Chemistry Igcse

Front 1

Solvents for chromatography

Back 1

Water and organic solvents (ethanol, propanol). Organic solvents are used because they dissolve substantes insoluble in water.

Front 2

Precautions for using organic solvents for chromatography

Back 2

Use lid to stop it from evaporating

Front 3

Industrial uses of isotopes

Back 3

U-235 used in nuclear power stations

Detecting level of liquid in container

Detecting leaks in underground water pipes

Controlling thickness of sheets of plastic, foil, paper, etc.

Front 4

Medical uses of isotopes

Back 4

Cobalt-60 used to treat internal cancer tumours.

Strontium-90 and Phosphorus-32 used to treat skin cancer tumours.

Gamma rays used to sterilize surgical equipment.

Iodine-131 used to check for dysfunctions in the thyroid gland

Front 5

Alkali metals react with air to form ________ and with water to form _________.

Back 5

Air=oxides

Water=hydroxides + hydrogen

Front 6

Name three trends for group 1 metals

Back 6

1. Reactivity increases down the group

2. Increase in density down the group

3. Decrease in melting point only (not boiling point)

Front 7

Name some physical properties of alkali metals

Back 7

Soft (can be cut with a knife)

Shiny when freshly cut

Low density

Low melting points

Front 8

Name some chemical properties of alkali metals

Back 8

React with air to form oxides

Have a valency of 1

React with water to form hydroxides and hydrogen

Front 9

Reaction of Mg with steam

Back 9

Mg + H²O(g) -----> MgO + H²

Front 10

State the three rules for electrolysis

Back 10

1. The positive ions will follow the series

2. For concentrated solutions, the halogens will be discharged first

3. For dilute solutions, and for solutions of sulfates and nitrates, the hydroxide ions will be discharged first.

Front 11

Where is V2O5 (vanadium pentoxide)

Back 11

In the production of sulfuric acid as a catalyst

Front 12

Where is finely divided iron used?

Back 12

In the Haber process, for the production of ammonia as a catalyst.

Front 13

State general properties of halogens.

Back 13

Reactivity decreases down the group

Gradual change in the states of matter

Colour darkens down the group

All have a valency of one

Mp and Bp increase down the group

Toxic

Poor conductors of heat

Poor conductors of electricity

Front 14

Map out the limestone cycle

Back 14

Front 15

Bronze is made of?

Back 15

Copper and tin

Front 16

Brass is made of?

Back 16

Copper and Zinc

Front 17

Give three reasons why aluminium is used for cables

Back 17

Low density (prevents sagging)

Good conductor of electricity

Resistant to corrosion

Front 18

How can leakage of power from cables be prevented?

Back 18

Using ceramic plates between the cable and the pylon

Front 19

How is electricity conducted in metals and in ionic solutions?

Back 19

Electrons carry the current in metals. Ions carry the charge in solutions.

Front 20

Where does oxidation and reduction occur in electrolysis?

Back 20

Reduction at cathode where metal is formed as ions gain electrons. Oxidation at anode where non metals are formed as ions lose electrons.

Front 21

Why do we electroplate?

Back 21

For attractive appearances

For protective coating (eg to prevent rusting)

Front 22

Three precautions for electroplating.

Back 22

-object must be clean and free of grease otherwise impurities will interfere with the coating giving a rough texture that will flake off

-object must be rotated to get the coating evenly on all sides

-must be carried out in careful control: temperature, current and the concentration of the electrolyte must be precise otherwise the layer will form too fast and flake off.

Front 23

What is the catalyst used for the decomposition of hydrogen peroxide? Write an equation for the reaction.

Back 23

Manganese (IV) oxide.

2H2O2(l) ----> 2H2O(l) + O2(g)

Front 24

Name the catalyst for the following industrial processes:

1. Manufacture of nitric acid (oxidation of ammonia)

2. Methanol to hydrocarbon conversion

Back 24

1. Platinum-rhodium

2. Zeolite

Front 25

How do catalysts work?

Back 25

A catalyst increases the rate of reaction by reducing the amount of energy needed for the bonds to break. This reduces the activation energy of the reaction and makes sure there are more frequent collisions.

Front 26

Give three equations for photography from silver bromide and provide an explanation for each.

Back 26

2AgBr ----> 2Ag + Br2

Silver bromide crystals decompose in the presence of sunlight. The precipitates (AgBr) darken in the presence of sunlight. This is a redox reaction:

- bromide ions lose electrons (they are oxidized) 2Br- -2e ----> Br2

-silver ions gain electrons (they are reduced) Ag+ +e -----> Ag

So, in photography, silver ions are converted into silver atoms. This is the basis of photography.

2Ag

Front 27

FLAME TESTS FOR:

• Li

• Na

• K

Back 27

Lithium is red.

Sodium is yellow.

Potassium is lilac.

Front 28

Colour of transition metals in solution:

• Copper

• Chromium

• Cobalt

• Chromate

• Dichromate

• Managanate

Back 28

Cooper is blue.

Chromium is green.

Cobalt is pink.

Chromate is yellow and dichromate is orange.

Managanate is purple.

Front 29

Five uses of limestone.

Back 29

1. Used with sand and NaCO3 to make glass.

2. Used to make cement.

3. Used for neutralization of acidified lakes and soils.

4. Used to make concrete.

5. Used to remove impurities in the production of steel.

Front 30

What type of reaction forms the basis of electrolysis and what type forms the basis of cells? Explain why.

Back 30

Electrolysis is endothermic: energy is being supplied. The energy is converted from electrical to chemical.

Simple cells involve exothermic relations: energy is being produced. We use chemical energy to produce electricity.

Front 31

Explain the principle of simple cells.

Back 31

Zn loses electrons (oxidation) to form ions therefore becoming the negative electrode.

Zn - 2e ----> Zn+²

Cu ions gain electrons (reduction) to become copper atoms.

Cu²+(aq) + 2e ----> Cu(s)

Overall reaction:

Zn(s) + Cu²+ (aq) ----> Zn²+(aq) + Cu(s)

Front 32

How does a fuel cell work?

Back 32

Forms the basis of an electrochemical cell. It consists of an oxidant which reacts with a fuel in a redox reaction to release energy.

Front 33

Explain the functioning of a fuel cell using hydrogen as fuel.

Back 33

When oxidized, hydrogen forms water. The fuel cell consists of two electrodes sandwiched around an electrolyte. Hydrogen is fed to the anode and Oxygen is fed to the cathode. A catalyst activates the hydrogen atoms to separate into protons (H+) and electrons. The electrons flow through the external circuit, creating a current while the protons travel through the electrolyte, to the cathode, where they unite with oxygen and the electrons to produce water.

A: 2H2 + 4e ----> 4H+

C: O2 + 4H+ +4e ----> 2H2O

OVERALL:

2H2 + O2 -----> 2H2O

Front 34

What are the products of complete combustion of hydrocarbons?

Back 34

Carbon dioxide and water. It is al also an exothermic reaction.

Front 35

Pros and cons of fuel cells with hydrogen as a fuel.

Back 35

PROS:

- Very clean production of energy.

- Less opportunity of losing potentially useful energy as there are few staged of energy production.

- No carbon dioxide emissions

- Nontoxic

CONS:

- Large fuel tank required

- Currently expensive

- Difficult to store in one place

Front 36

Rules for assigning oxidation numbers.

Back 36

1. Ox no for any uncombined element is 0. (Eg O2, Xe, C)

2. Ox no of a monoatomic ion equals to it's charge.

3. Ox no of fluorine in a compound is always -1.

4. Oxygen has ox no of -2 unless combined with F when it is +2 or when in peroxide when it is -1.

5. Ox no of hydrogen in most compounds is +1 (HCl) unless combined with metal where it is -1 (NaH)

6. In compounds, group 1, 2 and Aluminium elements have ox nos 1, 2 and 3 respectively.

7. Sum of ox nos in a neutral compound is 0.

8. Sum of ox nos in polyatomic ions is equal to it's charge.

Front 37

Define oxidation, and reduction

Back 37

Oxidation is loss of electrons. Reduction is gain.

Oxidation is gain of oxygen, reduction is loss.

Oxidation is increase in oxidation number, reduction is decrease.

Oxidation is loss of hydrogen, reduction is gain.

Front 38

Define oxidising agents. Give examples.

Back 38

Cause oxidation reactions to take place. Gain electrons from other atoms or ions and are therefore themselves reduced. Eg: potassium manganate, potassium dichromate, hydrogen peroxide and oxygen.

Front 39

Define reducing agents. Give examples.

Back 39

Cause reduction reactions to take place. Lose/donate electrons to other substances and are thus oxidized themselves. Eg: Carbon, hydrogen, carbon monoxide.

Front 40

Test for oxidizing agents.

Back 40

Addition of KI (Potassium Iodide)

Colourless to yellow brown.

2I- -2e ----> I2

Front 41

Test for reducing agents.

Back 41

Addition of acidified potassium manganate: purple to colourless.

Addition of acidified potassium dichromate: orange to green.

Front 42

Chemical test for water

Back 42

CuSO4.5H2O <-----> CuSO4 + 5H2O

(Blue is hydrated) to (white is anhydrous)

Front 43

Define reversible reactions.

Back 43

Reactions in which products can be converted back into their original state as a reactant.

Front 44

Define Le Chatelier's Principle.

Back 44

When a system in dynamic equilibrium is stressed, the system moves so as to oppose gbr change in order to restore the equilibrium.

Front 45

Factors affecting equilibrium positions.

Back 45

Concentration.

Pressure (only for systems in gaseous states): increasing the pressure shifts the equilibrium to the side with fewer moles.

Temperature: increase in temperature causes the system to favor an endothermic reaction. Decrease in temperature favours an exothermic reaction.

Front 46

State the conditions for the Haber process. Write the equation.

Back 46

450°C temperature.

200 atm.

Catalyst of finely divided iron.

2N + 3H2 ----> 2NH3

(Forward reaction is exothermic, so a lower temperature is preferred with increased pressure).

Front 47

State the equation for the Contact process. Write the conditions.

Back 47

450°C temperature

1 atm

Vanadium pentoxide as catalyst.

Exothermic reaction is the forward reaction so a lower temperature is preferred.

2SO2 + O2 ----> 2SO3

Front 48

Three uses of sulfur.

Back 48

• To vulcanize rubber (in order to make it harder and more elastic)

• As a sterilizing agent

• To manufacture matches, fireworks.

Front 49

Three uses of sulfur dioxide.

Back 49

• Used in the Contact process for the production of sulfuric acid.

• Used as a bleach in making paper from wood pulp.

• Used as a food preservative by killing bacteria.

Front 50

Three uses of sulfuric acid.

Back 50

• Used in the manufacture of paint and dyes.

• Used in producing soaps and detergents.

• Used in dyes, fabrics and polymers.

Front 51

Explain the manufacture of H2SO4.

Back 51

- sulfur burnt in air to produce sulfur dioxide. S + O2 ---> SO2

- sulfur further oxidized to sulfur trioxide in contact process.

- SO3 is dissolved in conc H2SO4. The solution formed means the acid can be transported in concentrated form and diluted on site. SO3 + H2SO4 ----> H2S2O7 (oleum)

- oleum is then diluted to make sulfuric acid. H2S2O7 + H2O ----> 2H2SO4

Front 52

Why isn't sulfur trioxide directly dissolved in water to produce sulfuric acid?

Back 52

The reaction is extremely exothermic and so very violent producing an acid mist which is detrimental to the environment.

Front 53

Uses of chlorine.

Back 53

• to kill bacteria in water

• in making bleach

• to make plastic PVC

Front 54

Uses of fluorine

Back 54

• in toothpaste

• in non stick coating for pans

Front 55

Uses of bromine

Back 55

• in pesticides

• in photography from it's salts

• flame retardants

• petrol additives

Front 56

Uses of iodine

Back 56

• photography paper and film

• as an antiseptic

Front 57

How is cement made? Also concrete? Mortar?

Back 57

Powdered limestone is mixed with clay in a rotary kiln to make cement. Cement is mixed with aggregate to make concrete. Ca(OH)2 is mixed with sand to make mortar.

Front 58

Explain the manufacturing of hydrogen through methane and water.

Back 58

1. DESULPHERIZATION: methane is purified to remove sulfurous impurities that could poison the catalyst.

2. STEAM REFORMER: methane reacts with steam to form CO and H2 in 700°C, 30 atm and nickel catalyst.

3. SHIFT REACTOR: CO reacts with steam to form CO2 and H2.

4. ABSORBER: mixture of CO2 and H2 passes through a basic solution which removes CO2.

5. HYDROGEN IS SERVED.

Front 59

What do displacement reactions for ammonia produce?

Back 59

When ammonia salts displace by more reactive bases, they produce ammonia and water.

Front 60

Uses of ammonia.

Back 60

• used to manufacture fertilizers

• used as a household cleansing agent

• used to produce nitric acid

Front 61

What type of reaction is breaking and forming bonds?

Back 61

Breaking binds takes in energy and so is endothermic. Forming binds gives out energy and is exothermic.

Front 62

How is lime produced?

Back 62

By the decomposition of calcium carbonate in a lime kiln. An open system where equilibrium is not possible as the carbon dioxide gas escapes. Carrying the reaction out in a closed container would make it reversible and not that much of limestone will be decomposed.

Front 63

Uses of aluminium

Back 63

In cables.

In aluminium foil

In buildings and windows