Section 2: Chemistry of the Elements - Shawon Notes

Chemistry > Section 2: Chemistry of the elements

a) The Periodic Table

2.1 understand the terms group and period

Groups are the vertical columns in the periodic table. Periods are the horizontal rows in the periodic table.

2.2 recall the positions of metals and non-metals in the Periodic Table

Metals are on the left side of the zig-zag line in the periodic table. Non-metal is on the right.

2.3 explain the classification of elements as metals or non-metals on the basis of their electrical conductivity and the acid-base character of their oxides

Metals:

  • Tend to be solids with high melting and boiling points, and with relatively high densities(but as with several of the properties in this list, there are exceptions- like mercury is a liquid.
  • Are shiny when they are polished and tend to be easily workable.
  • Are good conductors of electricity and heat
  • Form positive ions in their compounds
  • Have oxides which tend to be basic, reacting with acids to give a salt and water.

Non-metals:

  • Tend to have low melting and boiling points(with the exceptional of carbon and silicon)
  • Tend to be brittle as solids and even if they are crstalline, they dont have the same sort of shine as metals.
  • Don’t usually conduct electricity(with the except
  • Are poor conductors of heat
  • Tend to form negative ions and covalent compounds
  • Have oxides which are acidic or neutral

2.4 understand why elements in the same group of the Periodic Table have similar chemical properties

Elements in the same group have similar properties because they have the same number of electrons in their outer electron shell.

2.5 understand that the noble gases (Group 0) are a family of inert gases and explain their lack of reactivity in terms of their electronic configurations.

Physical properties:

  • Colorless gases
  • Monatomic: ie. They consist of only one atom
  • Densities and boiling points are in a trend. They increases if you go down.

Chemical properties:

  • They do not react to form ionsBecauseandthey alreadyso havedon’ttheirouterprodu shell full and is stable.
  • They are unlikely to form covalent compounds. Because it costs too much energy to rearrange the full energy levels to produce the single electron that an atom needs.

b) Group 1 elements - Lithium, Sodium an Potassium

2.6 describe the reactions of these elements with water and understand that the reactions provide a basis for their recognition as a family of elements

Lithium: 2Li (s) + 2H2O(l) ==> 2LiOH(aq) + H2(g)

Lithium’s melting point is higher and the heat isn’t

Potassium: 2K(s) +2H2O(l) ==> 2KOH(aq) + H2 (g)

Potassium’s reaction is faster than sodium’s. enough flame. The reaction often ends with the potassium spitting around.

Sodium: 2Na(s) + 2H2O(l) ==> 2NaOH(aq) + H2(g)

The sodium floats on water because it is less dense. It melts because its melting point is low and a lot of heat is produced in this reaction. A white trail is formed inside the water which is sodium hydroxide.

2.7 describe the relative reactivities of the elements in Group 1

Cs>Rb>K>Na>Li

2.8 explain the relative reactivities of the elements in Group 1 in terms of distance between the outer electrons and the nucleus.

Group 1 elements lose electrons to non-metals to react. As you go down the group, the reactivity of the elements increases. Because the number of electrons increase if we go down. Higher electrons can form more shells. Higher electron shell reduces the attraction between proton and electron. So elements can easily lose electrons.

c) Group 7 elements - Chlorine, Bromine and Iodine

2.9 recall the colours and physical states of the elements at room temperature

State in room temperature

Colours

F 2

Gas

yellow

Cl 2

Gas

Green

Br 2

Liquid

Dark red liquid –red brown vapour

I 2

solid

Dark grey solid- purple vapour

2.10 make predictions about the properties of other halogens in this group

Halogens get higher melting and boiling points and darker colour down the group.

2.11 understand the difference between hydrogen chloride gas and hydrochloric acid

Both hydrogen chloride gas and hydrochloric acid have the formula HCl. Hydrogen chloride is a gas but hydrochloric acid is a solution of hydrogen chloride in water.

Hydrogen chloride is acidic in water as it produce hydroxonium ions and chloride ions:

H2O(l) + HCl(g) ==> H3O+(aq) +Cl-(aq)

2.12 explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in methylbenzene

When hydrogen chloride is dissolved in water, it ionize i.e. it forms ions. That’s why it chloride is dissolved in methylbenzene, no ions are formed.

2.13 describe the relative reactivities of the elements in Group 7

The reactivity falls quickly as you go down the group.

F>Cl>Br>I

2.14 describe experiments to demonstrate that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts

Reacting chlorine with potassium bromide

If you add chlorine solution to colourless potassium solution, the solution becomes orange as bromine is formed.

2KBr(aq) + Cl2(aq) ==> 2KCl(aq) + Br2(aq)

Reacting chlorine with potassium iodide

Adding chlorine solution to potassium iodide solution gives a dark reddish-brown solution of iodine.

2KI(aq) + Cl2(aq) ==> 2KCl(aq) + I2(aq or s)

2.15 understand these displacement reactions as redox reactions.

In the reaction between chlorine and potassium bromide, bromine loosed electrons. That means it is oxidized. Chlorine gained electron which means it is reduced. When both oxidation and reduction takes place in the same reaction, it is said to be redox reaction.

d) Oxygen and Oxides

2.16 recall the gases present in air and their approximate percentage by volume

Gas

Amount in air (%)

Amount in air (fraction)

Nitrogen

78.1

About 4 / 5

Oxygen

21.0

About 1 / 5

Argon

0.9

Carbon dioxide

0.004

2.17 explain how experiments involving the reactions of elements such as copper, iron and phosphorus with air can be used to investigate the percentage by volume of oxygen in air

Using copper:

The apparatus originally contains 100cm3 of air. This is pushed forward and backward over the heated copper, which turns black as copper(II) oxide is formed. The volume of gas reduce as the oxygen is used up.

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

As the copper reacts, the Bunsen is moved along the tube so that it is always heating fresh copper. Eventually all the oxygen in the air is used up. The volume stops reduc the percentage composition of 21% in air.

Using rusting or iron:

A test tube is taken and damp iron wool is placed. This is inverted and placed in a beaker containing water. The tube is now left for weeks. Iron uses oxygen and water to form rust. As long as the wool is damp, the rusting will continue.

The water level rises as oxygen is used. This is marked with a rubber band. At the end of the experiment, we will see all the oxygen is used up. The original water level was 15cm3 and now it is 12 cm3. That means 3 cm3 is used up.

3 / 15 x 100 = 20%.

2.18 describe the laboratory preparation of oxygen from hydrogen peroxide, using manganese(IV) oxide as a catalyst

Oxygen is made in the lab from hydrogen peroxide solution using manganese(IV) oxide as a catalyst. The reaction is known as catalytic decomposition of hydrogen peroxide.

2H2O2(aq) ==> 2H2O(l) + O2(g)

Testing oxygen: Oxygen relight a glowing splint.

2.19 describe the reactions of magnesium, carbon and sulfur with oxygen in air, and the acid-base character of the oxides produced

Burning magnesium

Magnesium burns in air with a bright white flame to give a white, powdery ash of magnesium oxide. The flame is extremely bright in pure oxygen.

2Mg(s) + O2(g) ==> 2MgO(s)

Burning carbon

Carbon burns if it is heated very strongly in air or oxygen to give colourless carbon dioxide gas. The carbon may produce a small yellow-orange flame and perhaps some sparks. It depends on the purity of the carbon.

C(s) +O2(g) ==> CO2(g)

Burning sulfur

Sulfur burns in air with a tiny, almost invisible blue flame. In oxygen it burns much more strongly giving a bright blue flame. Poisonous, colourless sulfur dioxide gas is produced.

S(s) + O2(g) ==> SO2(g)

Acid base character of oxides:

Few metal oxides react or dissolve in water –to form alkaline solutions. Most metal oxides do not. Shaking a solid magnesium oxide with water doesn’t seem to dissolve. alkaline.

MgO(s) + H2O(l) ==> Mg(OH)2(s and aq)

Many non-metals react with water to give acidic solutions except water and carbon dioxide. For example sulfur dioxide reacts with water to give sulfurous acid. Sulfurous acid is fairly acidic.

H2O(l) + SO2(g) ==> H2SO3(aq)

2.20 describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid

Carbon dioxide is most easily made by the reaction between dilute hydrochloric acid and calcium carbonate in the form of marble chips. Carbon dioxide is collected in an inverted jar.

CaCO3(s) + 2HCl(aq) ==> CaCl2(aq) + CO2(g) + H2O(l)

Testing carbon dioxide:

Carbon dioxide turns lime water(calcium hydroxide solution) milky. It reacts to give a white precipitate of calcium carbonate.

Ca(OH)2(aq) + CO2(g) ==> CaCO3(s) + H2O(l)

With an excess of carbon dioxide, the precipitate dissolves again to give a colorless solution of calcium hydrogen carbonate.

CaCO3(s) + CO2(g) + H2O(l) ==> Ca(HCO3)2 (aq)

2.21 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper(II) carbonate

When metal carbonates are heated they become carbon dioxide and a metal. For example:

copper carbonate > copper oxide + carbon dioxide CuCO3 > CuO + CO2

2.22 describe the properties of carbon dioxide, limited to its solubility and density

Carbon dioxide is a colourless, odourless gas, denser than air, and slightly soluble in water.

2.23 explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density

It is used in carbonated drinks because it dissolves in water under pressure. When you open bottle, the pressure falls and the gas bubbles out of solution.

It is also used in fire extinguishers to put out electrical fires, or those caused by burning liquids, where using water could cause problems. The dense gas sinks onto the flames and prevents any more oxygen from reaching them.

2.24 understand that carbon dioxide is a greenhouse gas and may contribute to climate change.

Carbon dioxide prevents heat loss from earth. This causes to warm up the atmosphere and may lead to climate change.

e) Hydrogen and Water

2.25 describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron

Magnesium

With dilute hydrochloric acid

Mg(s) + HCl(aq) ==> MgCl2(aq) + H2(g)

When magnesium react with dilute hydrochloric acid, it forms magnesium chloride and hydrogen gas.

With dilute sulphuric acid

Mg(s) + H2SO4(aq) ==> MgSO4(aq) + H2(g)

When magnesium react with dilute sulphuric acid, it forms magnesium sulphate and hydrogen gas.

Aluminium

With dilute hydrochloric acid

Al(s) + HCl(aq) ==> AlCl3(aq) + H2(g)

When aluminium react with dilute hydrochloric acid, it forms aluminium trichloride and hydrogen gas.

With dilute sulphuric acid

Al(s) + H2SO4(aq) ==> Al2(SO4)3(aq) + H2(g)

When aluminium react with dilute sulphuric acid, it forms aluminium sulphate and hydrogen gas.

Iron

With dilute hydrochloric acid

Fe(s) + HCl(aq) ==> FeCl2(aq) + H2

When Iron react with hydrochloric acid, it forms Iron(III) chloride and hydrogen gas.

With dilute sulphuric acid

Fe(s) + H2SO4(aq) ==> FeSO4(aq) + H2

When iron react with sulphuric acid, it forms iron sulphate and hydrogen gas.

2.26 describe the combustion of hydrogen

2H2(g) + O2(g) ==> 2H2O(g)

Hydrogen gas burns in air or oxygen to form water(steam).

When hydrogen is pure, hydrogen burns quietly. However, when it is mixed with air or oxygen first, an explosion will occur if a spark or flame is applied.

Testing hydrogen

When a lighted splint is held at the mouth of the test tube that contains hydrogen gas, it will make a popping sound.

2.27 describe the use of anhydrous copper(II) sulfate in the chemical test for water

Water turns white anhydrous copper(II) sulphate blue. Anhydrous copper(II) sulphate lacks water of crystallisation. Dropping water in it replaces the water of crystallisation and therefore colour changes.

CuSO4(s) + 5H2O(l) ==> CuSO4x5H2O(s)

2.28 describe a physical test to show whether water is pure.

If water exactly freezes at 0oC and boils at exactly 100oC at 1 atmospheric pressure, then the water is pure.

f) Reactivity series

2.29 understand that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold

2.30 describe how reactions with water and dilute acids can be used to deduce the following order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper

Potassium, sodium, lithium, calcium react with water and acids.

Magnesium, aluminium, zinc ,iron ,tin, lead can reac Copper can’t react with any of them.

The more the vigorous the element is in the reaction, the more reactive the element is and the more things they can react with.

2.31 deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions

If metal is reactive than the metal in metal oxide, the metal will be displaced and reaction will occur in water. If metal is less reactive than the metal in metal oxide, no reaction will occur.

In this way we can work out their position in reactivity series.

2.32 understand oxidation and reduction as the addition and removal of oxygen respectively

Oxidation: Oxidation means addition of oxygen to an element Reduction: Reduction means removal of oxygen from an element

2.33 understand the terms redox, oxidising agent, reducing agent

Redox reaction: If in a reaction, both oxidation and reduction take place, it is called redox reaction.

Oxidizing agent: Oxidising agent is a substance that oxidizes others but reduce itself from the reaction. Reducing agent: Reducing agent is a substance that reduces others but oxides itself from the reaction.

Oxidising agents

Reducing agents

Bromine

carbon

Chlorine

Carbon monoxide

Concentrated sulphuric acid

Hydrogen

Nitric acid

Hydrogen sulphide

Oxygen

Metals

Potassium manganate(VII)

Potassium iodide

Potassium dichromate

Sulphur dioxide

Hydrogen peroxide

Ammonia

2.34 describe the conditions under which iron rusts

What is rusting?

When an object made of iron is exposed to moist air for some time, a reddish-brown substance slowly forms on the surface of the metal. The substance is called rust and have the chemical name hydrated iron (II) oxide.

The process is known as rusting or corrosion of iron. 4Fe(s) + 3O2(g) + 2xH2O(l) ==> 2Fe2O3xH 2O(s)

Conditions under which iron rusts:

· Both air and water are needed for rusting to occur

· The presence of sodium chloride increases the speed of rusting.

2.35 describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising

Grease, oil, paint and plastic prevent air and or water from coming into contact with iron. This means the reaction that

rusts iron can't occur.

Galvanizing is coating in zinc. This Zinc react in the air to form ZnCO3 which prevents air and or water from coming into contact with the iron.

2.36 understand the sacrificial protection of iron in terms of the reactivity series.

Magnesium or zinc are used as sacrificial protection of iron. If you keep them beside or attach them to iron, magnesium or zinc will corrode instead of iron. Because they are more reactive than iron. As long as magnesium or zinc is present, iron will not rust.

g) Test for ions and gases

2.37 describe tests for the cations:

i. Li+, Na+, K+, Ca2+ using flame tests

  • Lithum: red
  • Sodium: orange (so strong can mask other colours)
  • Potassium: lilac
  • Calcium: brick red

ii. NH4+, using sodium hydroxide solution and identifying the ammonia evolved

(NH4)2 SO4 + NaOH ==> NH3 + Na2SO4 + H2O

Ammonium sulphate + sodium hydroxide gives ammonia gas, sodium sulphate and water. Ammonia is a pungent smelling gas which turns red litmus paper blue.

iii. Cu2+, Fe2+ and Fe3+, using sodium hydroxide solution

  • Copper(ii) sulphate + sodium hydroxide > blue precipitate
  • Iron(ii) sulphate + sodium hydroxide > green precipitate
  • Iron(iii) sulphate + sodium hydroxide > brown precipitate

2.38 describe tests for the anions:

i. Cl - , Br - and I - , using dilute nitric acid and silver nitrate solution

  • Chloride ions + nitric acid + silver nitrate ==> white precipitate (silver chloride)
  • Bromide ions + nitric acid + silver nitrate ==> cream precipitate (silver bromide)
  • Iodide ions + nitric acid + silver nitrate ==> yellow precipitate (silver iodide)

ii. SO 4 2- , using dilute hydrochloric acid and barium chloride solution

  • SO4(2-) + HCl + Ba(2+) ==> white precipitate (barium sulphate)

iii. CO 3 2- , using dilute hydrochloric acid and identifying the carbon dioxide evolved

  • Carbonate + acid ==> salt + water + carbon dioxide
  • Carbon dioxide produced will turn lime water cloudy

2.39 describe tests for the gases:

i. hydrogen

  • Make popping sound when hydrogen is burned.

ii. oxygen

  • Relights a glowing splint.
iii. carbon dioxide
  • Make lime water cloudy

iv. ammonia

  • Have a pungent smell
  • Turns red litmus paper blue

v. chlorine.

  • Turns damp litmus paper white.