METALS

Bonding & Structure Properties Alloys Chemical reactions Reactivity series

Bonding in metal What is Metallic Bonding?

As observed, each valence electron is detached from its parent atom and can move about freely. The electrons are said to be delocalized. The metal is held together by strong electrostatic forces of attraction between positive nuclei and the delocalized electrons.This is often described as an “array of positive ions in a sea of delocalized electrons”.

Strength of metallic bond What determines the strength of metallic bond?1. No. of delocalized valence electrons

As the number of delocalized electrons increases, the electron density of the “sea” will increase. The remaining ‘ions’ will also experience an increase in the charge.

This will result in a stronger attractive force between the delocalized electrons and the positive metal ions.

2. Atomic radius of the metallic atoms As the size of the atoms gets smaller, the delocalized electrons

are closer to the positive nuclei. The forces of attraction between the delocalized electrons and

positive nuclei will become stronger.

Properties due to bonding Implications of the metallic bond1. High boiling point

Boiling point is a better guide to the strength of the metallic bond than melting point.

In molten state (melting point), the metallic bonds are weakened. Only the ordered structure is destroyed.

Metallic bonds are completely broken during boiling. Large amount of energy required to break the stronger bonds.

2. Electrical Conductivity No. of delocalized valence electrons ∝ Electrical conductivity With more delocalized electrons, there are more charge carriers Metals are able to conduct electricity in all states.

Structure of Metals Metals are giant structures of atoms held together by strong metallic bonds. The structure is often described as “Giant Metallic Lattice”

Properties due to structure Physical properties due to its structure 1. Malleability and ductility

Malleable (can be beaten into different shapes) Ductile (can be pulled out into wires) Ability of the atoms to roll over each other into new positions,

without breaking the metallic bonds

Alloys & Pure metals

Pure metal Mixtures of metals Same atomic radius Atoms of varying atomic radius

Lower tensile strength Harder (higher tensile strength)

Less corrosion resistant More corrosion resistant

Fixed melting/boiling point Melts and boils over a range of temperature

Comparing Reactivity of Metals

In your group, design an experiment to compare the reactivity of: Magnesium; Zinc; Silver; Copper; AluminiumYou are to discuss on the following:1. Reagents to use2. Hypothesized results (observations & equations) of

your designed experiment3. Accounting reasons for the observations made

Reactivity Series

Devise a mnemonic Is carbon metallic? Is hydrogen metallic?

Chemical Reactions 1. Reaction of metals with Oxygen Metals burn in oxygen to form metal oxides Metal oxides are mostly basic, Al, Zn and Pb being

amphoteric Sodium metal reacts with oxygen in air at room

temperature, forming sodium oxide. Sodium metal is kept under kerosene to prevent its reaction

with oxygen and moisture in air. 2Na (s) + O2 (g) 2NaO (s) Calcium does not react with oxygen at room temperature.

On heating, calcium reacts with oxygen in air to form calcium oxide

2Ca (s) + O2 (g) 2CaO (s)

Chemical Reactions 1. Reaction of metal with Oxygen Zinc metal burns in air on strong heating to form zinc oxide. 2 Zn (s) + O2 (g) 2 ZnO (s)

Iron metal does not burn. The hot metal glows in oxygen and gives off yellow spark

4Fe (s) + 3 O2 (g) 2 Fe2O3 (S)

Copper metal does not burn. On prolonged strong heating, metal eventually coats with a black layer.

2Cu (s) + O2 (g) 2CuO (s)

Other metals like Gold, Silver and Platinum do not react with oxygen in the air.

Reactivity Series 2. Reaction of metals with Water

metal + cold water metal hydroxide + hydrogen If the metal reacts with steam, the metal oxide and hydrogen gas are formed.

metal + steam metal oxide + hydrogen

As you go down the reactivity series, the reactions become less and less vigorous.

If the metal reacts with cold water, the metal hydroxide and hydrogen gas are formed.

Metals above hydrogen in the reactivity series react with water (or steam) to produce hydrogen.

Metals below hydrogen in the reactivity series do not react with water or steam.

Reactivity Series 2. Reaction of metals with Water Group I metals and Calcium react with cold water forming

alkaline metal hydroxide and hydrogen gas Ca (s) + 2H2O (l) Ca(OH)2 (aq) + H2 (g) Magnesium reacts slowly with cold water but rapidly with

steam to form magnesium oxide and hydrogen gas Mg (s) + H2O (g) MgO (s) + H2 (g) Zinc reacts slowly with steam to form zinc oxide and

hydrogen gas Zn (s) + H2O (g) ZnO (s) + H2 (g) Iron reacts slowly and reversibly with steam to form iron

oxide and hydrogen gas 2Fe (s) + 3H2O (g) ⇌ Fe2O3 (s) + 3H2 (g)

Reactivity Series 3. Reaction of metals with Acids Dilute acids react with metals depending on their

positions in the reactivity series. Metals below hydrogen in the series do not react

with dilute acids. Metals above hydrogen in the series react to

produce hydrogen gas. The higher the metal in the series, the more

vigorous the reaction. You would NEVER mix sodium or potassium with acids!

Solve the puzzle…..

Some experiments were conducted to place copper, nickel and silver in reactivity series order.

Experiment 1: A piece of copper was placed in some green nickel (II) sulfate solution. There was no change to either the copper or the solution.

Experiment 2: A coil of copper wire was suspended in some silver nitrate solution. The solution gradually changes from colourless to blue and silvery crystals appeared on the copper wire.

Place the metals in reactivity series order…..

4. Displacement

Sodium Calcium CopperChloride

Reactivity Series 5. Thermal stability Measures the stability of the compound towards

decomposition when heated. A compound with high thermal stability will not decompose

as well when heated. The more reactive the metal, the more stable will be its

metal carbonate. X2CO3 (s) X2O (s) + CO2 (g), where X is an alkali metal

You are provided with a mixture of 2 compounds: Copper (II) carbonate and Sodium carbonate. State the observations (if any) upon heating.CuCO3 (s) CuO (s) + CO2 (g)There will be an observed change in colour, from green to black.There will no change to the white sodium carbonate. White precipitate is produced when the gas bubbles through limewater.

Reactivity Series

Why is knowledge of the reactivity series important ?

Predicting using reactivity series

Manganese , Mn, lies between aluminium and zinc in the reactivity series and forms a 2+ ion. Solutions of manganese (II) salts are very pale pink (almost colourless).

(a) Use the reactivity series to predict whether manganese will react with copper (II) sulfate solution. If it will react, describe what you would see, name the products and write an equation for the reaction.

(b) Would you expect Manganese to react with steam? If yes, name the products of the reaction and write the equation.

Chemical of the Week ALUMINIUM

Aluminium is the third most abundant element in Earth’ crust, after oxygen and silicon. Aluminium is too reactive chemically to occur in nature as the free metal. It is found combined in 270 different minerals. The most common ore of aluminium is bauxite -hydrated aluminium oxide.

Chemical of the Week Predict the reaction of aluminium in COLD WATER

Predict the reaction of aluminium in STEAM

Predict the reaction of aluminium in AIR (O2)

Predict the reaction of aluminium in ACIDS

No reaction with cold water!

Shows a slow reaction (or no reaction) with steam!

Reacts readily to form a non-porous layer of aluminium oxide that coats round the metal surface!

Reacts readily with acids and bases!

Chemical of the Week Aluminium is high on the reactivity series, but shows a lack of reactivity in WATER. Aluminium reacts readily and fast in air (oxygen), forming a non-porous layer of oxide – aluminium oxide. This layer of oxide acts as a ‘PROTECTIVE LAYER’. It adheres onto the metal’s surface, protecting it from further reaction.Unlike rust (hydrated iron oxide) which ‘flakes off’ the surface and exposing the surface to further corrosion.

RUSTING Rust is a chemical compound known as Hydrated Iron (III) oxide, which reddish brown in colour.

Hydrated iron (III) oxide = Fe2O3. xH2O From the name of the compound, the 3 important conditions for rusting to take place are : IRON, OXYGEN and WATER (moisture) Because iron combines so readily with oxygen, pure iron is rarely found in nature.