Chemical Earth Summary

Chemistry Summary

The particle theory:

1. Matter is made of tiny particles

2. Particles of matter are in constant motion

3. Particles of matter are held together by very strong electric forces

4. There are empty spaces between the particles of matter that are very large compared to the particles themselves.

5. Each substance has unique particles that are different from the particles of other substances

6. Temperature affects the speed of the particles. The higher the temperature, the faster the speed of the particles.

The Biosphere, Lithosphere, Hydrosphere and Atmosphere contain mixtures of elements and compounds.

• Biosphere – consists of all living organisms and their life cycles. The most abundant elements are oxygen (60%), carbon (21%) and hydrogen (11%), which are mostly found within the DNA of living things and organic compounds such as amino acids, carbohydrates, fats and proteins.

• Atmosphere – is the layer of gas surrounding the earth and is a mixture composed mostly of uncombined lighter elements, although composition varies according to location or climate. The two most abundant elements are nitrogen (75.3%) and oxygen (23.2%), which occur in their gaseous forms or in compounds like carbon dioxide.

• Lithosphere – is made up of the crust and the solid top of the mantle, and is composed mostly of rocks and soil (mixtures of minerals), which can have a definite composition or a range of compositions. Elements such as oxygen (47%), silicon (28%), aluminium (8%) and iron (5%) occur in silicate, oxide and carbonate minerals. Pure elements such as gold also occur in their pure form (uncombined) in the lithosphere.

• Hydrosphere – is all the water on earth and is a discontinuous sphere. It has a varied composition (i.e., salt and fresh water), and is mainly made up of the compound water (H2O) but also contains many water soluble sulfates and soluble carbonates, and elements like chlorine and sodium dissolved as ions.

Separation Procedures: - usually separations require a number of processes or procedures to separate the constituent substances from a mixture. These separations are physical separations.

• Filtration – separates undissolved solids from liquids or gases by passing the mixture through a screen such as filter paper which is fine enough to collect the particles of the solid.

• Solution – usually used in combination with another separation method, and is based on the fact that some constituents in a mixture dissolve in a solvent such as water more readily than others. That is, a mixture is added to a solvent and can be separated through the fact that one constituent will dissolve more readily than the others (although to fully separate the mixture, another method such as filtration would have to be used).

• Evaporation – relies upon the varying evaporating points of the constituent substances within a mixture. A mixture is heated (in an evaporating basin and the process is usually sped up using a Bunsen burner) and one substance will evaporate, leaving the other substance behind.

• Crystallisation - depends on the components of the mixture having different solubilities in a selected liquid (usually water) at different temperatures. For example, a mixture of salt and baking powder are both soluble in hot water, but when the hot water (with the mixture dissolved in it) is cooled, the baking soda will crystallise because it is much less soluble at cooler temperatures.

• Sedimentation – occurs when solid particles are allowed to settle from water (or other liquids) or air. This occurs most readily when the solvent is not moving.

• Decantation – the process of pouring off a liquid above a solid which has been allowed to settle by sedimentation.

• Sieving – the process of separating solid particles of various sizes.

• Centrifugation – involves a centrifuge that spins and separates solids and liquids.

• Distillation – is effective where the constituent substances of a mixture have very different boiling points. The mixture is heated and the substance with the lowest boiling point boils, is cooled in a condenser and is collected as a pure liquid. The components with the higher boiling points remain in the distilling flask. Fractional distillation is also used to distill several liquids with similar boiling points. This requires many stages of vaporisation-condensation, which gradually separates the substances.

• Separating Gases In The Air – to separate gases, the liquefaction of the gases is needed, and then fractional distillation used.

• Separating Immiscible Liquids – a separating funnel is used, so the denser liquid runs though the tap in the funnel and then the tap can be turned off, leaving the layer of liquid that was on top of the denser liquid.

• Chromatography – the mixture is passed over the surface of an inert substance such as chromatography paper. The separation of the components in the mixture occurs because the components absorb, or cling, to the surface of the inert substance with different strengths.

• Magnetic separation – by passing a mixture between the poles of a strong electromagnet, substances can be extracted using their magnetic properties, thus separating the magnetic substances from the non-magnetic.

Classification of Elements:

Elements vary significantly in their physical and chemical properties.

|METALS |SEMI-METALS (also known as METALLOIDS) |NON-METALS |
|Lie on the left hand side of the periodic |Lie along the border between metals and |Lie to the right and side of the periodic |
|table |non-metals (B, Si, Ge, As, Sb, Te, Po) |table |
|Donate electrons (form positive ions) |Varies, depending on where it lies (to the |Receive electrons (form negative ions) |
| |right or the left of the dividing line) | |
|Good conductors of heat and electricity |Semi-metals have no defined properties, their |Insulators – do not conduct heat and |
| |properties are varied. |electricity (there are some exceptions) |
|High strength | |Low strength |
|Ductile (able to be stretched into a wire) and| |Brittle – not ductile or malleable. |
|malleable (able to be hammered into a flat | | |
|sheet) | | |
|Lustrous (shiny) | |Not lustrous (not shiny) |

Chemical change: MUST include the transfer of electrons, and therefore change the electron configuration of an element.

Physical change: does not change the electron configuration of the element.

Chemical property: a property that can only be determined through chemical changes/processes

Physical property: any property that can be determined WITHOUT the use of chemical reactions.

An element’s properties are based on the structure of the electrons in its shells.

Lewis dot diagrams: - only draw the valence (outer shell) of the element. See notes for details.

Structural formula: - bonds indicated by lines. Non-bonding pairs are not shown in a structural formula. See notes for details.

Discrete molecules: - which travel around as singular entities, do not bond with other molecules. For example, water molecules.

CHEMICAL BONDS:

Ionic Bonds: - are strong bonds that form between cations (positive charge) and anions (negative charge). They occur because of the attraction between the ions and they only form between non-metals and metals.

o Cations are always smaller than the neutral atom an anions are always larger than their parent ion. This is because in a cation, the positive charge of the nucleus draws the electrons in, and in an anion, the electrons are repulsed by the nucleus.

o Conduct electricity when their ions become mobile (like when dissolved in a polar solvent) and polarity (net charge on molecules) is needed to conduct electricity.

o Ionic compounds have certain properties because of the nature of their bonds...

|PROPERTY |COMMENT |REASON |
|Melting point |High |The strong attractions between opposite charged ions need high energy |
| | |levels to be overcome. |
|Electrical |No conductivity in a solid |In a solid state the ions are not mobile/free to move around |
|conductivity |state |In a molten state or aqueous solution the ions are highly mobile, and can |
| |Excellent conductivity if |therefore transmit the electric charge. |
| |molten or in an aqueous | |
| |solution/state | |
|Ductile |No- they are very brittle |As distortion occurs in the lattice, ions of the same charge come in close|
| | |proximity and repulsion, rather than an attractive force, is set up, |
| | |resulting in the collapse of the crystal structure. |
|Solubility |High in POLAR solvents* |Polar water molecules surround the ions at the edge of the crystal. This |
| | |shields the attraction between opposite charged ions in the crystal and |
| | |the water molecules carry the ion away from the lattice structures.** |

* LIKE DISSOLVES LIKE: polar solvents (like water) dissolve polar solutes (such as ionic compounds). Non-polar solvents dissolve non-polar solutes, but if the solute and the solvent are not of the same “polarity” (or are not BOTH polar or BOTH non-polar), the solvent will not dissolve the solute.

** The dissolution of ionic substances in polar solvents is a PHYSICAL CHANGE, as no electrons are transferred in the process.

Metallic Bonding: - refers to the bonding within metals. Usually it refers to the bonding between the atoms of one metallic element.

• Outer shell electrons are held very loosely in metals.

• A metal lattice consists of positive ions held in fixed positions, surrounded by a ‘sea’ of delocalised electrons, which hold the positive ions in place.

[pic]

Covalent Bonding: - involves the sharing of valence electrons, usually between two non-metallic elements. Covalent bonding results in discrete molecules, as well as covalent lattice formations. Covalent bonds are very strong, and can form with single, double or triple bonds.

• Co-ordinate covalent bonding – is where one atom supplies BOTH the electrons in the bonding pair. For example, ammonium. There is no distinction between that bond and the other bonds once the bond is formed, and any charge that may occur will be spread over the entire molecule.

• Covalent network solids – e.g., carbon as diamond and graphite, as well as silicon dioxide. All atoms are held together in a network lattice. In the case of diamond and SiO2, the tetrahedral structure with strong covalent bonds results in extremely hard solids and high melting points.

• Covalent molecular substances – are non-polar due to equal sharing of electrons (like in H2, F2 and O2. As a result, dispersion forces are weak, and they have very low melting and boiling points, yet it is difficult to break the covalent bonds themselves.

NAMING COMPOUNDS:

• Ionic compounds:

o State the metal first, adding the valency in roman numerals if it is located in the transitional area of the periodic table.

o State the non-metal, adding the suffix –ide.

o E.g., NH4Br – Ammonium Bromide

|Prefix |Valency |
|Mon(o)- |1 |
|Di- |2 |
|Tri- |3 |
|Tetra- |4 |
|Pent- |5 |
|Hex- |6 |
|Hept- |7 |

• Covalent compounds:

o Valencies are not so strict in covalent bonding, so the number of atoms in each molecule must be specified by adding prefixes.

o The first element identified does not need the prefix mono- if there is only one particle, however, if there is more than one particle, a prefix should be used.

o The most electronegative element (closest to fluorine) is always named last.

o E.g., PCl3 – Phosphorous trichloride and Cl2O7 – dichloride heptoxide

o Prefixes: (see table) ---------------------------------------->

|Number Of Carbons |Prefix Used |
|1 |Meth- |
|2 |Eth- |
|3 |Prop- |
|4 |But- |
|5 |Pent- |
|6 |Hex- |
|7 |Hept- |
|8 |Oct- |
|9 |Non- |
|10 |Dec- |

• Organic Compounds:

o The names of organic molecules are broken into two parts

o The prefix is in accordance to the number of carbons in the compound ------------------------------------------------------->

o The suffix for naming hydrocarbons is based on the bonds between the carbons in the chain.

▪ -ane : refers to single bonds

▪ -ene : refers to one double bond

▪ -yne : refers to one triple bond.

▪ -ol : refers to a hydrocarbon with a hydroxide (OH) group.

Note: the larger these hydrocarbon chains get, the higher the melting point, due to the greater dispersion forces acting along the hydrocarbon chains, joining them to other hydrocarbon chains.

• The more electronegative an element is, the keener it is to pull electrons towards its nucleus. The electronegativity of an element increases as we move towards the top right of the periodic table, i.e., the most electronegative element is fluorine. (as the noble gases are ignored, due to their unreactive noble gas configuration).

• The more reactive an element is, the keener it is to give away electrons. These are generally the metals. Reactivity increases as we move towards the bottom left corner of the periodic table, with the most reactive element being Francium.

• The more reactive (this term also encompasses electronegativity) an element is, the more eager it is to bond with other atoms through their eagerness to gain a noble gas configuration (through giving, receiving or sharing electrons). Therefore, the more reactive an element is (the further towards the extremes of the periodic table it lies); the more likely it is to exist as a bonded atom. The greater the reactivity of an element, the smaller the probability of it existing as an uncombined element.

• Ions are arranged around the nucleus in shells, which number 2, 8 and 8 electrons respectively. Within these shells are subshells (however, these are not gone into in depth at HSC level). Atoms attempt to gain a noble gas configuration (full outer shell, which is also known as a “valence” shell) and will lose or gain electrons depending on what requires the least energy.

Isotope: describes the same element (i.e., same number of protons in the nucleus) but with differing numbers of NEUTRONS in the nucleus, therefore altering its relative atomic mass.
E.g., Uranium-238 has 92 protons and 146 neutrons whilst Uranium-235 has 92 protons and 143 neutrons.

Allotrope: refers to the same element. However, the different allotropes of these elements have different physical properties due to the different ‘packing’ arrangement of atoms.
e.g., carbon – can be found in many different allotropes – soot, graphite, diamond, bucky bells and bucky tubes.

• The properties of elements can be very different from the compounds they form. For example, sodium chloride is ordinary table salt, but chlorine is a very deadly gas and sodium is a lustrous metal.

• The tetrahedral structure of electrons is maintained in molecules, although lone pairs will sometimes take the place of bonding pairs.

• INTERmolecular forces: between molecules.

• INTRAmolecular forces: within molecules.

Polar Molecules:

• All ionic bonds are polar.

• Covalent bonds depend on the symmetry of the molecule to determine the polarity (structure of electrons, and if they cancel each other out).

• Elements that are more electronegative have a negative polarity (as they hold electrons closer to the nucleus) and more reactive elements have a positive polarity (as they push electrons away).

• Dipole attractions are weaker than ionic bonds, as dipoles (the poles in polar molecules) are of weaker strength and therefore attraction. Refers to when molecules are held together by the attractions of their dipoles (an example of intermolecular forces).

• Dispersion forces are also examples of intermolecular forces. They arise from attractions between the positive nucleus of one atom and the negative electron cloud around another atom as well as the attraction between the temporary dipoles set up in adjacent molecules.

• Hydrogen bonds – (to be done in detail in another topic) –are intermolecular forces. Consider them as a special case of dipole attraction.

[pic]

Relative strengths of each bond. Note that covalent bonds are intramolecular forces, whereas dipole attractions and dispersion forces are intermolecular forces.

Reaction Rates:

• Are measured by the percentage of products per unit of time.

• (Collision Theory) – collision must occur between molecules for reactions to occur. However, they must collide with enough energy to start to break the bonds holding reacting molecules together, and must be oriented in the correct position for new bonds to form.

• Effective collisions – the higher the frequency of the collisions, the greater the probability of an effective collision and the reaction occurs faster. Activation energy refers to the minimum amount of energy needed for an effective collision, and the lower this energy level is, the faster the reaction occurs.

COMPARING THE PROPERTIES OF THE FOUR BONDING TYPES.

|Bonding Type |Melting Point |Conductivity |Solubility |Malleability/ |
| | | | |Ductility |
| | |Solid |Molten state | | |
|Metallic |High |High |High |None |Variable hardness, |
| | | | | |malleable |
|Ionic |High |None |High |Yes – in polar |Very hard and |
| | | | |substances |brittle |
|Covalent molecular |Low |None |None |Yes |Soft, brittle |
| | | | | |solids; are often |
| | | | | |liquids or gases at |
| | | | | |SRP. |
|Covalent network |Very high |None |None |None |Very hard, very |
| | | | | |brittle. |

For more information, and for justifications as to how these properties relate to bonding, see pages 94-98 in textbook (tables).

Boiling water to make steam is a physical change, as the molecules are still intact, the intermolecular forces are simply weakened and the water molecules become more mobile. However, in the electrolysis of water, involves the breaking of the strong bonds between the hydrogen and oxygen molecules, which requires greater amounts of energy. Water is passed through an apparatus, with two different charged rods. These attract the oxygen and hydrogen molecules and supply electrons, with which to make a noble gas configuration within these elements, thus separating the H2O molecules into separate molecules of O2 and H2 gas.

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(Right) Protons surrounded by a sea of delocalised electrons