SpecPoint Archives - Page 6 of 8 - TutorMyself Chemistry

2:11 describe the combustion of elements in oxygen, including magnesium, hydrogen and sulfur

Magnesium reacts with oxygen producing a bright white flame leaving behind a white ash of magnesium oxide.

magnesium + oxygen → magnesium oxide

2Mg (s) + O₂ (g) → 2MgO

MgO is a base, which can react with an acid to give a salt and water.

Hydrogen reacts with oxygen in an explosive reaction. This is the basis of the ‘squeak pop’ test for hydrogen in test tube. With larger quantities of hydrogen this explosion can be dangerous.

hydrogen + oxygen → water

2H₂ (g) + O₂ (g) → 2H₂O (l)

Sulfur reacts with oxygen producing a blue flame.

sulfur + oxygen → sulfur dioxide

S (s) + O₂ (g) → SO₂ (g)

When sulfur dioxide (SO₂) dissolves in water it forms an acidic solution of sulfurous acid:

SO₂ (g) + H₂O (l) → H₂SO₃ (aq)

Hydr0Gen2022-11-30T14:20:27+00:00Categories: (c) Gases in the atmosphere, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:11, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Key Reactions|

2:12 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates, including copper(II) carbonate

thermal decomposition is the process of breaking down by heating.

On heating metal carbonates thermal decompose into metal oxides and carbon dioxide.

Observation: green powder (CuCO₃) changes to a black powder (CuO)

Hydr0Gen2022-11-28T13:09:36+00:00Categories: (c) Gases in the atmosphere, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:12, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Key Reactions, Topic: Specification experiments|

2:13 know that carbon dioxide is a greenhouse gas and that increasing amounts in the atmosphere may contribute to climate change

Carbon dioxide (CO₂) is a greenhouse gas.

It absorbs infra-red radiation and therefore warms the atmosphere. This leads to global warming.

This may cause climate change.

Hydr0Gen2022-11-30T14:20:45+00:00Categories: (c) Gases in the atmosphere, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:13, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Alkanes & Crude Oil|

2:14 Practical: determine the approximate percentage by volume of oxygen in air using a metal or a non-metal

The following 3 experiments can be used to determine that oxygen (O₂) makes up approximately 20% by volume of air.

Copper

The copper is in excess and uses up the oxygen to form copper oxide (CuO).

All the oxygen in the air is therefore used up, and so the volume of the air decreases by about 20% (the percentage of oxygen in air).

Iron

The iron reacts with the oxygen in the air (rusting).

As long as the iron, oxygen and water are all in excess, the total volume of air enclosed by the apparatus decreases by about a fifth (20%) over several days.

Phosphorus

The phosphorus is lit with a hot wire.

It reacts with the oxygen in the air and causes the water level in the bell jar to rise by about 20%.

Hydr0Gen2023-09-18T11:06:27+00:00Categories: (c) Gases in the atmosphere, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:14, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Simple Molecules & Covalent Bonding, Topic: Specification experiments|

2:15 understand how metals can be arranged in a reactivity series based on their reactions with: water and dilute hydrochloric or sulfuric acid

Some metals are more reactive than others.

The order of reactivity can be determined by adding acid to different metals and observing the rate of reaction.

For example, when hydrochloric acid is added to iron (Fe) then bubbles of hydrogen are produced slowly. However, if the same acid is added to zinc (Zn) then bubbles will be produced more quickly. This tells us that zinc is more reactive than iron.

Instead of using acid, water can be used to test the relative reactivity of metals. However, many metals are too low in the reactivity series to react with water

Hydr0Gen2023-01-10T12:30:26+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:15, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Metal Reactivity & Halogens, Topic: Metals (D)|

2:16 understand how metals can be arranged in a reactivity series based on their displacement reactions between: metals and metal oxides, metals and aqueous solutions of metal salts

A metal will displace another metal from its oxide that is lower in the reactivity series. For example, a reaction with magnesium and copper (II) oxide will result in the magnesium displacing the copper from its oxide:

A metal will also displace another metal from its salt that is lower in the reactivity series. For example, the reaction between zinc and copper (II) sulfate solution will result in zinc displacing the copper from its salt:

The blue colour of the copper (II) sulfate solution fades as colourless zinc sulfate solution is formed.

Hydr0Gen2023-01-10T12:30:47+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:16, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Electrolysis, Topic: Metal Reactivity & Halogens, Topic: Metals (D)|

2:17 know the order of reactivity of these metals: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver, gold

A more reactive metal will displace a less reactive metal.

In addition a more reactive metal will react more vigorously than a less reactive metal.

For example, potassium takes a shorter time to react than sodium:

Hydr0Gen2023-01-10T12:31:04+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:17, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Metal Reactivity & Halogens, Topic: Metals (D)|

2:18 know the conditions under which iron rusts

Iron rusts when oxygen and water are present.

Reaction:

Hydr0Gen2023-01-10T12:31:16+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:18, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Metal Reactivity & Halogens, Topic: Metals (D)|

2:19a understand how the rusting of iron may be prevented by: barrier methods, galvanising

Barrier Methods: Rusting may be prevented by stopping the water and oxygen getting to the iron with a barrier of grease, oil, paint or plastic.

Galvanising: (coating in zinc) also prevents water and oxygen getting to the iron, but with galvanising even if the barrier is broken the more reactive zinc corrodes before the less reactive iron. During the process, the zinc loses electrons to form zinc ions.

Sam Mahal2022-12-12T11:13:19+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:19a, Chemistry, GCSE, GCSE_Chemistry_Single, SpecPoint, Topic: Metal Reactivity & Halogens|

2:19 understand how the rusting of iron may be prevented by: barrier methods, galvanising and sacrificial protection

Barrier Methods: Rusting may be prevented by stopping the water and oxygen getting to the iron with a barrier of grease, oil, paint or plastic.

Galvanising: (coating in zinc) also prevents water and oxygen getting to the iron, but with galvanising even if the barrier is broken the more reactive zinc corrodes before the less reactive iron. During the process, the zinc loses electrons to form zinc ions.

Sacrificial Protection: Zinc blocks are attached to iron boat hulls and underground pipelines to act as sacrificial anodes. Zinc is more reactive than iron, so oxygen in the air reacts with the zinc to form a layer of zinc oxide instead of the iron.

Hydr0Gen2023-01-10T12:31:29+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:19, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Metal Reactivity & Halogens, Topic: Metals (D)|

2:20 in terms of gain or loss of oxygen and loss or gain of electrons, understand the terms: oxidation, reduction, redox, oxidising agent, reducing agent, in terms of gain or loss of oxygen and loss or gain of electrons

Oxidation

Oxidation is the loss of electrons. For example a sodium atom (Na) loses an electron to become a sodium ion (Na⁺). Another example is a chloride ion (Cl⁻) losing an electron to become a chlorine atom (Cl).
Another definition of oxidation is the gain of oxygen. For example if carbon combines with oxygen to form carbon dioxide, the carbon is being oxidised.

Reduction

Reduction is the gain of electrons. For example a sodium ion (Na⁺) gains an electron to become a sodium atom (Na). Another example is a chlorine atom (Cl) gaining an electron to become a chloride ion (Cl⁻).
Another definition of reduction is the loss of oxygen. For example when aluminium oxide is broken down to produce aluminium and oxygen, the aluminium is being reduced.

Redox: A reaction involving oxidation and reduction.

A good way to remember the definitions of oxidation and reduction in terms of electrons is:

OILRIG : Oxidation Is the Loss of electrons and Reduction Is the Gain of electrons

Oxidising agent: A substance that gives oxygen or removes electrons (it is itself reduced).

Reducing agent: A substance that takes oxygen or gives electrons (it is itself oxidised).

Hydr0Gen2023-04-19T10:58:30+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:20, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Electrolysis, Topic: Metal Reactivity & Halogens, Topic: Metals (D)|

2:21 practical: investigate reactions between dilute hydrochloric and sulfuric acids and metals (e.g. magnesium, zinc and iron)

Metals which are above hydrogen in the reactivity series will react with dilute hydrochloric or sulfuric acid to produce a salt and hydrogen.

metal + acid → salt + hydrogen

For example:

magnesium + hydrochloric acid → magnesium chloride + hydrogen

Mg (s) + 2HCl (aq) → MgCl₂ (aq) + H₂ (g)

This is a displacement reaction.

There is a rapid fizzing and a colourless gas is produced. This gas pops with a lighted splint, showing the gas is hydrogen.

The reaction mixture becomes warm as heat is produced (exothermic).

The magnesium disappears to leave a colourless solution of magnesium chloride.

If more reactive metals are used instead of magnesium the reaction will be faster so the fizzing will be more vigorous and more heat will be produced.

Hydr0Gen2023-01-10T12:32:04+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:21, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Metal Reactivity & Halogens, Topic: Metals (D), Topic: Specification experiments|

2:22 (Triple only) know that most metals are extracted from ores found in the Earth’s crust and that unreactive metals are often found as the uncombined element

Most metals are found in the Earth’s crust combined with other elements. Such compounds are found in rocks called ore, rocks from which it is worthwhile to extract a metal.

A few very unreactive metals, such as gold, are found native which means they are found in the Earth’s crust as the uncombined element.

Hydr0Gen2019-05-13T15:38:43+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:22, Chemistry, ChemistrySpecPointsNew2019, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Electrolysis, Triple|

2:23 (Triple only) explain how the method of extraction of a metal is related to its position in the reactivity series, illustrated by carbon extraction for iron and electrolysis for aluminium

Extraction of a metal from its ore typically involves removing oxygen from metal oxides.

If the ore contains a metal which is below carbon in the reactivity series then the metal is extracted by reaction with carbon in a displacement reaction.

If the ore contains a metal which is above carbon in the reactivity series then electrolysis (or reaction with a more reactive metal) is used to extract the metal.

Hydr0Gen2019-05-13T15:38:34+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:23, Chemistry, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Electrolysis, Triple|

2:24 (Triple only) be able to comment on a metal extraction process, given appropriate information

Extraction of a metal from its ore typically involves removing oxygen from metal oxides.

If the ore contains a metal which is below carbon in the reactivity series then the metal is extracted by reaction with carbon in a displacement reaction.

If the ore contains a metal which is above carbon in the reactivity series then electrolysis (or reaction with a more reactive metal) is used to extract the metal.

Hydr0Gen2019-05-13T15:38:25+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:24, Chemistry, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Electrolysis, Triple|

2:25 (Triple only) explain the uses of aluminium, copper, iron and steel in terms of their properties the types of steel will be limited to low-carbon (mild), high-carbon and stainless

Aluminium
Use	Property
Aircrafts and cans	Low density / resists corrosion
Power cables	Conducts electricity / ductile
Pots and pans	Low density / strong (when alloyed) / good conductor of electricity and heat

Aluminium resists corrosion because it has a very thin, but very strong, layer of aluminium oxide on the surface.

Copper
Use	Property
Electrical wires	very good conductor of electricity and ductile
Pots and pans	very good conductor of heat / very unreactive / malleable
Water pipes	unreactive / malleable
Surfaces in hospitals	antimicrobial properties / malleable

Iron
Use	Property
Buildings	Strong
Saucepans	Conducts heat / high melting point / malleable

Steel
Type of steel	Iron mixed with	Some uses
Mild steel	up to 0.25% carbon	nails, car bodies, ship building, girders
High-carbon steel	0.6%-1.2% carbon	cutting tools, masonry nails
Stainless steel	Chromium (and nickel)	cutlery, cooking utensils, kitchen sinks

Mild steel is a strong material that can easily be hammered into various shapes (malleable). It rusts easily.

High-carbon steel is harder than mild steel but more brittle (not as malleable).

Stainless steel forms a strong, protective oxide layer so is very resistant to corrosion.

Hydr0Gen2019-06-21T14:45:07+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:25, Chemistry, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Structure & Bonding, Triple|

2:26 (Triple only) know that an alloy is a mixture of a metal and one or more elements, usually other metals or carbon

An alloy is a mixture of a metal with, usually, other metals or carbon.

For example, brass is a alloy of copper and zinc, and steel is an alloy of iron and carbon.

Hydr0Gen2019-06-21T14:45:47+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:26, Chemistry, ChemistrySpecPointsNew2019, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Structure & Bonding, Triple|

2:27 (Triple only) explain why alloys are harder than pure metals

Alloys are harder than the individual pure metals from which they are made.

In an alloy, the different elements have slightly different sized atoms. This breaks up the regular lattice arrangement and makes it more difficult for layers of ions to slide over each other.

Hydr0Gen2019-06-21T14:46:10+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:27, Chemistry, ChemistrySpecPointsNew2019, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Structure & Bonding, Triple|

2:28a describe the use of litmus to distinguish between acidic and alkaline solutions

Indicator	Colour in acidic solution	Colour in alkaline solution
Litmus	Red	Blue

Sam Mahal2022-12-12T11:16:42+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:28a, Chemistry, GCSE, GCSE_Chemistry_Single, SpecPoint, Topic: Key Reactions|

2:28 describe the use of litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions

Indicators are substances which change colour in response to a change in pH (acid or alkali).

Indicator	Colour in acidic solution [H⁺]	Colour in alkaline solution [OH^-]
Litmus	Red	Blue
Methyl orange	Red	Yellow
Phenolphthalein	Colourless	Pink

Methyl orange is orange in a neutral solution

Hydr0Gen2023-01-13T22:06:03+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:28, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Key Reactions|

2:29 understand how to use the pH scale, from 0–14, can be used to classify solutions as strongly acidic (0–3), weakly acidic (4–6), neutral (7), weakly alkaline (8–10) and strongly alkaline (11–14)

The pH scale ranges from 0 to 14, and tells you how acidic or how alkaline a solution is.

	strongly acidic	weakly acidic	neutral	weakly alkaline	strongly alkaline
pH	0-3	4-6	7	8-10	11-14

Hydr0Gen2023-01-10T11:38:25+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:29, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Key Reactions|

2:30 describe the use of Universal Indicator to measure the approximate pH value of an aqueous solution

An indicator is a substance that has more than one colour form depending on the pH.

Universal indicator is a mixture of different dyes which change colour in a gradual way over a range of pH.

Hydr0Gen2023-01-10T11:39:19+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:30, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Key Reactions|

2:31 know that acids in aqueous solution are a source of hydrogen ions and alkalis in a aqueous solution are a source of hydroxide ions

An acid is source of hydrogen ions (H⁺).

An alkali is source of hydroxide ions (OH^–).

Hydr0Gen2023-01-10T11:39:28+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:31, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Key Reactions|

2:32 know that bases can neutralise acids

Metal oxides, metal hydroxides and ammonia (NH₃) are called bases.

Bases neutralise acids by combining with the hydrogen ions in them.

The key reaction is:

acid + base → salt + water

An example of this is:

sulfuric acid + copper oxide → copper sulfate + water

H₂SO₄ + CuO → CuSO₄ + H₂O

Hydr0Gen2023-09-18T11:07:11+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:32, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Key Reactions|

2:33 (Triple only) describe how to carry out an acid-alkali titration

Titration is used to find out precisely how much acid neutralises a certain volume of alkali (or vice versa).

The diagram shows the titration method for a neutralisation reaction between hydrochloric acid and sodium hydroxide, using phenolphthalein as an indicator. The indicator changes colour when neutralisation occurs.

The conical flask is swirled to mix the solutions each time alkali is added. When reading the burette it is important to be aware that the numbers on the scale increase from top to bottom. Readings are usually recorded to the nearest 0.05cm³ so all readings should be written down with 2 decimal places. The second decimal place is given as a ‘0’ if the level of the solution is on a line, or ‘5’ if it is between the lines. The volume of alkali added is calculated by subtracting the final reading from the initial reading. Various indicators can be used such as phenolphthalein or methyl orange. However universal indicator should not be used since it has a wide range of colours rather than one specific colour change so it would be unclear when the precise endpoint of titration was achieved.

This process is repeated a number of times. The first time it is done roughly to get a good approximation of how much alkali needs to be added. On subsequent attempts, the alkali is added very slowly when approaching the correct volume.

Hydr0Gen2019-05-13T15:37:44+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:33, Chemistry, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Key Calculations, Topic: Specification experiments, Triple|

2:34 know the general rules for predicting the solubility of ionic compounds in water: common sodium, potassium and ammonium compounds are soluble, all nitrates are soluble, common chlorides are soluble, except those of silver and lead(II), common sulfates are soluble, except for those of barium, calcium and lead(II), common carbonates are insoluble, except for those of sodium, potassium and ammonium, common hydroxides are insoluble except for those of sodium, potassium and calcium (calcium hydroxide is slightly soluble)

Salt	Solubility	Exceptions
sodium (Na⁺), potassium (K⁺) and ammonium (NH₄⁺)	soluble	none
nitrates (NO₃^-)	soluble	none
chlorides (Cl^-)	soluble	silver chloride (AgCl) and lead (II) chloride (PbCl₂)
sulfates (SO₄^2-)	soluble	barium sulfate (BaSO₄), calcium sulfate (CaSO₄) and lead (II) sulfate (PbSO₄)
carbonates (CO₃^2-)	insoluble	sodium carbonate (Na₂CO₃), potassium carbonate (K₂CO₃) and ammonium carbonate ((NH₄)₂CO₃)
hydroxides (OH^-)	insoluble	sodium hydroxide (NaOH), potassium hydroxide (KOH) and calcium hydroxide (Ca(OH)₂) (calcium hydroxide is slightly soluble)

Hydr0Gen2023-01-10T11:40:19+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:34, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D)|

2:35 understand acids and bases in terms of proton transfer

An acid is a proton (H⁺) donor.

A base is a proton (H⁺) acceptor.

A proton is the same as a hydrogen ion. A good way to think about that is to realise that a hydrogen atom is just one proton and zero neutrons surrounded by only one electron. If that atom becomes an ion by the removal of the electron, then only one proton is left.

When sulfuric acid reacts with copper (II) oxide (CuO):

Cu²⁺O²⁻ (s) + H₂SO₄ (aq) → Cu²⁺ (aq) + SO₄²⁻ (aq) + H₂O (l)

H₂SO₄ is an acid. It donates protons (H⁺) to CuO, the base.

Hydr0Gen2023-01-10T11:41:24+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:35, Chemistry, ChemistrySpecPointsNew2019, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D)|

2:36 understand that an acid is a proton donor and a base is a proton acceptor

An acid is a proton donor.

A base is a proton acceptor.

A proton is the same as a hydrogen ion. A good way to think about that is to realise that a hydrogen atom is just one proton and zero neutrons surrounded by only one electron. If that atom becomes an ion by the removal of the electron, then only one proton is left.

Hydr0Gen2023-01-10T11:41:38+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:36, Chemistry, ChemistrySpecPointsNew2019, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D)|

2:37 describe the reactions of hydrochloric acid, sulfuric acid and nitric acid with metals, bases and metal carbonates (excluding the reactions between nitric acid and metals) to form salts

Acid reactions summary

alkali + acid → water + salt

base + acid → water + salt

carbonate + acid → water + salt + carbon dioxide

metal + acid → salt + hydrogen

To assist remembering this list, many pupils find it useful to remember this horrid looking but very effective mnemonic:

AAWS

BAWS

CAWS CoD

MASH

Acids are a source of hydrogen ions (H⁺) when in solution. When the hydrogen in an acid is replaced by a metal, the compound is called a salt. The name of the salt depends on the acid used. For example if sulfuric acid is used then a sulfate salt will be formed.

Parent acid	Formula	Salt	Formula ion
sulfuric acid	H₂SO₄	sulfate	SO₄^2-
hydrochloric acid	HCl	chloride	Cl^-
nitric acid	HNO₃	nitrate	NO₃^-

Acid + Alkali and Acid + Base

A base is a substance that can neutralise an acid, forming a salt and water only.

Alkalis are soluble bases. When they react with acids, a salt and water is formed. The salt formed is often as a colourless solution. Alkalis are a source of hydroxide ions (OH⁻) when in solution.

alkali + acid → water + salt

base + acid → water + salt

Examples of acid + alkali reactions:

sodium hydroxide + hydrochloric acid → sodium chloride + water
NaOH (aq) + HCl (aq) → NaCl (aq) + H₂O (l)

potassium hydroxide + sulfuric acid → potassium sulfate + water
2KOH (aq) + H₂SO₄ (aq) → K₂SO₄ (aq) + 2H₂O (l)

Example of an acid + base reaction:

CuO (s) + H₂SO₄ (aq) → CuSO₄ (aq) + H₂O (l)

Acid + Carbonate

carbonate + acid → water + salt + carbon dioxide

A carbonate is a compound made up of metal ions and carbonate ions. Examples of metal carbonates are sodium carbonate, copper carbonate and magnesium carbonate.

When carbonates react with acids, bubbling is observed which is the carbon dioxide being produced. If the acid is in excess the carbonate will disappear.

Examples of acid + carbonate reactions:

calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxide
CaCO₃ (s) + 2HCl (aq) → CaCl₂ (aq) + H₂O (l) + CO₂ (g)

potassium carbonate + hydrochloric acid → potassium chloride + water + carbon dioxide
K₂CO₃ (aq) + 2HCl (aq) → 2KCl (aq) + H₂O (l) + CO₂ (g)

Acid + Metal

metal + acid → salt + hydrogen

Metals will react with an acid if the metal is above hydrogen in the reactivity series.

When metals react with acids, bubbling is observed which is the hydrogen being produced. If the acid is in excess the metal will disappear.

Examples of acid + metal reactions:

magnesium + sulfuric acid → magnesium sulfate + hydrogen
Mg (s) + H₂SO₄ (aq) → MgSO₄ (aq) + H₂ (g)

aluminium + hydrochloric acid → aluminium chloride + hydrogen
2Al (s) + 6HCl (aq) → 2AlCl₃ (aq) + 3H₂ (g)

copper + hydrochloric acid → no reaction (since copper is below hydrogen in the reactivity series)

Hydr0Gen2023-01-10T11:41:49+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:37, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Key Reactions, Topic: Specification experiments|

2:38 know that metal oxides, metal hydroxides and ammonia can act as bases, and that alkalis are bases that are soluble in water

A base is a substance that neutralises an acid by combining with the hydrogen ions in them to produce water.

A base usually means a metal oxide, a metal hydroxide or ammonia.

Alkalis are bases which are soluble in water.

Some metal oxides are soluble in water and react with it to form solutions of metal hydroxides. For example:

Na₂O (s) + H₂O (l) → 2NaOH (aq)

Apart from this and other group 1 oxides (such as potassium oxide) most other metal oxides are not soluble in water.

One exception is calcium oxide which does dissolve slightly in water to form calcium hydroxide (known as limewater):

CaO (s) + H₂O (l) → Ca(OH)₂ (aq)

Ammonia is another base. Ammonia reacts with water to form ammonium ions and hydroxide ions:

NH₃ (aq) + H₂O (l) ⇋ NH₄⁺ (aq) + OH⁻ (aq)

All the solutions produced here contain hydroxide ions (OH⁻) so they are all alkalis.

Hydr0Gen2023-01-10T11:42:29+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:38, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Key Reactions|

2:39 describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an insoluble reactant

Excess Solid Method:

Preparing pure dry crystals of copper sulfate (CuSO₄) from copper oxide (CuO) and sulfuric acid (H₂SO₄)

Step	Explanation
Heat acid (H₂SO₄) in a beaker	Speeds up the rate of reaction
Add base (CuO) until in excess (no more copper oxide dissolves) and stir with glass rod	Neutralises all the acid
Filter the mixture using filter paper and funnel	Removes any excess copper oxide
Gently heat the filtered solution (CuSO₄)	To evaporate some of the water
until crystals form on a glass rod	Shows a hot saturated solution formed
Allow the solution to cool so that hydrated crystals form	Copper sulfate less soluble in cold solution
Remove the crystals by filtration	Removes crystals
Dry by leaving in a warm place	Evaporates the water

Hydr0Gen2023-01-10T11:42:40+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:39, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D)|

2:40 (Triple only) describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an acid and alkali

Titration Method:

Preparing pure dry crystals of sodium chloride (NaCl) from hydrochloric acid (HCl) and sodium hydroxide (NaOH)

Before the salt preparation is carried out using the below method, the volume of acid that exactly reacts with 25cm³ of the alkali is found by titration using methyl orange indicator.

Step	Explanation
Pipette 25cm³ of alkali (NaOH) into a conical flask	Accurately measures the alkali (NaOH)
Do not add indicator	Prevents contamination of the pure crystals with indicator
Using the titration values, titrate the known volume acid (HCl) into conical flask containing alkali	Exactly neutralises all of the alkali (NaOH)
Transfer to an evaporating basin & heat the solution	Forms a hot saturated solution (NaCl(aq))
Allow the solution to cool so that hydrated crystals form	Sodium chloride is less soluble in cold water
Remove the crystals by filtration and wash with distilled water	Removes any impurities
Dry by leaving in a warm place	Evaporates the water

(Note – This process could be reversed with the acid in the pipette and the alkali in the burette)

How to select the right method for preparing a salt:

Hydr0Gen2019-05-13T15:37:35+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:40, Chemistry, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Specification experiments, Triple|

2:41 (Triple only) describe an experiment to prepare a pure, dry sample of an insoluble salt, starting from two soluble reactants

Precipitation Method:

Preparing pure dry crystals of silver chloride (AgCl) from silver nitrate solution (AgNO₃) and potassium chloride solution (KCl)

Step	Explanation
Mix the two salt solutions together in a beaker	Forms a precipitate of an insoluble salt (AgCl)
Stir with glass rod	Make sure all reactants have reacted
Filter using filter paper and funnel	Collect the precipitate (AgCl)
Wash with distilled water	Removes any the other soluble salts (KNO₃)
Dry by leaving in a warm place	Evaporates the water

Hydr0Gen2019-05-13T15:37:24+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:41, Chemistry, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Specification experiments, Triple|

2:42 practical: prepare a sample of pure, dry hydrated copper(II) sulfate crystals starting from copper(II) oxide

Excess Solid Method:

Preparing pure dry crystals of copper sulfate (CuSO₄) from copper oxide (CuO) and sulfuric acid (H₂SO₄)

Step	Explanation
Heat acid (H₂SO₄) in a beaker	Speeds up the rate of reaction
Add base (CuO) until in excess (no more copper oxide dissolves) and stir with glass rod	Neutralises all the acid
Filter the mixture using filter paper and funnel	Removes any excess copper oxide
Gently heat the filtered solution (CuSO₄)	To evaporate some of the water
until crystals form on a glass rod	Shows a hot saturated solution formed
Allow the solution to cool so that hydrated crystals form	Copper sulfate less soluble in cold solution
Remove the crystals by filtration	Removes crystals
Dry by leaving in a warm place	Evaporates the water

Hydr0Gen2023-09-18T11:07:48+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:42, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Key Reactions, Topic: Specification experiments|

2:43 (Triple only) practical: prepare a sample of pure, dry lead(II) sulfate

Objective: prepare a pure, dry sample of lead (II) sulfate (PbSO₄).

Preparing a pure, dry sample of lead (II) sulfate (PbSO₄) from lead (II) nitrate solution (Pb(NO₃)₂) and sodium sulfate solution (Na₂SO₄).

Pb(NO₃)₂ (aq) + Na₂SO₄ (aq) → PbSO₄ (s) + 2NaNO₃ (aq)

Mix similar volumes lead nitrate solution and sodium sulfate solution in a beaker. The precise volumes do not matter since any excess will be removed later.
A white precipitate of lead (II) sulfate will form.
The reaction mixture is filtered.
The residue left on the filter paper is washed with distilled water several times to remove impurities.
The residue is then moved to a warm oven to dry.

Hydr0Gen2019-05-13T15:37:14+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:43, Chemistry, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Specification experiments, Triple|

2:44a describe tests for these gases: hydrogen, carbon dioxide

Test for hydrogen gas (H₂)

Test: use a lit splint
Result: burns with a squeaky pop

Test for carbon dioxide (CO₂)

Test: bubble through limewater
Result: limewater turns cloudy

Hydr0Gen2019-06-26T09:45:31+00:00Categories: (h) Chemical tests, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:44a, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, SubSpecPoint, Topic: Key Reactions|

2:44 describe tests for these gases: hydrogen, oxygen, carbon dioxide, ammonia, chlorine

Tests for gases

Gas	Test	Result if gas present
hydrogen (H₂)	Use a lit splint	Gas pops
oxygen (O₂)	Use a glowing splint	Glowing splint relights
carbon dioxide (CO₂)	Bubble the gas through limewater	Limewater turns cloudy
ammonia (NH₃)	Use red litmus paper	Turns damp red litmus paper blue
chlorine (Cl₂)	Use damp litmus paper	Turns damp litmus paper white (bleaches)

Hydr0Gen2023-01-10T11:43:33+00:00Categories: (h) Chemical tests, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:44, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Metal Reactivity & Halogens, Topic: Simple Molecules & Covalent Bonding, Topic: Specification experiments|

2:45 describe how to carry out a flame test

A flame test is used to show the presence of certain metal ions (cations) in a compound.

A platinum or nichrome wire is dipped into concentrated hydrochloric acid to remove any impurities.
The wire is dipped into the salt being tested so some salt sticks to the end.
The wire and salt are held in a non-luminous (roaring) bunsen burner flame.
The colour is observed.

Properties of the platinum or nichrome wire is:

Inert
High melting point

Hydr0Gen2023-02-03T11:12:43+00:00Categories: (h) Chemical tests, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:45, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Specification experiments|

2:46 know the colours formed in flame tests for these cations: Li⁺ is red, Na⁺ is yellow, K⁺ is lilac, Ca²⁺ is orange-red, Cu²⁺ is blue-green

When put into a roaring bunsen burner flame on a nichrome wire, compounds containing certain cations will give specific colours as follows.

Ion	Colour in flame test
lithium (Li⁺)	red
sodium (Na⁺)	yellow
potassium (K⁺)	lilac
calcium (Ca²⁺)	orange-red
copper (II) (Cu²⁺)	blue-green

Hydr0Gen2023-01-10T11:44:14+00:00Categories: (h) Chemical tests, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:46, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Specification experiments|

2:47 describe tests for these cations: NH₄⁺ using sodium hydroxide solution and identifying the gas evolved, Cu²⁺, Fe²⁺ and Fe³⁺ using sodium hydroxide solution

Underneath are the tests for : ammonium ion test, copper (II) ion test, iron (II) ion test, iron (III) ion test

Describe tests for the cation NH₄⁺, using sodium hydroxide solution and identifying the ammonia evolved

Describe tests for the cations Cu²⁺, Fe²⁺ and Fe³⁺, using sodium hydroxide solution

First, add sodium hydroxide (NaOH), then observe the colour:

Hydr0Gen2023-01-10T11:44:30+00:00Categories: (h) Chemical tests, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:47, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Specification experiments|

2:48a describe a test for CO₃²⁻ using hydrochloric acid and identifying the gas evolved

Describe tests for anions: Carbonate ions (CO₃^2-)

Sam Mahal2022-12-12T11:10:04+00:00Categories: (h) Chemical tests, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:48b, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Specification experiments|

2:48 describe tests for these anions: Cl⁻, Br⁻ and I⁻ using acidified silver nitrate solution, SO₄²⁻ using acidified barium chloride solution, CO₃²⁻ using hydrochloric acid and identifying the gas evolved

Describe tests for anions: Halide ions (Cl^–, Br^– and I^–)

Underneath are the tests for:

chloride ion test, bromide ion test, iodide ion test, sulfate ion test, carbonate ion test

Describe tests for anions: Sulfate ions (SO₄²^–)

Describe tests for anions: Carbonate ions (CO₃^2-)

Hydr0Gen2023-01-10T11:44:43+00:00Categories: (h) Chemical tests, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:48, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Acids & Salts, Topic: Acids & Salts (D), Topic: Specification experiments|

2:49 describe a test for the presence of water using anhydrous copper(II) sulfate

Add anhydrous copper (II) sulfate (CuSO₄) to a sample.

If water is present the anhydrous copper (II) sulfate will change from white to blue.

Hydr0Gen2023-01-10T11:15:29+00:00Categories: (h) Chemical tests, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:49, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Equilibria, Topic: Reversible Reactions (D), Topic: Simple Molecules & Covalent Bonding, Topic: Specification experiments|

2:50 describe a physical test to show whether a sample of water is pure

If the sample is pure water it will boil at 100^oC

Hydr0Gen2023-09-18T11:08:18+00:00Categories: (h) Chemical tests, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: 2:50, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Key Reactions, Topic: Specification experiments|

3:01 know that chemical reactions in which heat energy is given out are described as exothermic, and those in which heat energy is taken in are described as endothermic

Exothermic: chemical reaction in which heat energy is given out.

Endothermic: chemical reaction in which heat energy is taken in.

(So, in an exothermic reaction the heat exits from the chemicals so temperature rises)

Hydr0Gen2023-01-10T11:58:37+00:00Categories: (a) Energetics, 3 Physical chemistry, Edexcel iGCSE Chemistry|Tags: 3:01, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Crude oil to plastics (D), Topic: Energetics, Topic: Energetics (D), Topic: Equilibria, Topic: Fundamentals, Topic: Reversible Reactions (D)|

3:02 describe simple calorimetry experiments for reactions such as combustion, displacement, dissolving and neutralisation

Calorimetry allows for the measurement of the amount of energy transferred in a chemical reaction to be calculated.

EXPERIMENT1: Displacement, dissolving and neutralisation reactions

Example: magnesium displacing copper from copper(II) sulfate

Method:

50 cm³ of copper(II) sulfate is measured and transferred into a polystyrene cup.
The initial temperature of the copper sulfate solution is measured and recorded.
Magnesium is added and the maximum temperature is measured and recorded.
The temperature rise is then calculated. For example:

Initial temp. of solution (^oC)	Maximium temp. of solution (^oC)	Temperature rise (^oC)
24.2	56.7	32.5

Note: mass of 50 cm³ of solution is 50 g

The cup used is polystyrene because:

polystyrene is an insulator which reduces heats loss

EXPERIMENT2: Combustion reactions

To measure the amount of energy produced when a fuel is burnt, the fuel is burnt and the flame is used to heat up some water in a copper container

Example: ethanol is burnt in a small spirit burner

Method:

The initial mass of the ethanol and spirit burner is measured and recorded.
100cm³ of water is transferred into a copper container and the initial temperature is measured and recorded.
The burner is placed under of copper container and then lit.
The water is stirred constantly with the thermometer until the temperature rises by, say, 30 ^oC
The flame is extinguished and the maximum temperature of the water is measured and recorded.
The burner and the remaining ethanol is reweighed. For example:

Mass of water (g)	Initial temp of water (^oC)	Maximum temp of water (^oC)	Temperature rise (^oC)	Initial mass of spirit burner + ethanol (g)	Final mass of spirit burner + ethanol (g)	Mass of ethanol burnt (g)
100	24.2	54.2	30.0	34.46	33.68	0.78

The amount of energy produced per gram of ethanol burnt can also be calculated:

Hydr0Gen2023-01-10T11:58:51+00:00Categories: (a) Energetics, 3 Physical chemistry, Edexcel iGCSE Chemistry|Tags: 3:02, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Crude oil to plastics (D), Topic: Energetics, Topic: Energetics (D), Topic: Equilibria, Topic: Reversible Reactions (D), Topic: Specification experiments|

3:03 calculate the heat energy change from a measured temperature change using the expression Q = mcΔT

Calorimetry allows for the measurement of the amount of energy transferred in a chemical reaction to be calculated.

EXPERIMENT1: Displacement, dissolving and neutralisation reactions

Example: magnesium displacing copper from copper(II) sulfate

Method:

50 cm³ of copper(II) sulfate is measured and transferred into a polystyrene cup.
The initial temperature of the copper sulfate solution is measured and recorded.
Magnesium is added and the maximum temperature is measured and recorded.
The temperature rise is then calculated. For example:

Initial temp. of solution (^oC)	Maximium temp. of solution (^oC)	Temperature rise (^oC)
24.2	56.7	32.5

Note: mass of 50 cm³ of solution is 50 g

EXPERIMENT2: Combustion reactions

To measure the amount of energy produced when a fuel is burnt, the fuel is burnt and the flame is used to heat up some water in a copper container

Example: ethanol is burnt in a small spirit burner

Method:

The initial mass of the ethanol and spirit burner is measured and recorded.
100cm³ of water is transferred into a copper container and the initial temperature is measured and recorded.
The burner is placed under of copper container and then lit.
The water is stirred constantly with the thermometer until the temperature rises by, say, 30 ^oC
The flame is extinguished and the maximum temperature of the water is measured and recorded.
The burner and the remaining ethanol is reweighed. For example:

Mass of water (g)	Initial temp of water (^oC)	Maximum temp of water (^oC)	Temperature rise (^oC)	Initial mass of spirit burner + ethanol (g)	Final mass of spirit burner + ethanol (g)	Mass of ethanol burnt (g)
100	24.2	54.2	30.0	34.46	33.68	0.78

The amount of energy produced per gram of ethanol burnt can also be calculated:

Hydr0Gen2023-03-06T17:22:19+00:00Categories: (a) Energetics, 3 Physical chemistry, Edexcel iGCSE Chemistry|Tags: 3:03, Chemistry, Double, GCSE, GCSE_Chemistry_Single, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Crude oil to plastics (D), Topic: Energetics, Topic: Energetics (D), Topic: Key Calculations|

3:04 calculate the molar enthalpy change (ΔH) from the heat energy change, Q

Hydr0Gen2023-01-10T11:59:12+00:00Categories: (a) Energetics, 3 Physical chemistry, Edexcel iGCSE Chemistry|Tags: 3:04, Chemistry, Double, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Crude oil to plastics (D), Topic: Energetics, Topic: Energetics (D), Topic: Key Calculations|

3:05 (Triple only) draw and explain energy level diagrams to represent exothermic and endothermic reactions

The symbol ΔH is used to represent the change in heat (or enthalpy change) of a reaction.

ΔH is measured in kJ/mol (kilojoules per mole).

The change in heat (enthalpy change) can be represented on an energy level diagram. ΔH must also labelled.

In an exothermic reaction, the reactants have more energy than the products.

Energy is given out in the form of heat which warms the surroundings.

ΔH is given a negative sign, because the reactants are losing energy as heat, e.g ΔH = -211 kJ/mol.

In an endothermic reaction, the reactants have less energy than the products.

Energy is taken in which cools the surroundings.

ΔH is given a positive sign, because the reactants are gaining energy, e.g ΔH = +211 kJ/mol.

Hydr0Gen2019-05-13T15:37:05+00:00Categories: (a) Energetics, 3 Physical chemistry, Edexcel iGCSE Chemistry|Tags: 3:05, Chemistry, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Energetics, Topic: Equilibria, Triple|

3:06 (Triple only) know that bond-breaking is an endothermic process and that bond-making is an exothermic process

During chemical reactions, the bonds in the reactants must be broken, and new ones formed to make the products.

Breaking bonds need energy and therefore is described as endothermic.

Energy is released when new bonds are made and therefore is described as exothermic.

If bonds are both broken and made during chemical reactions, why can a reaction overall be describe as either exothermic or endothermic?

Example: hydrogen reacts with oxygen producing water. Overall energy is released and therefore the reaction is exothermic.

The reaction is exothermic because the energy needed to break the bonds is less than the energy released in making new bonds.

If a reaction is endothermic then the energy needed to break the bonds is more than the energy released in making new bonds.

Hydr0Gen2019-05-13T15:36:56+00:00Categories: (a) Energetics, 3 Physical chemistry, Edexcel iGCSE Chemistry|Tags: 3:06, Chemistry, GCSE, GCSE_Chemistry_SpecPoint, SpecPoint, Topic: Energetics, Triple|