C2 topic 4 notes

Topic 4 – Groups in the periodic table • CHEMICAL CLASSIFICATION • The type of bonding affects the properties of the material…(see Topics 2 and 3) Type of Ionic Simple molecular Giant molecular bonding covalent covalent How when a metal reacts with between atoms of between atoms of bonds a non-metal: non-metal elements: non-metal elements form • Metals lose electrons • Electrons are • Electrons are to become positive shared between shared between cations atoms so they end atoms so they up with stable, end up with a • Non-metals gain full outer shells stable, full electrons to become outer shells negative anions • Oppositely charged ions attract Examples sodium chloride water, oxygen diamond Strength strong Strong covalent Strong across all of bonds bonds between atoms in a structure atoms, but weak forces holding separate molecules together Melting high - solids at room low – most are high – solids at and temperature liquids or gases at room temperature boiling room temperature points Solubility many dissolve in water some dissolve in Insoluble in water water Do they no no, except graphite • yes when molten or conduct in an aqueous electricity? solution • no when solid • Properties of metallic substances: • The atoms in metals are held together by metallic bonding (see below) – this gives metals different properties from other types of substances…: o Metals are good conductors of electricity and heat o Metals are solids at room temperature (àmetallic bonds are strong), except for mercury, which is a liquid at room temperature o Metals don’t dissolve in water o Metals are malleable (i.e can be hammered into shape) • àmetals have many uses…e.g they’re used to make cars, buildings, tools etc…

• METALLIC BONDING AND TRANSITION METALS • Metallic bonding: • Metal atoms form positive ions, which are held close together in a regular arrangement by a ‘sea’ of outer shell electrons: o The term ‘sea of electrons’ is used because electrons in the outer shells of metal atoms are free to move through the structure

o The electrons aren’t located in specific atomsàthey’re called ‘delocalised electrons’

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Metals conduct electricity: o Delocalised electrons move around randomly in all directions between the positive ions o If a potential difference (i.e voltage) is applied across a piece of metal, all the delocalised electrons start to move in the same direction o This movement of electrons is an electric current Metals are malleable: o If a large force is applied, the layers of positive ions in a metal can slide over each other o The positive ions are still held together by the sea of electronsàthe metal spreads out (changes shape) instead of breaking Transition metals: Most metals are transition metals Most transition metals have high melting and boiling points and form coloured compounds Transition metals are in the central block of the periodic table…:

ALKALI METALS The alkali metals are found in group 1 of the periodic table: o àthey have 1 electron in their outer shell… o àto end up with a full outer shell they must lose their outer electron o àthey form ions with a charge of +1 The atoms in alkali metals are held together by metallic bonding Alkali metals are solids at room temperature, but they have low melting points compared to other metals Alkali metals are soft metalsàcan be cut with a knife All alkali metals react with water to form a metal hydroxide and hydrogen gas: o E.g lithium + water à lithium hydroxide + hydrogen o 2Li (s) + 2H2O (l) à 2LiOH (aq) + H2 (g)

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o Note: all metal hydroxides are alkaline Reactivity: If lithium is dropped in water, it just floats on the water and fizzes (that’s the hydrogen being produced) until the reaction is finished Sodium reacts more strongly with water: o Sodium has a lower melting pointàthe reaction produces enough heat to melt the metal o This forms a molten ball of sodium that whizzes around the surface of the water (releasing hydrogen) until the reaction is finished Potassium reacts even more strongly and the hydrogen produced during the reaction catches fire, producing a lilac flame So the reactivity of alkali metals increases as you go down group 1 Explanation for this: o The elements at the bottom of group 1 have more electrons than the elements at the top of the groupàthey have more electron shellsàthe electron in the outer shell is further from the nucleus o The attraction between positive and negative charges (i.e between the nucleus and the outer electron) is weaker when charges are further apart o àouter electron in a potassium atom (configuration: 2.8.8.1) is not held as strongly as the outer electron in a lithium atom (configuration: 2.1) o Metals react by losing their outer electron, forming ions with a +1 charge and a full outer shell o Potassium loses its outer electron more easilyàit’s more reactive HALOGENS The halogens are the elements in group 7: o àhave 7 electrons in their outer shell… o àto form a full outer shell they must gain one electron o àthey form ions with a charge of -1 At room temperature…: o Fluorine is a pale yellow gas o Chlorine is a yellow-green gas o Bromine is a brown liquid o Iodine is a grey solid Reactivity of the halogens: The pattern of reactivity for the halogens is the opposite to that of the alkali metals - i.e halogens become less and less reactive as you go down group 7: o This is because halogens react by gaining an electron - this is easier with fewer electron shells (because outer electrons are closer to the nucleus) o So fluorine is the most reactive halogen HALOGEN REACTIONS Reactions with metals: All halogens react with metals to form metal halides (the word ‘halide’ means that the compound contains only metal ions and ions of one of the halogens): o E.g potassium + bromine à potassium bromide o 2K (s) + Br2 (l) à 2KBr (s) Note: chlorine forms ‘chloride’, fluorine forms ‘fluoride’, iodine forms ‘iodide’ Reactions with hydrogen: Halogens react with hydrogen gas to form hydrogen halides: o E.g hydrogen + chlorine à hydrogen chloride

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o H2 (g) + Cl2 (g) à 2HCl(g) Note: hydrogen halides form acids when they dissolve in water (e.g when hydrogen chloride dissolves in water it forms hydrochloric acid) Displacement reactions: More reactive halogens can ‘displace’ less reactive halogens from their compounds Less reactive halogens cannot ‘displace’ more reactive halogens from their compounds (in this case, no reaction takes place) àdisplacement reactions can be used to work out the reactivity of different halogen elements…

E.g chlorine is more reactive than bromine (as it is higher up group 7)àwill displace bromine from a bromide…: o sodium bromide + chlorine à sodium chloride + bromine o 2NaBr (aq) + Cl2 (g) à NaCl (aq) + Br2 (aq) o Note: the solution is initially colourless but will turn an orange-brown colour as bromine is formed E.g2 if chlorine is added to a compound of fluorine there will be no reaction because chlorine is less reactive than fluorine (àcan’t displace it) NOBLE GASES The noble gases are in group 0 of the periodic table Elements in other groups gain full outer shells by forming ions (i.e by losing or gaining electrons) or by forming covalent bonds (i.e by sharing electrons) However, noble gases already have full outer shellsàcompared to other elements, noble gases are inert (very unreactive) Properties of noble gases: All noble gases are gases at room temperatureàthey have low boiling points Particles in a gas are spread far apartànoble gases have low densities As you down group 0: o the boiling points of noble gases increase o the densities of noble gases increase Discovery of the noble gases: Chemists noticed that the density of pure nitrogen made in chemical reactions was less than the density of nitrogen extracted from the air It was hypothesised that the nitrogen extracted from air also contained a denser gas Experiments were done to find the identity of this dense gas – it turned out to be argon The other noble gases were discovered soon after Uses of the noble gases: The noble gases are useful because they are unreactive: o Xenon and argon were used inside filament lamps, instead of air, to stop the hot filament reacting with oxygen and burning away o Argon and helium are used in welding - they form a blanket over the hot metalàpreventing it from reacting with oxygen in the air Argon is non-flammableàused in fire-extinguisher systems Helium has a low densityàused for filling balloons and airships When electric current is passed through a tube filled with neon under low pressure, coloured light is producedàneon is often used in fluorescent lamps and advertising displays