Chemistry Unit 2

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C2 1.1 Chemical Bonding

  • Compounds are formed when two or more elements react together.
  • The atoms of elements join together by sharing electrons or by transferring electrons to achieve stable electronic structures. Atoms of the noble gases have stable electronic structures.
  • COVALENT BONDING is when atoms of non-metallic elements join together by sharing electrons.
  • IONIC BONDING is when metallic elements react with non-metallic elements to produce ionic compounds. The metal atoms lose electrons to form positive ions. The atoms of non-metals gain electrons to form negative ions. The ions have the stable electronic structure of a noble gas. The oppositely charged ions attract each other in the ionic compound - this is ionic bonding.
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C2 1.2 Ionic Bonding

  • Ionic bonding holds oppositely charged ions together in giant structures. The giant structure of ionic compounds is very regular because the ions all pack together neatly, like marbles in a box.
  • In all directions, there are strong electrostatic forces of attraction. Each ion in the giant structure or lattice is surrounded by ions with the opposite charge and so is held firmly in place.
  • Sodium chloride (NaCl) contains equal number of sodium ions and chloride ions as shown by its formula. The sodium ions and chloride ions alternate to form a cubic lattice.
  • Dot and cross diagrams can be used to represent the atoms and ions involved in forming ionic bonds.
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C2 1.3 Formulae of Ionic Compounds

  • When an ionic compound is formed, it becomes neutral. If we know the charge on each ion in a compound we can work out its formula by balancing the charges. Sodium chloride is NaCl (one Na+ ion for every one Cl- ion), but calcium chloride is CaCl2 (one Ca2+ ion for every two Cl- ions).
  • The charge on simple ions formed by elements in the main groups of the periodic table can be worked out from the number of the group. For transition metals the charge on the ion is shown by the Roman numeral in the name of the compoundm for example iron(II) sulfate contains Fe2+. (In the examination you will have a data sheet showing the charges of ions.)
  • Some ions are made up of more than one element, for example carbonate ions are CO32- and hydroxide ions are OH-. If we need to multiply these ions to write a formula, we use brackets. The formula of calcium carbonate is CaCo3, and the formula of calcium hydroxide is Ca(OH)2
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C2 1.4 Covalent Bonding

  • The atoms of non-metals need to gain electrons to achieve stable electronic structures. They can do this by sharing electrons with other atoms. Each shared pair of electrons strongly attracts the two atoms, forming a covalent bonds. Substances that have atoms held together by covalent bonding are called molecules.
  • Atoms of elements in Group 7 need to gain one electron and so form a single covalent bond. Atoms of elements in Group 6 need to gain two electrons and so from two covalent bonds. Atoms of elements in Group 5 can form three bonds and those in Group 4 can form four bonds.
  • A covalent bond acts only between the two atoms it bonds to each each other, and so many covalently bonded substances consist of small molecules. Some atoms that can form several bonds, like corbon, can join together in giant covalent structures, These giant covalent structures are sometimes referred to as macromolecules.
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C2 1.5 Metals

  • The atoms in a metallic element are all the same size. They form giant structures in which layers of atoms are arranged in regular patterns. You can make models of metal structures by putting lots of small same-sized spheres, like marbles, together.
  • METALLIC BONDING is when metal atoms pack together the electrons in the highest energy level (the outer electrons) delocalise and can move freely between atoms. This produces a lattice of positive ions in a 'sea' of moving electrons. The delocalised electrons strongly attract the positive ions and hold the giant structure together.
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C2 2.1 Giant Ionic Structures

  • Ionic compounds have giant structures in which many strong electrostatic forces hold the ions together. This means they are solids at room temperature. A lot of energy is needed to overcome the ionic bonds to melt the solids. Therefore ionic compounds have high melting points and high boiling points.
  • However, when an ionic compound has been melted the ions are free to move. This allows them to carry electrical charge, so the liquids conduct electricity. Some ionic solids dissolve in water because water molecules can split up the lattice. The ions are free to move in the solutions and so they also conduct electricity.

Ionic compounds cannot conduct electricity when solid because the ions can only vibrate about fixed positions; they cannot move around. The compounds must be melted or dissolved in water for the ions to be able to move about freely.

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C2 2.2 Simple Molecules

  • The atoms within a molecule are held together by strong covalent bonds. The bonds act only between the atoms within the molecule, and so simple molecules have little attraction for each other. Substances made of simple molecules have relatively low melting points and boiling points. They do not conduct electricity because molecules have no overall charge and so cannot carry electrical charge.
  • INTERMOLECULAR FORCES are the forces of attraction molecules, and are weak. These forces are overcome when a molecular substance melts or boils. This means that substances made of small molecules have low melting point and boiling points. Those with the smallest molecules, like H2, Cl2 and CH4, have the weakest intermolecular forces and are gases at room temperature. Larger molecules have stronger attractions and so may be liquid at room temperature, like Br2 and C6H14, or solids with low melting points.
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C2 2.3 Giant Covalent Structures

  • Atoms of some elements can form several covalent bonds. These atoms can join together in giant covalent structures (sometimes called macromolecules). Every atom in the structure is joined to several other atoms by strong covalent bonds. It takes an enormous amount of energy to break down the lattice and so these substances have very high melting points.
  • Diamond is a form of carbon that has a regular three-dimensional giant structure. Every carbon atom is covalently bonded to four other carbon atoms. This makes diamond hard and transparent. The compound silicon dioxide (silica) has a similar structure.
  • Graphite is a form of carbon in which the atoms are covalently bonded to three other carbon atoms in giant flat two-dimensional layers. There are no covalent bonds between the layers and so they slide over each other, making graphite slippery and grey.                                             

Differences

Similarities

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C2 2.3 Giant Covalent Structures (2)

Bonding in Graphite and Fullerenes

In graphite each carbon atom bonds covalently to three other carbon atoms forming a flat sheet of hexagons. One electron from each carbon atom is delocalised, rather like electrons in a metal. These delocalised electrons allow graphite to conduct heat and electricity.

There are only weak intermolecular forcs between the layers in graphite, so the layers can slide over each other quite easily.

FULLERENES are large molecules formed from hexagonal rings of carbon atoms. The rings join together to form cage-like shapes with different numbers of carbon atoms, some of which are nano-sized. Scientists are finding many applications for fullerenes, including drug delivery into the body, lubricants, catalysts and reinforcing materials.

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C2 2.4 Giant Metallic Structures

  • Metal atoms are arranged in layers. When a force is applied the layers of atoms can slide over each other. They can move into a new position without breaking apart, so the metal bends or stretches into a new shape. This means that metals are useful for making wires, rods and sheet materials.
  • Alloys are mixtures of metals, or metals mixed with other elements. The different sized atoms in the mixture distort the layers in the metal structure and make it more difficult for them to slide over each other. This makes alloys harder than pure metals.
  • Shape memory alloys can be bent or deformed into a different shape. When they are heated they return to their original shape. They can be used in many ways such as dental braces.

Metal structures have delocalised electrons. Metals are good conductors of heat and electricity because the delocalised electrons move throughout the giant metallic lattice and can transfer energy quickly.

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C2 2.5 The Properties of Polymers

  • The properties of a polymer depend on the monomers used to make it, and the conditions we use to carry out the reaction. Poly(propene) is made from propene and softens at a higher temperature than poly(ethene) and high density (HD) poly(ethene) are made using different catalysts and different reaction conditions. HD poly(ethene) has a higher softening temperature and is stronger than LD poly(ethene).
  • Poly(ethene) is an example of a thermosoftening polymer. It is made up of individual polymer chains that are tangled together. When it is heated it becomes soft and harder agains when it cools. This means it can be heated to mould it into shape and it can be remoulded by heating it again.
  • Other polymers called thermosetting polymers do not melt or soften when we heat them. These polymers set hard when they are first moulded because strong covalent bonds form cross-links between their polymer chains. The strong bonds hold the polymer chain in position.
  • In thermosoftening polymers the forces between the polymer chains are weak. When we heat the polymer, these weak intermolecular forces are broken and the polymer becomes soft. When the polymer cools down, the intermolecular forces bring the polymer molecules back together so the polymer hardens again.
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C2 2.6 Nanoscience

  • Nanoscience is a new and exciting area of science. When atoms are arranged into very small particles they behave differently to ordinary materials made of the same atoms. A nanometre is one billionth of a metre and nanoparticles are a few nanometres in size. They contain a few hundred atoms arranged in a particular way. Their very small sizes give them very large surface areas and new properties that can make them very useful materials.
  • Nanotechnology uses nanoparticles as highly selective sensors, very efficient catalysts, new coatings, new cosmetics such as sun screens and deodorants, and to give construction materials special properties.
  • If nanoparticles are used more and more there will be a greater risk of them finding their way into the air and into our bodies. This could have unpredictable consequences for our health and the environment. More research needs to be done to find out their effects.
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C2 3.1 The Mass of Atoms

  • Protons and neutrons have the same mass and so the relative masses of a proton and a neutron are both one.
  • The mass of an electron is very small compared with a proton or a neutron, and so the mass of an atom is made up almost entirely of its protons and neutrons. The total number of protons and neutrons in an atom is called its mass number.
  • Atoms of the same element all have the same atomic number. The number of protons and electrons in an atom must always be the same, but there can be different numbers of neutrons.
  • Atoms of the same element with different numbers of neutrons are called isotopes.
  • The number of neutrons in an atom is equal to its mass no. minus its atomic no.(http://img.sparknotes.com/figures/C/c3554c42f955f947b125020c2c00d830/carbon.gif)
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C2 3.2 Masses of Atoms and Moles

  • Atoms are too small to weigh and so we use relative atomic masses (Ar) in calculations. Relative atomic masses are often shown in periodic tables. In the lab, substances are usually weighed in grams. The relative atomic mass of an element in grams is called one mole of atoms of the element.

Relative Atomic Mass

  • We use an atom of C as a standard atom and compare the masses of all other atoms with this. The relative atomic mass of an element (Ar) is an average value that depends on the isotopes the element contains. However, when rounded to a whole number it is often the same as the mass number of the main isotope of the element.

Relative Formula Mass

  • The relative formula mass (Mr) of a substance is found by adding up the relative atomic masses of the atoms in its formula.
  • The relative formula mass of a substance in grams is called one mole of that substance. Using moles of substances is useful when we need to work out how much of a substance reacts or how much product we will get.
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C2 3.3 Percentages and Formulae

  • We can calculate the percentage of any of the elements in a compound from the formula of the compound. Divide the relative atomic mass of the element by the relative formula mass of the compound and multiply the answer by 100 to convert it to a percentage. This can be useful when deciding if a compound is suitable for a particular purpose or to identify a compound.

Working Out the Formula of a Compound from its Percentage Composition

  • The empirical formula is the simplest ratio of the atoms of ions in a compound. It is the formula used for ionic compounds, but for covalent compounds it is not always the same as the molecular formula. For example, the molecular formula of ethene is C2H6, but its empirical formula is CH3.
  • We can calculate the empirical formula of a compound from its percentage composition:
  • Divide the mass of each element in 100g of the compound by its Ar to give the ratio of atoms.
  • Then convert this to the simplest whole number ratio.
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