
• Anions (-) have gained electrons, making them
  negatively charged (eg O2- gain of 2e-)

41
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 (-)

42
Ionic bonding

• Electrostatic attraction between oppositely
  charged particles

• 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)

cation
anion
Example Mg2 Cl- ? MgCl2
43
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.

44
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 e- to 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.

45
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
46
Lewis Dot Structures

• Lewis dot structures are a way of representing
  the valence e- configuration 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).

47
8.2 The Chemical Earth

• Focus 4
• Energy is required to extract elements from their
  naturally occurring sources

48
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.

49
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!

50
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!

51
Physical vs. Chemical -Water

• Water- a physical change
• Boiling water is an example of relatively weak
  intermolecular forces (Hydrogen bonds) breaking.
• Energy required 44 kJ/mol

• Water-a chemical change
• Electrolysis of water involves the breaking of
  very strong covalent bonds between H and O atoms.
• Energy required 286 kJ/mol

52
Physical vs. Chemical -Water

• Electrolysis breaking covalent bonds

• Boiling breaking H bonds

H bond
53
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
54
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
55
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).

56
Synthesis ReactionsEnergy 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
57
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)

58
Decomposition and SynthesisEveryday 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.

59
8.2 The Chemical Earth

• Focus 5
• The properties of elements and compounds are
  determined by their bonding and structure

60
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
61
Properties of elements, compounds and mixtures
are very different

• Aluminium
• Physical properties
• M.P. 6600C
• Density 2.7 g/cm3
• Conductivity 37.8 x106 Sm-1 (20 C)
• Chemical properties
• 4Al(s) 3O2(g) ? 2Al2O3(s)
• 2Al(s) 6HCl(g) ? 2AlCl3(aq) 3H2(g)
• Oxygen
• Physical properties
• M.P. -2190C
• Density 0.0013 g/cm3
• Conductivity non-conductor

• Bauxite ore (Al2O3?xH2O)
• Physical properties
• M.P. 20450C
• Density 3.5-4 g/cm3
• Conductivity non-conductor
• Chemical properties
• Al2O3(s) 2NaOH 3H2O(l) ? 2NaAl(OH)4(aq)
  (Bayer process)
• 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.

62
Structure of metals

• Metals can be described as three-dimensional
  lattices of positive ions in a sea of delocalised
  electrons.

63
Covalent Bonding

• Covalent Molecular-strong bonds, weak
  intermolecular forces holding molecules together

• Covalent Network-covalent bonding lattice that
  extends indefinitely throughout the crystal

64
Ionic bonding

• No discreet molecules are formed in ionic bonding
  due to electrostatic forces holding the atoms
  together. They form a continuing 3D lattice.

65
Properties associated with bond typesMetallic
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.

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Properties associated with bond typesIonic
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.

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Properties associated with bond typesCovalent
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.

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Properties associated with bond typesCovalent
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.

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The Chemical Earth

• Compiled by Robert Slider (2006)
• Please share this resource with others