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

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:

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. The sample is injected into the injector port and is vaporised. It is swept along by an inert carrier gas through a temperature controlled, looped metal or glass column. The stationary phase is crucial to the separation. Where the stationary phase is a liquid it is coated onto the surface of an inert material (often glass spheres). As the separated components emerge from the column, their presence is detected and measured, usually by a flame ionisation detector. The result is translated into a series of peaks on a printout. The retention time, Rt, a measure of how long the component takes to emerge from the column, is used to identify the component, and the area under the peak is related to the amount of the compound present. It is necessary to prepare a calibration curve to allow the concentration to of the component to be calculated. Calibration compares the reading from an instrument against an accurately known standard. Here the peak area is measured and the quantity of interest is concentration. The peak area is plotted against a set of known concentrations of samples of the identical compound, called standards, to prepare a calibration curve. Calibration curves need to be prepared for every procedure because the rate at which components travel is affected by a large number of factors.

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 There are other separation techniques that are not mentioned here, such as sedimentation, crystallization, centrifugation and decantation. These can be mentioned in class discussions.

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

25

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. liquid gas solid

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)

36 Structure of the Atom Isotopes Name # p # n # e Hydrogen 1 Deuterium
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 Deuterium Tritium 2 Carbon 12 6 Carbon 13 7 Carbon 14 8 Uranium 235 92 143 Uranium 238 146

37 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.

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 Electron configurations for the noble gases (note the full s and p orbitals i.e. full octet in valence shell except for He) HeliumElement #2 1s2 NeonElement #10 1s2 2s2 2p6 ArgonElement #18 1s2 2s2 2p6 3s2 3p6 KryptonElement #36 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 XenonElement #54 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 RadonElement #86 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 Note: For Interest Only! You are not required to learn this information for the HSC

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 - + Example: Mg2+ + Cl-  MgCl2
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 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 - +

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 Water Chlorine Hydrogen Chloride
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 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
Boiling – breaking H bonds Electrolysis – breaking covalent bonds O H H O H O H H O H + + O H O O H 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 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

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 Synthesis Everyday Applications
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

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 S·m-1 (20 °C) Chemical properties 4Al(s)+ 3O2(g)  2Al2O3(s) 2Al(s) + 6HCl(g)  2AlCl3(aq) + 3H2(g) Oxygen: M.P. = -2190C Density = g/cm3 Conductivity = non-conductor Bauxite ore (Al2O3·xH2O) Physical properties M.P. = 20450C Density = g/cm3 Conductivity = non-conductor Chemical properties Al2O3(s) + 2NaOH + 3H2O(l)  2Na[Al(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 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.

66 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.

67 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.

68 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.

69 The Chemical Earth Compiled by: Robert Slider (2006)
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