THE CHEMICAL EARTH 1. THE LIVING AND NON-LIVING COMPONENT OF THE EARTH CONTAIN MIXTURES 1.1 Construct and balance equations of chemical reactions Text book page 7-73 1.2 Identify the difference between elements, compounds and mixtures in terms of the particle theory. Homogeneous – uniform composition throughout e.g. pure water, sugar, aluminium, petrol Heterogeneous – non-uniform/variable composition throughout e.g. strawberry jam, wood A mixture: Matter Pure substances – constant composition Mixtures – variable composition Elements – not separable into smaller substances Compounds – two or more elements chemically combined in fixed proportions Solutions – homogeneous mixtures (uniform composition and properties throughout Solutions – heterogeneo us mixtures – variable composition and properties throughout
Transcript
THE CHEMICAL EARTH
1. THE LIVING AND NON-LIVING COMPONENT OF THE EARTH CONTAIN MIXTURES
1.1 Construct and balance equations of chemical reactions
Text book page 7-73
1.2 Identify the difference between elements, compounds and mixtures in terms of the particle theory.
Homogeneous – uniform composition throughoute.g. pure water, sugar, aluminium, petrol
Can usually be separated into 2 or more pure substances Can be homogeneous or heterogeneous Properties change due to variable composition E.g. sea water, air, coffee
A pure substance:
Matter
Pure substances – constant composition Mixtures – variable composition
Elements – not separable into smaller substances
Compounds – two or more elements chemically combined in fixed proportions
Solutions – homogeneous mixtures (uniform composition and properties throughout
Solutions – heterogeneous mixtures – variable composition and properties throughout
Some can be decomposed into simpler substances, others cannot Is homogeneous (e.g. crystals of sugar) Has properties such as appearance, density, colour, melting and boiling points which are constant
throughout the whole sample Properties do not change Fixed composition E.g. table salt, sugar, copper, aluminium, diamond, gold
An element is a pure substance that cannot be decomposed into simpler substances
A compound is a pure substance which can be decomposed into simpler substancese.g. water, table salt (sodium chloride), sugar, alcohol
- Made up of 2 or more elements- Always has the elements present in the same ration by mass
Changes of state
- Solid liquid (melting)- Liquid gas (vaporisation, evaporation, boiling)- Gas liquid (condensation, liquidation)- Liquid solid (freezing, solidification)- Solid gas (sublimation)
Properties used to identify pure substances
Colour Physical state Melting/boiling points Density Electrical conductivity Solubility in different liquids Mechanical properties
Density = mass/volumemass per unit volumemeasured in g/mL
1.3 Identify that the biosphere, lithosphere, hydrosphere and atmosphere contain mixtures of elements and compounds
Hydrosphere – consists of different mixtures, compound water e.g. rivers, lakes, fresh water, sea water, ground water
Lithosphere – diverse range of mixtures. Rocks, sand, soils, mineral ores, coal, oil, natural gasmixtures contain predominantly compounds
1.4 Identify and describe procedures that can be used to separate naturally occurring mixtures of:- solids of different sizes-solids and liquids-dissolved solids in liquids-liquids-gases
Separation of solids of different sizes: sieving
Solids and liquids: filtration
Liquid/solution that passes through filter paper is the filtrate Sedimentation if the process in which solids settle to the bottom of container Decantation is process of pouring off the liquid and leaving the solid undisturbed at the bottom of container
Dissolved solids in liquids
Vaporizing off the liquid (evaporating, boiling) Only keep solid
Distillation
Separating two or more liquids from one another or separating the liquids from solids Liquids of sufficiently different boiling points Process in which a solution or mixture of liquids is boiled with the vapour formed being condensed back to a
liquid in a different part of apparatus and so separating from the mixture Liquid changes to vapour, rises up neck of flask and diffuse down the side arm and into the water-cooled
condenser, where the vapour is cooled and condensed back to a liquid, which is collected in the beaker Distillate: liquid collected (one with lower boiling point)
Volatile – able to be converted to vapourmore volatile of two liquids is the one with the lower boiling point
Fractional distillation
Separate liquids with boiling points fairly close together Expensive Has a fractionating column Repeated condensations and vaporization up the column, effectively giving many separate distillations E.g. crude oil
Immiscible liquids
Immiscible liquids – do not mix, do not form homogeneous liquid, two layers E.g. water and oil Separating funnel More dense liquid comes out
Separation based on solubility
If one solid is soluble in a particular solvent while the others are not
Sufficient solvent is added to the mixture to dissolve the soluble component; then the insoluble component/s are filtered off
The soluble liquid is recovered by evaporating the filtrate to dryness
Separating gases
Property used to separate: Differences in boiling points or differences in solubilities in liquids (water) Fractional distillation: if gases are similar boiling points
Separation method Property used in the separationSieving Particle sizeVaporisation (evaporation, boiling) Liquid has a much lower boiling point that the solidDistillation Big difference in boiling pointsFractional distillation Significant but small difference in boiling pointsFiltration One substance is a solid, the other a liquid or solutionAdding a solvent, then filtration One substance is soluble in the chosen solvent, while the
others are notSeparating funnel Immiscible liquids
1.5 Assess separation techniques for their suitability in separating examples of earth materials, identifying the differences in properties which enable these separations
Ex 11-12 (CC)
1.6 Describe situations in which gravimetric analysis supplies useful data for chemists and other scientists
Chemical Analysis
Qualitative Analysis – what substances are present?
Quantitative Analysis – How much of each substance is present? Percentage composition
Gravimetric Analysis – involves weighing
Volumetric analysis – involves measuring the volumes of solutions
Reasons for gravimetric analysis
Determine composition of soil in a particular location to see if it is suitable for growing a certain crop To determine the amount of particular substances in water/air to decide how polluted the samples are
E.g. A team of geologists discovered a new mineral in a remote desert location; it was a mixture of barium sulfate and magnesium sulfate. Its composition was determined as follows. They first ground up a 3.61g sample with water, magnesium sulfate dissolves, barium sulfate does not. The barium sulfate was filtered off, dried and its mass determined to be 1.52g. They evaporated the filtrate to dryness to recover the magnesium sulfate and determined its mass to be 2.07g. Calculate the percentage composition of the sample.
Percentage of barium sulfate = mass of barium sulfate present / total mass of sample x 100= 1.52 / 3.61 x 100=42%
Percentage of magnesium sulfate = 100-42 = 58%
1.7 Apply systematic naming of inorganic compounds as they are introduced in the laboratory
1.8. Identify IUPAC names for carbon compounds as they are encountered
2. ALTHOUGH MOST ELEMENTS ARE FOUND IN COMBINATIONS ON EARTH, SOME ELEMENTS ARE FOUND UNCOMBINED
2.1 Explain the relationship between the reactivity of an element and the likelihood of it existing as an uncombined element
Most elements are chemically reactive – when they come into contact with certain other elements they react to from compounds
The more reactive and element, the less change of finding it uncombined Sodium, potassium, calcium, magnesium, fluorine and chlorine are very reactive elements Elements that occur uncombined are gold, silver, platinum, sulphur, and noble gases – argon, helium
2.2 Classify elements as metals, non-metals and semi-metals according to their physical properties
Metals
Are solids at room temperature (except mercury) Have a shiny or lustrous appearance Are good conductors of heat and electricity Are malleable (able to be rolled into sheets) and ductile (able to be drawn into wires) E.g. aluminium cobalt, copper, gold, iron, lead, magnesium, nickel, potassium, sodium, silver, tin, zinc (LHS of
periodic table)
Non metals
E.g. argon, bromine, chlorine, hydrogen, iodine, nitrogen, oxygen
****Mercury has a shiny appearance and is a good conductor of electricity but is a liquid
****Carbon in the form of graphite is a fair conductor of electricity and it is a solid but resembles the non-metals more than the metals
Semi-metals (metalloids)– e.g. boron, silicon, germanium, arsenic, antimony, tellurium
2.3 Account for the uses of metals and non-metals in terms of their physical properties
Physical properties and uses of elements
Melting point, density, electrical conductivity, hardness and tensile strength Uses of iron, aluminium, copper, lead – building materials (cars, planes, machinery, electrical wiring,
household goods) Aluminium – used for making aircraft because of low density and adequate mechanical strength Iron – high tensile strength, used in cars and trains, moderately high density Copper – electrical wiring, high electrical conductivity Tungsten – filaments in electric light bulbs because of its high melting point CHEMICAL REACTIVITY AND COST ARE IMPORTANT FACTORS TO CONSIDER Carbon as graphite – significant electrical conductivity, dry lubricant, slipper/soft nature electrodes in
common dry cells (batteries) Carbon as diamond – jewellery, extremely hard, high refractive index Liquid nitrogen – cooling agent, suitability of its freezing and boiling points
2.6 Process information from secondary sources and use a periodic table to present information about the classification of elements as: - metals/non-metals/semi-metals, solids/liquids/gases at 25 degrees and normal atmospheric pressure
Similar properties in each vertical column (groups) Transition elements, middle of table Horizontal rows are called periods There is a gradual change in properties as we go across any one period Non-metals occur near the top right of the table Semi-metals : diagonal band separating metals from non-metals (boron, silicon, germanium, As, Sb, Te) Two liquids at room temperature: mercury, bromine 11 gases, rest are solids
*****Group headings indicate how many electrons are in the outer shell of each element.
3. ELEMENTS IN EARTH MATERIALS ARE PRESENT MOSTLY AS COMPOUNDS BECAUSE OF INTERACTIONS AT THE ATOMIC LEVEL
3.1 Identify that matter is made up of particles that are continuously moving and interacting
Particle nature of matter
All matter is made up of small particles Solids – particles are packed closely toget
Solid
Definite volume and shape Difficult to compress Low kinetic energy Particles vibrate at same spot Strong forces between particles Greater densities than liquids
Liquid
Some rotational/translational movement Takes shape of container Difficult to compress Moderate kinetic energy
Gas
Particles are spread out Expand to fill volume available – rapid translational movement Takes shape of container Easily compressed High kinetic energy No significant forces between particles
3.2 Describe qualitatively the energy levels of electrons in atoms
Energy increases as number of shell increase.
Maximum electrons in each shell level – 2, 8, 18, 32, 50
The arrangement of electrons in energy levels is called the electron configuration of the atom.
SEQUENCE FOR PUTTING ELECTRONS IN ENERGY LEVELS/SHELLS
DRAW DIAGRAM HERE
3.3 Describe the atoms in terms of mass number and atomic number
Atomic model
Atomic number – number of protons Atomic mass – number of protons + neutrons Number of electrons in outer shell – column number Number of shells- rows down (periods)
An atom is the smallest particle of an element which is still recognizable as that element.
An atom consists of an extremely small dense nucleus or core which contains the bulk of the mass of the atom and carries positive electrical charges.
The nucleus is surrounded by a cloud of rapidly moving extremely light particles carrying negative charges – electrons.
The amount of negative charge carried by these rapidly/randomly moving electrons is equal to the amount of positive charge on the nucleus so that the atom is neutral overall.
Protons are small positively charge particles.
Neutrons are small neutral particles
3.4 Describe the formation of ions in terms of atoms gaining or losing electrons
Stable electron configurations
Noble gases undergo no chemical reactions – completely filled or semi-filled energy levels. They have extremely stable electronic configurations.
Alkali metals are very reactive 0 have one more electron in the next outer shell than the nearby noble gases Alkali metals tend to lose one electron to obtain the electron configuration of the nearby noble gas
e.g. sodium loses on electron to become like neon Halogens gain one electron to achieve the electron configuration of the nearby noble gas (ATOMS TEND TO LOSE OR GAIN ELECTRONS IN ORDER TO BECOME LIKE THE NEARBY NOBLE GASES)
Valence electrons
The electrons in the incompletely filled highest energy level (outermost shell) are called valence electrons Outermost energy level is called the valence shell Noble gases have no valence electrons
3.5 Apply the Periodic Table to predict the ions formed by atoms of metals and non-metals
Formation of Ions
Ionic bonding is the outright transfer of electrons from one atom to another to from what are called ions – positively or negatively charged particles
E.g. sodium and chlorine combine to from the compound sodium chloride.Sodium loses one atom to become like neon, chlorine gains one atom to become like argon. One electron is transferred from a sodium atom to a chlorine atom. When the neutral sodium atom loses one electron it becomes positively charged positive ion – cationNa becomes Na+, Cl becomes Cl-
Positive ion – cation, negative ion – anion Strong electrostatic attraction between positive and negative ions Between metal and non-metal Metals generally from positive ions (cations), non-metals generally from negative ions (anions)
SIMPLE IONS – electrically charged species formed when atoms gain or lose electrons
POLYTOMIC IONS – electrically charged groups of atoms
3.6 Apply Lewis electron dot structure to:- the formation of ions-the electron sharing in some simple molecules
3.7 Describe the formation of ionic compounds in terms of the attraction of ions of opposite charges
If one atom wants to gain electrons while the other wants to lose, the compound will be ionic. If both want to gain electrons, then the compound will be covalent.
E.G.
3.8 Describe molecules as particles which can move independently of each other3.9 Distinguish between molecules containing one atom (the noble gases) and molecules with more than one atom.
Molecules:
Compounds – substance composed of two or more different types of atoms. Consist of two or more elements combined together in definite proportions by mass.
Molecule – two or more atoms bound together, capable of existing alone; smallest particle of a substance that is capable of separate existence.
E.g. hydrogen and oxygen from the compound water. Water is made up of molecules. Each molecule of water has two atoms of hydrogen and one atom of oxygen. All molecules of water are identical.
Molecule of elements
Oxygen exists as pairs. A diatomic molecule, 2 oxygen atoms chemically bonded together. This goes for all gas elements except for noble gases.
These molecules can be broken into separate atoms in chemical reactions. Noble gases e.g. helium, neon, argon can exist as independent atoms. Noble gases consist of monatomic
molecules.
3.10 Describe the formation of covalent molecules in terms of sharing electrons
Covalent bonding
Formed between pairs of atoms by the atoms sharing electrons. Between non-metal and non-metal Covalent bonds are as strong as ionic bonds Covalent bonds are a strong electrostatic force between the shared pair of electrons and the nucleus of
the atoms. Anything with H – hydrogen, is covalently bonded
E.g. 1 -2 chlorine atoms combine to form a chlorine molecule. They will share a pair of electrons. E.g. 2 Hydrogen and chlorine combine to from hydrogen chloride.
Covalent/simple molecular
Between atom covalent bond Strong electrostatic force between the nucleus and shared pair of electrons Low boiling and melting point Change of state involves breaking the weak intermolecular forces E.g. hydrogen gas, carbon dioxide, water Most molecular covalent dissolve in water, except for acids (e.g HCl)
Covalent network
E.g. graphite, diamond, quartz No intermolecular forces Only strong covalent bonds exist
High boiling and melting point >2000 degrees
Covalency and the Periodic Table
Covalent bonding occurs when both of the elements forming the compound need to gain electrons to attain noble gas configurations
Elements in the centre and to the right tend to form covalent bonds Number of covalent bonds is the number of electrons that an atom needs to gain to acquire noble gas
configuration Position of an element on the Periodic Table tells us how many electrons it needs to gain to achieve noble
gas configuration
3.11 Construct formulae for compounds formed from:-ions-atoms sharing electrons
Formulae for ionic compounds
We can deduce the charge on the ions they form from their position in the Periodic table Criss cross method
e.g.
Formulae for Covalent compounds
3.13 Construct ionic equations showing metal and non-metal atoms forming ions
Half equations
4. ENERGY IS REQUIRED TO EXTRACT ELEMENTS FROM THEIR NATURALLY OCCURRING SOURCES
4.1 Identify the differences between physical and chemical change in terms of rearrangement of particles
A change in which no new substance is formed is called a physical change
Physical changes
Changing the state of a substance Changing physical appearance Dissolving a solid in a liquid Separating mixtures
A change in which at least one new substance is formed is called a chemical change
e.g.
Heating green copper carbonate to form a black solid and colourless gas Burning silvery magnesium ribbon to form a white powder
Signs of a chemical change
Gas is evolved A precipitate if formed Change in colour Significant change in temperature Disappearance of a solid Odour is produced
In a chemical change the starting substances are called reactants.Substances that are formed are called products.
In all chemical changes-mass is conserved (balance equations)-the number of atoms of each type is conserved.
The law conservation of mass/matter: Matter can be neither created nor destroyed, but merely changed from one form to another.
4.2 Summarise the differences between the boiling and electrolysis of water as an example of the difference between chemical and physical change.
Electrolysis of water (decomposition)
Two 10ml measuring cylinders filled with water are inverted over a pair of inert electrodes in a beaker of water containing sulphuric acid (acid is necessary because water is a poor conductor of electricity)
The electrodes are connected to a voltage source and a current is allowed to flow
After several hours we can see a colourless gas has collected in each cylinder with the volume above the negative electrode being twice that above the positive electrode.
Pop test shows that the larger volume of gas is hydrogen, and splint test shows that the other gas is oxygen. Shows a chemical change.
Differences between boiling and electrolysis of water
Electrolysis produces two new substances (hydrogen and oxygen gas), boiling does not produce any new substance – just converts liquid water to gaseous water
Electrolysis is difficult to reverse, boiling is easily reversed Electrolysis requires much more energy that boiling
Boiling of water Decomposition of waterBonds broken Weak intermolecular forces Covalent bondsAmount of energy required Less MoreChemical reaction No YesPhysical reaction Yes NoNew substance formed No Yes
Breaking bonds requires energymaking bonds releases energy
4.3 Identify light, heat and electricity as the common forms of energy that may be released or absorbed during the decomposition or synthesis of substances and identify examples of these changes occurring in everyday life.
Heating copper nitrate
Copper nitrate is a blue solid and is soluble in water When it is heated in a crucible, a brown gas is evolved and a black solid remains The black solid is insoluble in water (different substance from copper nitrate) Copper nitrate has be decomposed into 2 substances – a brown gas (nitrogen dioxide) and a black solid
(copper oxide) , and a third substance A chemical change has occurred
Decomposing substances with electricity
Electrolysis of molten lead bromide
Lead bromide is a white crystalline solid Is a pure substance, melts at 373 degrees without decomposing If an electric current is passed though molten lead bromide at about 400 degrees, a choking brown vapour is
formed at the positive electrode and drops of silvery liquid at the negative electrode. Lead bromide has been decomposed into molten lead (silvery liquid0 and brown bromine gas A chemical reaction has occurred
Decomposing pure substances with light
Some pure substances can be decomposed by sunlight Silver chloride is a white solid, when exposed to sunlight it turns purple, then black After several hours of exposure, the sample has a smaller mass than the starting material Sunlight has decomposed silver chloride into a black solid and a gas (chlorine) Decomposition of silver compounds by light is the basis of photography
Direct combination reactions
A compound is formed from its elements Elements combine directly to form compounds and energy is released E.g. hydrogen and chlorine gases react vigorously to form hydrogen chloride gas, HCl Silvery magnesium burns in air (oxygen) to form a white powder, magnesium oxide MgO. This reaction
liberates much light energy and also releases a lot of heat.
Everyday applications
Decomposition reactions
In air bags in motor cars, sodium azide is decomposed to sodium and nitrogen gas Calcium carbonate (limestone) is decomposed to calcium oxide and carbon dioxide by heating it to make
lime, cement and glass Aluminium is extracted commercially by electrolysing molten aluminium oxide
Direct combination reactions
Rusting of iron and steel to from iron (III) oxide – from direct combination of iron and oxygen Burning of coke (carbon) – releases heat energy that can be used Lightning – creates a localised high temperature that nitrogen and oxygen gases combine to from gaseous
nitric oxide, NO.
4.4 Explain that the amount of energy needed to separate atoms in a compound is an indication of the strength of the attraction, or bond between them.
Explanation for energy changes
Decomposing a compound into atoms requires a large input of energy because it is necessary to overcome the strong chemical bonds holding the atoms together in compounds
There are strong electrostatic attractions (ionic bonds) holding ions together in ionic compounds and strong covalent bonds holding atoms together in covalent molecules and in covalent lattices
The stronger the chemical bonding in a compound, the more energy that is required to break the compound into atoms
4.7 Analyse and present information to model the boiling of water and the electrolysis of water tracing the movements of and changes in arrangements of molecules
In terms of particles
Boiling water does not alter the particles/molecules: it just separates them from one another: the water vapour contains the same water molecules as the liquid
Electrolysis breaks the particles water molecules are broken up and hydrogen and oxygen molecules are formed
Physical changes just rearrange the particles without changing their nature Chemical changes break up particles (molecules) and rearrange the atoms into new substances Conservation of mass
5. THE PROPERTIES OF ELEMENTS AND COMPOUNDS ARE DETERMINE BY THEIR BONDING AND STRUCTURE
5.1 Identify differences between physical and chemical properties of elements, compounds and mixtures
Physical properties
Melting and boiling points Density Appearance Electrical conductivity Heat conductivity Hardness
Chemical properties
Ease of decomposition by heat Effect of light Reactivity with other substances such as oxygen, chlorine and sulphur
5.2 Describe the physical properties used to classify compounds as ionic or covalent molecular of covalent network
Properties of covalent molecular and ionic substances
Ionic substances are always compounds. Covalent molecular substances can be elements as well as compounds.
Ionic substances Covalent molecular substancesSolids at room temperature At room temperature are generally gases (e.g. nitrogen
gas), or liquids (e.g. methanol) or solidsHigh melting points >400 degrees. High boiling points >1000 degrees.
Hard and brittle When solid they are softAs solids do not conduct electricity Pure covalent substances do not conduct electricity
either as solids or as liquidsWhen molten or in aqueous solution they do conduct electricity
In aqueous solution do not conduct electricity (unless they actually react with water to form ions)
5.3 Distinguish between metallic, ionic and covalent bonds
Covalent molecular
Held together by covalent bonds (atoms sharing pairs of electrons)
Low melting and boiling points; many are liquids or gases at room temperature Non-conductors of electricity in both the solid and liquid states Form solids with waxy appearance
Covalent network (covalent lattice)
Include elements such as carbon and silicon. 3D network of strong covalent bonds between the atoms which hold the lattice together.
Very high melting and boiling points Non-conductors of electricity in the solid and liquid states (no delocalized electrons) Extremely hard and brittle Chemically inert Insoluble in water and most other solvents
Ionic compounds
Oppositely charge ions held together by electrostatic attraction and arranged in regular 3D lattices. Metal combined with non-metal.
Hard and brittle
Solid
Melting Point
High
Not covalent
molecular substance
Conducitivty of solid
Conducting
METALLIC SUBSTANCE
Non-conuducting
Conductivity of molten state
Conducting
IONIC SUBSTANCE
Non-conducting
COVALENT NETWORK SUBSTANCE
Low
COVALENT MOLECULAR SUBSTANCE
does not conduct in solid and molten state
Non-conductors of electricity in the solid state. Good conductor of electricity when molten or in aqueous solution.
High melting and boiling points.
Metals
Relatively high densities Good conductors of heat/electricity Malleable (can be beaten into sheets) and ductile (can be drawn into wires) Shiny surface (lustrous) Relatively high melting points
Element Bonding NatureNon-metals Covalent Sharing of electrons
Bonding is to strong electrostatic attraction between the positive ions and the shared pair of electrons
Metals and non-metals Ionic Transfer of electrons and formation of cations and anions.Bonding is due to strong electrostatic attraction between the positive and negative ions
Metals Metallic Bonding is due to strong electrostatic attraction between the positive ions and the delocalized electrons.
5.4 Describe metals as three dimensional lattices of ions in a sea of electrons
Metallic bonding
Metals with the exception of mercury are solids at room temperature Relatively high melting points, fairly hard Good conductors of electricity Consist of an orderly 3D array of positive ions held together by a mobile ‘sea’ of delocalised electrons Valence electrons break away from their atoms, leaving behind positive ions Free electrons (delocalised electrons) no long belong to particular atoms, they move randomly through the
lattice, are shared by numerous positive ions and provide the chemical bonding that holds the crystal together
Delocalised electrons moving freely through the lattice allow metals to be good conductors of electricity. Electric current through a metallic wire is a flow of electrons Metals can bent, are ductile, and malleable – possible because when the orderly array of positive ions is
sheared , the mobile electrons are able to adjust to the new arrangement of positive ions and hold the whole assemble together
5.5 Describe ionic compounds in terms of repeating three-dimensional lattices of ions5.6 Explain why the formula for an ionic compound is an empirical formula
Ionic substances consists of orderly arrays of positively and negatively ions in a repeating 3D lattice For ionic compounds, specify the ratios in which the atoms (or ions) are present, not the composition of
discrete molecules Such formulae that give the ratio by atoms of elements of a compound rather than the actual number of
atoms in a molecule are called empirical formulas Formulae for ionic compounds are therefore always empirical formulae because there are no molecules
5.7 Identify common elements that exist as molecules or as covalent lattices
Element as covalent molecules or covalent lattices
Hydrogen gas, chlorine gas, oxygen gas are diatomic gases Br2 is a diatomic liquid while I2 is a diatomic solid Carbon exists as a diamond which is a 3D lattice and as graphite which is a 2D lattice Semi metals B, Si, Ge, As, Sb and Te closely approximate to covalent lattices through their bonding electrons
are not as firmly localised as in diamond. The noble gases (He, Ne, Ar, Kr, Xe, Rn exist as monatomic molecules; no chemical bonding
Covalent network solids in the earth
Many substances in the lithosphere are covalent network solids:
Sand and quartiz – silicon dioxide Gemstones may be silicon dioxide with traces of impurities which provide colour while others e.g. emerald,
aquamarine, topaz, garnet are silicates/alumina which are covalent lattices with some ionic parts incorporated
Mica, talc are also silicate lattices Clays are alumino silicate lattices, again with some ionic portions
5.8 Explain the relationship between the properties of conductivity and hardness and the structure of ionic, covalent molecular and covalent network structures
Properties of ionic substances
Melting an ionic substance means breaking up the orderly arrangement of ions As the electrostatic forces between ions are strong, much energy and high temperature is need to break the
bonds Boiling an ionic substance means producing a vapour that consists of well-separated ion pairs. This requires
an even greater amount of energy The strong electrostatic attraction between pairs of ions makes ionic substances hard If the orderly array of ions is disturbed by applying a strong force, then ions of the same charge come close
together. They then repel each other and this causes the crystal to shatter. This means that ionic crystals are brittle.
Solid ionic compounds do not conduct electricity because in the solid the ion are tightly bound into an orderly array and therefore unable to move towards a charged electrode.
Not all ionic substances are soluble. When ionic substances melt, the orderly arrangement of ions is broken up and ions can move about relatively freely molten ionic substances conduct electricity
When ionic substances are dissolved in water, the crystals are broken up and the ions are free to move about and hence conduct electricity
Properties of covalent molecular substances
Bonding forces holding atoms together within a covalent molecule are strong, intermolecular forces are weak
Boiling involves separating molecules from one another breaks intermolecular forces not covalent tbonds Melting just disrupts the orderly arrangement of molecules The stronger the melting and boiling point of covalent molecular compounds, the higher the melting and
boiling points Covalent molecular are neutral species, they cannot conduct electricity either as pure substance or in
solution Some covalent substances when mixed with water actually react and form ions e.g. HCl these conduct
electricity
Covalent network solids
Covalent bonding extends indefinitely throughout the whole crystal Covalent lattice solids/covalent lattices Lattice: infinite orderly array of particles Carbon in the form of diamond is a covalent network solid Since carbon is in group 4, each carbon atom is covalently bonded to four other carbon atoms Silica is also a covalent lattice (silicon dioxide) Silicon wants to form 4 covalent bonds, oxygen wants to form 2. Therefore in silica each silicon atom is
covalently bonded to 4 oxygen atoms, and each oxygen atom is bonded to 2 silicon atoms to form an infinite network of covalent bonds
The chemical formula for covalent lattice compound represent the ration which the atoms are present in the compound – empirical formula
Melting covalent lattices breaks many covalent bonds that are very strong. This process requires lots of energy and only occurs at high temperatures. Covalent lattices have extremely high melting points above 1000 degrees
Except for graphite, covalent network solids do not conduct electricity because they do not contain any ions and all their electrons are tied up either being held by individual atoms, or shared by pairs of atoms (no free electrons)
